Multi Agency Radiological Laboratory Analytical Protocols Manual (MARLAP) Analysis

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NUREG-1576
EPA 402-B-04-001B
NTIS PB2004-105421

Multi-Agency Radiological
Laboratory Analytical Protocols Manual
Volume II: Chapters 10 – 17 and Appendix F

PLANNING

MARLAP
ASSESSMENT

IMPLEMENTATION

July 2004

Disclaimer
References within this manual to any specific commercial product, process, or service by trade
name, trademark, manufacturer, or otherwise does not necessarily imply its endorsement or
recommendation by the United States Government. Neither the United States Government nor
any agency or branch thereof, nor any of their employees, makes any warranty, expressed or
implied, nor assumes any legal liability of responsibility for any third party’s use, or the results
of such use, of any information, apparatus, product, or process disclosed in this manual, nor
represents that its use by such third party would not infringe on privately owned rights.

NUREG-1576
EPA 402-B-04-001B
NTIS PB2004-105421

Multi-Agency Radiological
Laboratory Analytical Protocols Manual
(MARLAP)
Part II: Chapters 10 – 17
Appendix F
(Volume II)

United States Environmental Protection Agency
United States Department of Defense
United States Department of Energy
United States Department of Homeland Security
United States Nuclear Regulatory Commission
United States Food and Drug Administration
United States Geological Survey
National Institute of Standards and Technology

July 2004

FOREWORD
MARLAP is organized into two parts. Part I, consisting of Chapters 1 through 9, is intended
primarily for project planners and managers. Part I introduces the directed planning process
central to MARLAP and provides guidance on project planning with emphasis on radioanalytical
planning issues and radioanalytical data requirements. Part II, consisting of Chapters 10 through
20, is intended primarily for laboratory personnel and provides guidance in the relevant areas of
radioanalytical laboratory work. In addition, MARLAP contains seven appendices—labeled A
through G—that provide complementary information, detail background information, or concepts
pertinent to more than one chapter. Six chapters and one appendix are immediately followed by
one or more attachments that the authors believe will provide additional or more detailed
explanations of concepts discussed within the chapter. Attachments to chapters have letter
designators (e.g, Attachment “6A” or “3B”), while attachments to appendices are numbered (e.g.,
“B1”). Thus, “Section B.1.1” refers to section 1.1 of appendix B, while “Section B1.1” refers to
section 1 of attachment 1 to appendix B. Cross-references within the text are explicit in order to
avoid confusion.
Because of its length, the printed version of MARLAP is bound in three volumes. Volume I
(Chapters 1 through 9 and Appendices A through E) contains Part I. Because of its length, Part II
is split between Volumes II and III. Volume II (Chapters 10 through 17 and Appendix F) covers
most of the activities performed at radioanalytical laboratories, from field and sampling issues
that affect laboratory measurements through waste management. Volume III (Chapters 18
through 20 and Appendix G) covers laboratory quality control, measurement uncertainty and
detection and quantification capability. Each volume includes a table of contents, list of
acronyms and abbreviations, and a complete glossary of terms.
MARLAP and its periodic revisions are available online at www.epa.gov/radiation/marlap and
www.nrc.gov/reading-rm/doc-collections/nuregs/staff/sr1576/. The online version is updated
periodically and may differ from the last printed version. Although references to material found
on a web site bear the date the material was accessed, the material available on the date cited may
subsequently be removed from the site. Printed and CD-ROM versions of MARLAP are
available through the National Technical Information Service (NTIS). NTIS may be accessed
online at www.ntis.gov. The NTIS Sales Desk can be reached between 8:30 a.m. and 6:00 p.m.
Eastern Time, Monday through Friday at 1-800-553-6847; TDD (hearing impaired only) at 703487-4639 between 8:30 a.m. and 5:00 p.m Eastern Time, Monday through Friday; or fax at 703605-6900.
MARLAP is a living document, and future editions are already under consideration. Users are
urged to provide feedback on how MARLAP can be improved. While suggestions may not
always be acknowledged or adopted, commentors may be assured that they will be considered
carefully. Comments may be submitted electronically through a link on EPA’s MARLAP web
site (www.epa.gov/radiation/marlap).
JULY 2004

III

MARLAP

CONTENTS (VOLUME II)
Page
List of Figures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

XVIII

List of Tables . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
Acronyms and Abbreviations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

XX

XXIII

Unit Conversion Factors . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . XXXI
10 Field and Sampling Issues That Affect Laboratory Measurements . . . . . . . . . . . . . . . . . 10-1
Part A: Generic Issues . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-1
10.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-1
10.2 Field Sampling Plan: Non-Matrix-Specific Issues . . . . . . . . . . . . . . . . . . . . . . . . . 10-3
10.2.1 Determination of Analytical Sample Size . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-3
10.2.2 Field Equipment and Supply Needs . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-3
10.2.3 Selection of Sample Containers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-4
10.2.3.1 Container Material . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-4
10.2.3.2 Container Opening and Closure . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-5
10.2.3.3 Sealing Containers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-5
10.2.3.4 Precleaned and Extra Containers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-5
10.2.4 Container Label and Sample Identification Code . . . . . . . . . . . . . . . . . . . . . . 10-6
10.2.5 Field Data Documentation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-7
10.2.6 Field Tracking, Custody, and Shipment Forms . . . . . . . . . . . . . . . . . . . . . . . . 10-8
10.2.7 Chain of Custody . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-9
10.2.8 Field Quality Control . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-10
10.2.9 Decontamination of Field Equipment . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-10
10.2.10 Packing and Shipping . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-11
10.2.11 Worker Health and Safety Plan . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-12
10.2.11.1 Physical Hazards . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-13
10.2.11.2 Biohazards . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-15
Part B: Matrix-Specific Issues That Impact Field Sample Collection, Processing, and
Preservation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-16
10.3 Liquid Samples . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-17
10.3.1 Liquid Sampling Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-18
10.3.2 Liquid Sample Preparation: Filtration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-18
10.3.2.1 Example of Guidance for Ground-Water Sample Filtration . . . . . . . 10-19
10.3.2.2 Filters . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-21
10.3.3 Field Preservation of Liquid Samples . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-22
10.3.3.1 Sample Acidification . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-22
10.3.3.2 Non-Acid Preservation Techniques . . . . . . . . . . . . . . . . . . . . . . . . . . 10-23
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Contents
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10.3.4 Liquid Samples: Special Cases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.3.4.1 Radon-222 in Water . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.3.4.1 Milk . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.3.5 Nonaqueous Liquids and Mixtures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.4 Solids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.4.1 Soils . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.4.1.1 Soil Sample Preparation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.4.1.2 Sample Ashing . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.4.2 Sediments . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.4.3 Other Solids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.4.3.1 Structural Materials . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.4.3.2 Biota: Samples of Plant and Animal Products . . . . . . . . . . . . . . . . . .
10.5 Air Sampling . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.5.1 Sampler Components and Operation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.5.2 Filter Selection Based on Destructive Versus Nondestructive Analysis . . . .
10.5.3 Sample Preservation and Storage . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.5.4 Special Cases: Collection of Gaseous and Volatile Air Contaminants . . . . .
10.5.4.1 Radioiodines . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.5.4.2 Gases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.5.4.3 Tritium Air Sampling . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.5.4.4 Radon Sampling in Air . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.6 Wipe Sampling for Assessing Surface Contamination . . . . . . . . . . . . . . . . . . . .
10.6.1 Sample Collection Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.6.1.1 Dry Wipes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.6.1.2 Wet Wipes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.6.2 Sample Handling . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
10.6.3 Analytical Considerations for Wipe Material Selection . . . . . . . . . . . . . . . .
10.7 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

10-25
10-25
10-26
10-26
10-28
10-29
10-29
10-30
10-30
10-31
10-31
10-31
10-34
10-34
10-35
10-36
10-36
10-36
10-37
10-38
10-39
10-41
10-42
10-42
10-43
10-44
10-44
10-45

11 Sample Receipt, Inspection, and Tracking . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
11.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
11.2 General Considerations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
11.2.1 Communication Before Sample Receipt . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
11.2.2 Standard Operating Procedures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
11.2.3 Laboratory License . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
11.2.4 Sample Chain-of-Custody . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
11.3 Sample Receipt . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
11.3.1 Package Receipt . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
11.3.2 Radiological Surveying . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
11.3.3 Corrective Action . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
MARLAP

VI

11-1
11-1
11-1
11-1
11-3
11-4
11-4
11-5
11-5
11-6
11-8

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Page
11.4 Sample Inspection . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11-8
11.4.1 Physical Integrity of Package and Sample Containers . . . . . . . . . . . . . . . . . . . 11-8
11.4.2 Sample Identity Confirmation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11-9
11.4.3 Confirmation of Field Preservation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11-9
11.4.4 Presence of Hazardous Materials . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11-9
11.4.5 Corrective Action . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11-10
11.5 Laboratory Sample Tracking . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11-11
11.5.1 Sample Log-In . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11-11
11.5.2 Sample Tracking During Analyses . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11-11
11.5.3 Storage of Samples . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11-12
11.6 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11-13
12 Laboratory Sample Preparation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-1
12.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-1
12.2 General Guidance for Sample Preparation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-2
12.2.1 Potential Sample Losses During Preparation . . . . . . . . . . . . . . . . . . . . . . . . . 12-2
12.2.1.1 Losses as Dust or Particulates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-2
12.2.1.2 Losses Through Volatilization . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-3
12.2.1.3 Losses Due to Reactions Between Sample and Container . . . . . . . . . . 12-5
12.2.2 Contamination from Sources in the Laboratory . . . . . . . . . . . . . . . . . . . . . . . . 12-6
12.2.2.1 Airborne Contamination . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-7
12.2.2.2 Contamination of Reagents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-7
12.2.2.3 Contamination of Glassware and Equipment . . . . . . . . . . . . . . . . . . . 12-8
12.2.2.4 Contamination of Facilities . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-8
12.2.3 Cleaning of Labware, Glassware, and Equipment . . . . . . . . . . . . . . . . . . . . . . 12-8
12.2.3.1 Labware and Glassware . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-8
12.2.3.2 Equipment . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-10
12.3 Solid Samples . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-12
12.3.1 General Procedures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-12
12.3.1.1 Exclusion of Material . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-14
12.3.1.2 Principles of Heating Techniques for Sample Pretreatment . . . . . . . 12-14
12.3.1.3 Obtaining a Constant Weight . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-23
12.3.1.4 Subsampling . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-24
12.3.2 Soil/Sediment Samples . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-27
12.3.2.1 Soils . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-28
12.3.2.2 Sediments . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-28
12.3.3 Biota Samples . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-28
12.3.3.1 Food . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-29
12.3.3.2 Vegetation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-29
12.3.3.3 Bone and Tissue . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-30
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12.3.4 Other Samples . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.4 Filters . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.5 Wipe Samples . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.6 Liquid Samples . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.6.1 Conductivity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.6.2 Turbidity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.6.3 Filtration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.6.4 Aqueous Liquids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.6.5 Nonaqueous Liquids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.6.6 Mixtures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.6.6.1 Liquid-Liquid Mixtures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.6.6.2 Liquid-Solid Mixtures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.7 Gases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.8 Bioassay . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.9 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.9.1 Cited Sources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
12.9.2 Other Sources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

12-30
12-30
12-31
12-32
12-32
12-32
12-33
12-33
12-34
12-35
12-35
12-35
12-36
12-36
12-37
12-37
12-43

13 Sample Dissolution . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-1
13.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-1
13.2 The Chemistry of Dissolution . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-2
13.2.1 Solubility and the Solubility Product Constant, Ksp . . . . . . . . . . . . . . . . . . . . 13-2
13.2.2 Chemical Exchange, Decomposition, and Simple Rearrangement Reactions . 13-3
13.2.3 Oxidation-Reduction Processes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-4
13.2.4 Complexation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-5
13.2.5 Equilibrium: Carriers and Tracers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-6
13.3 Fusion Techniques . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-6
13.3.1 Alkali-Metal Hydroxide Fusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-9
13.3.2 Boron Fusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-11
13.3.3 Fluoride Fusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-12
13.3.4 Sodium Hydroxide Fusion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-12
13.4 Wet Ashing and Acid Dissolution Techniques . . . . . . . . . . . . . . . . . . . . . . . . . . 13-12
13.4.1 Acids and Oxidants . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-13
13.4.2 Acid Digestion Bombs . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-20
13.5 Microwave Digestion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-21
13.5.1 Focused Open-Vessel Systems . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-21
13.5.2 Low-Pressure, Closed-Vessel Systems . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-22
13.5.3 High-Pressure, Closed-Vessel Systems . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-22
13.6 Verification of Total Dissolution . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-23
13.7 Special Matrix Considerations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-23
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13.7.1 Liquid Samples . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
13.7.2 Solid Samples . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
13.7.3 Filters . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
13.7.4 Wipe Samples . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
13.8 Comparison of Total Dissolution and Acid Leaching . . . . . . . . . . . . . . . . . . . . .
13.9 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
13.9.1 Cited References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
13.9.2 Other Sources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

13-23
13-24
13-24
13-24
13-25
13-27
13-27
13-29

14 Separation Techniques . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-1
14.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-1
14.2 Oxidation-Reduction Processes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-2
14.2.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-2
14.2.2 Oxidation-Reduction Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-3
14.2.3 Common Oxidation States . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-6
14.2.4 Oxidation State in Solution . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-10
14.2.5 Common Oxidizing and Reducing Agents . . . . . . . . . . . . . . . . . . . . . . . . . . 14-11
14.2.6 Oxidation State and Radiochemical Analysis . . . . . . . . . . . . . . . . . . . . . . . . 14-13
14.3 Complexation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-18
14.3.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-18
14.3.2 Chelates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-20
14.3.3 The Formation (Stability) Constant . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-22
14.3.4 Complexation and Radiochemical Analysis . . . . . . . . . . . . . . . . . . . . . . . . . 14-23
14.3.4.1 Extraction of Laboratory Samples and Ores . . . . . . . . . . . . . . . . . . . . 14-23
14.3.4.2 Separation by Solvent Extraction and Ion-Exchange Chromatography 14-23
14.3.4.3 Formation and Dissolution of Precipitates . . . . . . . . . . . . . . . . . . . . . 14-24
14.3.4.4 Stabilization of Ions in Solution . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-24
14.3.4.5 Detection and Determination . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-25
14.4 Solvent Extraction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-25
14.4.1 Extraction Principles . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-25
14.4.2 Distribution Coefficient . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-26
14.4.3 Extraction Technique . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-27
14.4.4 Solvent Extraction and Radiochemical Analysis . . . . . . . . . . . . . . . . . . . . . . 14-30
14.4.5 Solid-Phase Extraction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-32
14.4.5.1 Extraction Chromatography Columns . . . . . . . . . . . . . . . . . . . . . . . . 14-33
14.4.5.2 Extraction Membranes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-34
14.4.6 Advantages and Disadvantages of Solvent Extraction . . . . . . . . . . . . . . . . . 14-35
14.4.6.1 Advantages of Liquid-Liquid Solvent Extraction . . . . . . . . . . . . . . . 14-35
14.4.6.2 Disadvantages of Liquid-Liquid Solvent Extraction . . . . . . . . . . . . . 14-35
14.4.6.3 Advantages of Solid-Phase Extraction Media . . . . . . . . . . . . . . . . . . 14-35
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14.4.6.4 Disadvantages of Solid-Phase Extraction Media . . . . . . . . . . . . . . . .
14.5 Volatilization and Distillation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.5.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.5.2 Volatilization Principles . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.5.3 Distillation Principles . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.5.4 Separations in Radiochemical Analysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.5.5 Advantages and Disadvantages of Volatilization . . . . . . . . . . . . . . . . . . . . .
14.5.5.1 Advantages . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.5.5.2 Disadvantages . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.6 Electrodeposition . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.6.1 Electrodeposition Principles . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.6.2 Separation of Radionuclides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.6.3 Preparation of Counting Sources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.6.4 Advantages and Disadvantages of Electrodeposition . . . . . . . . . . . . . . . . . .
14.6.4.1 Advantages . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.6.4.2 Disadvantages . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.7 Chromatography . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.7.1 Chromatographic Principles . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.7.2 Gas-Liquid and Liquid-Liquid Phase Chromatography . . . . . . . . . . . . . . . . .
14.7.3 Adsorption Chromatography . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.7.4 Ion-Exchange Chromatography . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.7.4.1 Principles of Ion Exchange . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.7.4.2 Resins . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.7.5 Affinity Chromatography . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.7.6 Gel-Filtration Chromatography . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.7.7 Chromatographic Laboratory Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.7.8 Advantages and Disadvantages of Chromatographic Systems . . . . . . . . . . .
14.7.8.1 Advantages . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.7.8.2 Disadvantages . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.8 Precipitation and Coprecipitation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.8.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.8.2 Solutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.8.3 Precipitation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.8.3.1 Solubility and the Solubility Product Constant, Ksp . . . . . . . . . . . . . .
14.8.3.2 Factors Affecting Precipitation . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.8.3.3 Optimum Precipitation Conditions . . . . . . . . . . . . . . . . . . . . . . . . . .
14.8.4 Coprecipitation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.8.4.1 Coprecipitation Processes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.8.4.2 Water as an Impurity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
14.8.4.3 Postprecipitation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
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14-36
14-36
14-36
14-36
14-38
14-39
14-40
14-40
14-40
14-41
14-41
14-42
14-43
14-43
14-43
14-43
14-44
14-44
14-45
14-45
14-46
14-46
14-48
14-54
14-54
14-55
14-56
14-56
14-56
14-56
14-56
14-57
14-59
14-59
14-64
14-69
14-69
14-70
14-74
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14.8.4.4 Coprecipitation Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-75
14.8.5 Colloidal Precipitates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-78
14.8.6 Separation of Precipitates . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-81
14.8.7 Advantages and Disadvantages of Precipitation and Coprecipitation . . . . . . 14-82
14.8.7.1 Advantages . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-82
14.8.7.2 Disadvantages . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-82
14.9 Carriers and Tracers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-82
14.9.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-82
14.9.2 Carriers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-83
14.9.2.1 Isotopic Carriers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-83
14.9.2.2 Nonisotopic Carriers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-84
14.9.2.3 Common Carriers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-85
14.9.2.4 Holdback Carriers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-89
14.9.2.5 Yield of Isotopic Carriers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-89
14.9.3 Tracers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-90
14.9.3.1 Characteristics of Tracers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-92
14.9.3.2 Coprecipitation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-93
14.9.3.3 Deposition on Nonmetallic Solids . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-93
14.9.3.4 Radiocolloid Formation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-94
14.9.3.5 Distribution (Partition) Behavior . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-95
14.9.3.6 Vaporization . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-95
14.9.3.7 Oxidation and Reduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-96
14.10 Analysis of Specific Radionuclides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-97
14.10.1 Basic Principles of Chemical Equilibrium . . . . . . . . . . . . . . . . . . . . . . . . 14-97
14.10.2 Oxidation State . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-100
14.10.3 Hydrolysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-100
14.10.4 Polymerization . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-102
14.10.5 Complexation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-103
14.10.6 Radiocolloid Interference . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-103
14.10.7 Isotope Dilution Analysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-104
14.10.8 Masking and Demasking . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-105
14.10.9 Review of Specific Radionuclides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-109
14.10.9.1 Americium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-109
14.10.9.2 Carbon . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-114
14.10.9.3 Cesium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-116
14.10.9.4 Cobalt . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-119
14.10.9.5 Iodine . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-125
14.10.9.6 Neptunium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-132
14.10.9.7 Nickel . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-136
14.10.9.8 Plutonium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-139
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14.10.9.9 Radium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-148
14.10.9.10 Strontium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-155
14.10.9.11 Sulfur and Phosphorus . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-160
14.10.9.12 Technetium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-163
14.10.9.13 Thorium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-169
14.10.9.14 Tritium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-175
14.10.9.15 Uranium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-180
14.10.9.16 Zirconium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-191
14.10.9.17 Progeny of Uranium and Thorium . . . . . . . . . . . . . . . . . . . . . . . . . . 14-198
14.11 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-201
14.12 Selected Bibliography . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-218
14.12.1 Inorganic and Analytical Chemistry . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-218
14.12.2 General Radiochemistry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-219
14.12.3 Radiochemical Methods of Separation . . . . . . . . . . . . . . . . . . . . . . . . . . 14-219
14.12.4 Radionuclides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-220
14.12.5 Separation Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-222
Attachment 14A Radioactive Decay and Equilibrium . . . . . . . . . . . . . . . . . . . . . . . . . . 14-223
14A.1 Radioactive Equilibrium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-223
14A.1.1 Secular Equilibrium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-223
14A.1.2 Transient Equilibrium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-225
14A.1.3 No Equilibrium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-226
14A.1.4 Summary of Radioactive Equilibria . . . . . . . . . . . . . . . . . . . . . . . . . 14-227
14A.1.5 Supported and Unsupported Radioactive Equilibria . . . . . . . . . . . . . . . . 14-228
14A.2 Effects of Radioactive Equilibria on Measurement Uncertainty . . . . . . . . . 14-229
14A.2.1 Issue . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-229
14A.2.2 Discussion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-229
14A.2.3 Examples of Isotopic Distribution: Natural, Enriched, and Depleted
Uranium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-231
14A.3 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-232
15 Quantification of Radionuclides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
15.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
15.2 Instrument Calibrations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
15.2.1 Calibration Standards . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
15.2.2 Congruence of Calibration and Test-Source Geometry . . . . . . . . . . . . . . . . . .
15.2.3 Calibration and Test-Source Homogeneity . . . . . . . . . . . . . . . . . . . . . . . . . . .
15.2.4 Self-Absorption, Attenuation, and Scattering Considerations for Source
Preparations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
15.2.5 Calibration Uncertainty . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
15.3 Methods of Source Preparation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
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15.3.1 Electrodeposition . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-8
15.3.2 Precipitation/Coprecipitation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-11
15.3.3 Evaporation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-12
15.3.4 Thermal Volatilization/Sublimation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-15
15.3.5 Special Source Matrices . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-16
15.3.5.1 Radioactive Gases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-16
15.3.5.2 Air Filters . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-17
15.3.5.3 Swipes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-18
15.4 Alpha Detection Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-18
15.4.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-18
15.4.2 Gas Proportional Counting . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-20
15.4.2.1 Detector Requirements and Characteristics . . . . . . . . . . . . . . . . . . . . . 15-20
15.4.2.2 Calibration and Test Source Preparation . . . . . . . . . . . . . . . . . . . . . . 15-25
15.4.2.3 Detector Calibration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-25
15.4.2.4 Troubleshooting . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-27
15.4.3 Solid-State Detectors . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-29
15.4.3.1 Detector Requirements and Characteristics . . . . . . . . . . . . . . . . . . . . 15-30
15.4.3.2 Calibration- and Test-Source Preparation . . . . . . . . . . . . . . . . . . . . . 15-33
15.4.3.3 Detector Calibration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-33
15.4.3.4 Troubleshooting . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-34
15.4.3.5 Detector or Detector Chamber Contamination . . . . . . . . . . . . . . . . . 15-35
15.4.3.6 Degraded Spectrum . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-37
15.4.4 Fluorescent Detectors . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-38
15.4.4.1 Zinc Sulfide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-38
15.4.4.2 Calibration- and Test-Source Preparation . . . . . . . . . . . . . . . . . . . . . 15-40
15.4.4.3 Detector Calibration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-41
15.4.4.4 Troubleshooting . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-41
15.4.5 Photon Electron Rejecting Alpha Liquid Scintillation (PERALS®) . . . . . . . 15-42
15.4.5.1 Detector Requirements and Characteristics . . . . . . . . . . . . . . . . . . . . 15-42
15.4.5.2 Calibration- and Test-Source Preparation . . . . . . . . . . . . . . . . . . . . . 15-44
15.4.5.3 Detector Calibration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-45
15.4.5.4 Quench . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-45
15.4.5.5 Available Cocktails . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-46
15.4.5.6 Troubleshooting . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-46
15.5 Beta Detection Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-46
15.5.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-46
15.5.2 Gas Proportional Counting/Geiger-Mueller Tube Counting . . . . . . . . . . . . . 15-49
15.5.2.1 Detector Requirements and Characteristics . . . . . . . . . . . . . . . . . . . . 15-49
15.5.2.2 Calibration- and Test-Source Preparation . . . . . . . . . . . . . . . . . . . . . 15-53
15.5.2.3 Detector Calibration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-54
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15.5.2.4. Troubleshooting . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-57
15.5.3 Liquid Scintillation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-57
15.5.3.1 Detector Requirements and Characteristics . . . . . . . . . . . . . . . . . . . . 15-58
15.5.3.2 Calibration- and Test-Source Preparation . . . . . . . . . . . . . . . . . . . . . 15-61
15.5.3.3 Detector Calibration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-62
15.5.3.4 Troubleshooting . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-68
15.6 Gamma Detection Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-68
15.6.1 Sample Preparation Techniques . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-70
15.6.1.1 Containers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-71
15.6.1.2 Gases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-71
15.6.1.3 Liquids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-72
15.6.1.4 Solids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-72
15.6.2 Sodium Iodide Detector . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-73
15.6.2.1 Detector Requirements and Characteristics . . . . . . . . . . . . . . . . . . . . 15-73
15.6.2.2 Operating Voltage . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-76
15.6.2.3 Shielding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-76
15.6.2.4 Background . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-76
15.6.2.5 Detector Calibration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-77
15.6.2.6 Troubleshooting . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-77
15.6.3 High Purity Germanium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-78
15.6.3.1 Detector Requirements and Characteristics . . . . . . . . . . . . . . . . . . . . 15-78
15.6.3.2 Gamma Spectrometer Calibration . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-82
15.6.3.3 Troubleshooting . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-84
15.6.4 Extended Range Germanium Detectors . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-88
15.6.4.1 Detector Requirements and Characteristics . . . . . . . . . . . . . . . . . . . . 15-89
15.6.4.2 Detector Calibration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-89
15.6.4.3 Troubleshooting . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-90
15.6.5 Special Techniques for Radiation Detection . . . . . . . . . . . . . . . . . . . . . . . . . 15-90
15.6.5.1 Other Gamma Detection Systems . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-90
15.6.5.2 Coincidence Counting . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-91
15.6.5.3 Anti-Coincidence Counting . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-93
15.7 Specialized Analytical Techniques . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-94
15.7.1 Kinetic Phosphorescence Analysis by Laser (KPA) . . . . . . . . . . . . . . . . . . . 15-94
15.7.2 Mass Spectrometry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-95
15.7.2.1 Inductively Coupled Plasma-Mass Spectrometry . . . . . . . . . . . . . . . 15-96
15.7.2.2 Thermal Ionization Mass Spectrometry . . . . . . . . . . . . . . . . . . . . . . . 15-99
15.7.2.3 Accelerator Mass Spectrometry . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-100
15.8 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-101
15.8.1 Cited References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-101
15.8.2 Other Sources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-115
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16 Data Acquisition, Reduction, and Reporting for Nuclear Counting Instrumentation . . . . 16-1
16.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-1
16.2 Data Acquisition . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-2
16.2.1 Generic Counting Parameter Selection . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-3
16.2.1.1 Counting Duration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-4
16.2.1.2 Counting Geometry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-5
16.2.1.3 Software . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-5
16.2.2 Basic Data Reduction Calculations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-6
16.3 Data Reduction on Spectrometry Systems . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-8
16.3.1 Gamma-Ray Spectrometry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-9
16.3.1.1 Peak Search or Identification . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-10
16.3.1.2 Singlet/Multiplet Peaks . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-13
16.3.1.3 Definition of Peak Centroid and Energy . . . . . . . . . . . . . . . . . . . . . . . . 16-14
16.3.1.4 Peak Width Determination . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-15
16.3.1.5 Peak Area Determination . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-17
16.3.1.6 Calibration Reference File . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-19
16.3.1.7 Activity and Concentration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-20
16.3.1.8 Summing Considerations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-21
16.3.1.9 Uncertainty Calculation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-22
16.3.2 Alpha Spectrometry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-23
16.3.2.1 Radiochemical Yield . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-27
16.3.2.2 Uncertainty Calculation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-28
16.3.3 Liquid Scintillation Spectrometry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-29
16.3.3.1 Overview of Liquid Scintillation Counting . . . . . . . . . . . . . . . . . . . . . . 16-29
16.3.3.2 Liquid Scintillation Spectra . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-29
16.3.3.3 Pulse Characteristics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-29
16.3.3.4 Coincidence Circuitry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-30
16.3.3.5 Quenching . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-30
16.3.3.6 Luminescence . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-31
16.3.3.7 Test-Source Vials . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-31
16.3.3.8 Data Reduction for Liquid Scintillation Counting . . . . . . . . . . . . . . . 16-31
16.4 Data Reduction on Non-Spectrometry Systems . . . . . . . . . . . . . . . . . . . . . . . . . . 16-32
16.5 Internal Review of Data by Laboratory Personnel . . . . . . . . . . . . . . . . . . . . . . . . 16-36
16.5.1 Primary Review . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-37
16.5.2 Secondary Review . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-37
16.6 Reporting Results . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-38
16.6.1 Sample and Analysis Method Identification . . . . . . . . . . . . . . . . . . . . . . . . . 16-38
16.6.2 Units and Radionuclide Identification . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-38
16.6.3 Values, Uncertainty, and Significant Figures . . . . . . . . . . . . . . . . . . . . . . . . . 16-39
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16.7 Data Reporting Packages . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
16.8 Electronic Data Deliverables . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
16.9 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
16.9.1 Cited References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
16.9.2 Other Sources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

16-39
16-41
16-41
16-41
16-44

17 Waste Management in a Radioanalytical Laboratory . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-1
17.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-1
17.2 Types of Laboratory Wastes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-1
17.3 Waste Management Program . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-2
17.3.1 Program Integration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-3
17.3.2 Staff Involvement . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-3
17.4 Waste Minimization . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-3
17.5 Waste Characterization . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-6
17.6 Specific Waste Management Requirements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-6
17.6.1 Sample/Waste Exemptions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-9
17.6.2 Storage . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-9
17.6.2.1 Container Requirements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-10
17.6.2.2 Labeling Requirements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-10
17.6.2.3 Time Constraints . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-11
17.6.2.4 Monitoring Requirements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-11
17.6.3 Treatment . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-12
17.6.4 Disposal . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-12
17.7 Contents of a Laboratory Waste Management Plan/Certification Plan . . . . . . . . 17-13
17.7.1 Laboratory Waste Management Plan . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-13
17.7.2 Waste Certification Plan/Program . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-14
17.8 Useful Web Sites . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-15
17.9 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-17
17.9.1 Cited References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-17
17.9.2 Other Sources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-17

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Appendix (Volume II)
Appendix F Laboratory Subsampling . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . F-1
F.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . F-1
F.2 Basic Concepts . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . F-2
F.3 Sources of Measurement Error . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . F-3
F.3.1 Sampling Bias . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . F-4
F.3.2 Fundamental Error . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . F-5
F.3.3 Grouping and Segregation Error . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . F-6
F.4 Implementation of the Particulate Sampling Theory . . . . . . . . . . . . . . . . . . . . . . . . . . . F-9
F.4.1 The Fundamental Variance . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . F-10
F.4.2 Scenario 1 – Natural Radioactive Minerals . . . . . . . . . . . . . . . . . . . . . . . . . . . F-10
F.4.3 Scenario 2 – Hot Particles . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . F-11
F.4.4 Scenario 3 – Particle Surface Contamination . . . . . . . . . . . . . . . . . . . . . . . . . F-13
F.5 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . F-15
F.6 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . F-16
Glossary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . End of volume

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List of Figures (Volume II)
Figure 10.1 Example of chain-of-custody record . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-9
Figure 11.1 Overview of sample receipt, inspection, and tracking . . . . . . . . . . . . . . . . . . . . . 11-2
Figure 12.1 Degree of error in laboratory sample preparation relative to other activities . . . 12-1
Figure 12.2 Laboratory sample preparation flowchart (for solid samples) . . . . . . . . . . . . . . 12-13
Figure 14.1 Ethylene diamine tetraacetic acid (EDTA) . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-20
Figure 14.2 Crown ethers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-21
Figure 14.3 The behavior of elements in concentrated hydrochloric acid on cation-exchange
resins . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-52
Figure 14.4 The behavior of elements in concentrated hydrochloric acid on anion-exchange
resins . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-53
Figure 14.5 The electrical double layer. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-79
Figure 14A.1 Decay chain for 238U . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-224
Figure 14A.2 Secular equilibrium of 210Pb/210Bi . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-225
Figure 14A.3 Transient equilibrium of 95Zr/95Nb . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-226
Figure 14A.4 No equilibrium of 239U/239Np . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-227
Figure 15.1 Alpha plateau generated by a 210Po source on a GP counter using P-10 gas . . . 15-23
Figure 15.2 Gas proportional counter self-absorption curve for 230Th . . . . . . . . . . . . . . . . . 15-28
Figure 15.3 Beta plateau generated by a 90Sr/Y source on a GP counter using P-10 gas . . . 15-52
Figure 15.4 Gas proportional counter self-absorption curve for 90Sr/Y . . . . . . . . . . . . . . . . 15-56
Figure 15.5 Representation of a beta emitter energy spectrum . . . . . . . . . . . . . . . . . . . . . . . 15-65
Figure 15.6 Gamma-ray interactions with high-purity germanium . . . . . . . . . . . . . . . . . . . 15-70
Figure 15.7 NaI(Tl) spectrum of 137Cs . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-75
Figure 15.8 Energy spectrum of 22Na . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-80
Figure 15.9 Different geometries for the same germanium detector and the same sample in
different shapes or position . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-83
Figure 15.10 Extended range coaxial germanium detector . . . . . . . . . . . . . . . . . . . . . . . . . . 15-88
Figure 15.11 Typical detection efficiencies comparing extended range with a normal coaxial
germanium detector . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-90
Figure 15.12 Beta-gamma coincidence efficiency curve for 131I . . . . . . . . . . . . . . . . . . . . . . 15-93
Figure 16.1 Gamma-ray spectrum . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-9
Figure 16.2 Gamma-ray analysis flow chart and input parameters . . . . . . . . . . . . . . . . . . . . 16-11
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Figure 16.3
Figure 16.4
Figure 16.5
Figure 16.6

JULY 2004

Low-energy tailing . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
Photopeak baseline continuum . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
Photopeak baseline continuum-step function . . . . . . . . . . . . . . . . . . . . . . . . . .
Alpha spectrum (238U, 235U, 234U, 239/240Pu, 241Am) . . . . . . . . . . . . . . . . . . . . . . .

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16-17
16-18
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List of Tables (Volume II)
Table 10.1 Summary of sample preservation techniques . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-25
Table 11.1 Typical topics addressed in standard operating procedures related to sample receipt,
inspection, and tracking . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11-3
Table 12.1
Table 12.2
Table 12.3
Table 12.4

Examples of volatile radionuclides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-4
Properties of sample container materials . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-5
Examples of dry-ashing temperatures (platinum container) . . . . . . . . . . . . . . . . 12-23
Preliminary ashing temperature for food samples . . . . . . . . . . . . . . . . . . . . . . . 12-29

Table 13.1 Common fusion fluxes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-7
Table 13.2 Examples of acids used for wet ashing . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13-13
Table 13.3 Standard reduction potentials of selected half-reactions at 25 EC . . . . . . . . . . . 13-14
Table 14.1 Oxidation states of elements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-8
Table 14.2 Oxidation states of selected elements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-10
Table 14.3 Redox reagents for radionuclides . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-13
Table 14.4 Common ligands . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-19
Table 14.5 Radioanalytical methods employing solvent extraction . . . . . . . . . . . . . . . . . . . 14-32
Table 14.6 Radioanalytical methods employing extraction chromatography . . . . . . . . . . . . 14-33
Table 14.7 Elements separable by volatilization as certain species . . . . . . . . . . . . . . . . . . . 14-37
Table 14.8 Typical functional groups of ion-exchange resins . . . . . . . . . . . . . . . . . . . . . . . 14-49
Table 14.9 Common ion-exchange resins . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-50
Table 14.10 General solubility behavior of some cations of interest . . . . . . . . . . . . . . . . . . 14-58
Table 14.11 Summary of methods for utilizing precipitation from homogeneous solution . 14-68
Table 14.12 Influence of precipitation conditions on the purity of precipitates . . . . . . . . . . 14-69
Table 14.13 Common coprecipitating agents for radionuclides . . . . . . . . . . . . . . . . . . . . . . 14-76
Table 14.14 Coprecipitation behavior of plutonium and neptunium . . . . . . . . . . . . . . . . . . 14-78
Table 14.15 Atoms and mass of select radionuclides equivalent to 500 dpm . . . . . . . . . . . 14-83
Table 14.16 Masking agents for ions of various metals . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-106
Table 14.17 Masking agents for anions and neutral molecules . . . . . . . . . . . . . . . . . . . . . 14-108
Table 14.18 Common radiochemical oxidizing and reducing agents for iodine . . . . . . . . 14-129
Table 14.19 Redox agents in plutonium chemistry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-142
Table 14A.1 Relationships of radioactive equilibria . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14-228
Table 15.1 Radionuclides prepared by coprecipitation or precipitation . . . . . . . . . . . . . . . . 15-12
Table 15.2 Nuclides for alpha calibration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-20
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Table 15.3 Typical gas operational parameters for gas proportional alpha counting . . . . . . 15-22
Table 15.4 Nuclides for beta calibration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-48
Table 15.5 Typical operational parameters for gas proportional beta counting . . . . . . . . . . 15-50
Table 15.6 Typical FWHM values as a function of energy . . . . . . . . . . . . . . . . . . . . . . . . . 15-79
Table 15.7 Typical percent gamma-ray efficiencies for a 55 percent HPGe detector with various
counting geometries . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15-83
Table 15.8 AMS detection limits for selected radionuclides . . . . . . . . . . . . . . . . . . . . . . . 15-100
Table 16.1 Units for data reporting . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16-39
Table 16.2 Example elements of a radiochemistry data package . . . . . . . . . . . . . . . . . . . . . 16-40
Table 17.1 Examples of laboratory-generated wastes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17-2

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ACRONYMS AND ABBREVIATIONS
AC . . . . . . . . .
ADC . . . . . . . .
AEA . . . . . . . .
AL . . . . . . . . .
AMS . . . . . . . .
ANSI . . . . . . .
AOAC . . . . . .
APHA . . . . . . .
APS . . . . . . . .
ARAR . . . . . .
ASL . . . . . . . .
ASQC . . . . . . .
ASTM . . . . . .
ATD . . . . . . . .

alternating current
analog to digital convertor
Atomic Energy Act
action level
accelerator mass spectrometry
American National Standards Institute
Association of Official Analytical Chemists
American Public Health Association
analytical protocol specification
applicable or relevant and appropriate requirement (CERCLA/Superfund)
analytical support laboratory
American Society for Quality Control
American Society for Testing and Materials
alpha track detector

BGO . . . . . . . . bismuth germanate [detector]
BNL . . . . . . . . Brookhaven National Laboratory (DOE)
BOA . . . . . . . . basic ordering agreement
CAA . . . . . . . .
CC . . . . . . . . .
CEDE . . . . . . .
CERCLA . . . .
c.f. . . . . . . . . .
cfm . . . . . . . . .
CFR . . . . . . . .
CL . . . . . . . . .
CMPO . . . . . .
CMST . . . . . . .
CO . . . . . . . . .
COC . . . . . . . .
COR . . . . . . . .
cpm . . . . . . . . .
cps . . . . . . . . .
CRM . . . . . . . .
CSU . . . . . . . .
CV . . . . . . . . .
CWA . . . . . . .
CWLM . . . . . .
JULY 2004

Clean Air Act
charcoal canisters
committed effective dose equivalent
Comprehensive Environmental Response, Compensation, and Liability Act of
1980 (“Superfund”)
carrier free [tracer]
cubic feet per minute
Code of Federal Regulations
central line (of a control chart)
[octyl(phenyl)]-N,N-diisobutylcarbonylmethylphosphine oxide
Characterization, Monitoring, and Sensor Technology Program (DOE)
contracting officer
chain of custody
contracting officer’s representative
counts per minute
counts per second
(1) continuous radon monitor; (2) certified reference material
combined standard uncertainty
coefficient of variation
Clean Water Act
continuous working level monitor
XXIII

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Acronyms and Abbreviations
d ...........
D...........
DAAP . . . . . . .
DC . . . . . . . . .
DCGL . . . . . . .
DHS . . . . . . . .
DIN . . . . . . . . .
DL . . . . . . . . .
DoD . . . . . . . .
DOE . . . . . . . .
DOELAP . . . .
DOT . . . . . . . .
DOP . . . . . . . .
dpm . . . . . . . .
DPPP . . . . . . .
DQA . . . . . . . .
DQI . . . . . . . . .
DQO . . . . . . . .
DTPA . . . . . . .
DVB . . . . . . . .

day[s]
homogeneous distribution coefficient
diamylamylphosphonate
direct current
derived concentration guideline level
U.S. Department of Homeland Security
di-isopropylnaphthalene
discrimination limit
U.S. Department of Defense
U.S. Department of Energy
DOE Laboratory Accreditation Program
U.S. Department of Transportation
dispersed oil particulate
disintegrations per minute
dipentylpentylphosphonate
data quality assessment
data quality indicator
data quality objective
diethylene triamine pentaacetic acid
divinylbenzene

Ee . . . . . . . . . .
Eβmax . . . . . . . .
EDD . . . . . . . .
EDTA . . . . . . .
EGTA . . . . . . .
EMEDD . . . . .
EPA . . . . . . . .
ERPRIMS . . .
ESC . . . . . . . .
eV . . . . . . . . . .

emission probability per decay event
maximum beta-particle energy
electronic data deliverable
ethylene diamine tetraacetic acid
ethyleneglycol bis(2-aminoethylether)-tetraacetate
environmental management electronic data deliverable (DOE)
U.S. Environmental Protection Agency
Environmental Resources Program Management System (U.S. Air Force)
expedited site characterization; expedited site conversion
electron volts

FAR . . . . . . . .
FBO . . . . . . . .
FDA . . . . . . . .
FEP . . . . . . . . .
fg . . . . . . . . . .
FOM . . . . . . . .
FWHM . . . . . .
FWTM . . . . . .

Federal Acquisition Regulations, CFR Title 48
Federal Business Opportunities [formerly Commerce Business Daily]
U.S. Food and Drug Administration
full energy peak
femtogram
figure of merit
full width of a peak at half maximum
full width of a peak at tenth maximum

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Acronyms and Abbreviations
GC . . . . . . . . .
GLPC . . . . . . .
GM . . . . . . . . .
GP . . . . . . . . .
GUM . . . . . . .
Gy . . . . . . . . . .

gas chromatography
gas-liquid phase chromatography
Geiger-Mueller [detector]
gas proportional [counter]
Guide to the Expression of Uncertainty in Measurement (ISO)
gray[s]

h ...........
H0 . . . . . . . . . .
HA, H1 . . . . . . .
HDBP . . . . . . .
HDEHP . . . . .
HDPE . . . . . . .
HLW . . . . . . .
HPGe . . . . . . .
HPLC . . . . . . .
HTRW . . . . . .

hour[s]
null hypothesis
alternative hypothesis
dibutylphosphoric acid
bis(2-ethylhexyl) phosphoric acid
high-density polyethylene
high-level [radioactive] waste
high-purity germanium
high-pressure liquid chromatography; high-performance liquid chromatography
hazardous, toxic, and radioactive waste

IAEA . . . . . . .
ICRU . . . . . . .
ICP-MS . . . . .
IPPD . . . . . . . .
ISO . . . . . . . . .
IUPAC . . . . . .

International Atomic Energy Agency
International Commission on Radiation Units and Measurements
inductively coupled plasma-mass spectroscopy
integrated product and process development
International Organization for Standardization
International Union of Pure and Applied Chemistry

k . . . . . . . . . . . coverage factor
keV . . . . . . . . . kilo electron volts
KPA . . . . . . . . kinetic phosphorimeter analysis
LAN . . . . . . . .
LANL . . . . . . .
LBGR . . . . . . .
LCL . . . . . . . .
LCS . . . . . . . .
LDPE . . . . . . .
LEGe . . . . . . .
LIMS . . . . . . .
LLD . . . . . . . .
LLNL . . . . . . .
LLRW . . . . . .
LLRWPA . . . .
JULY 2004

local area network
Los Alamos National Laboratory (DOE)
lower bound of the gray region
lower control limit
laboratory control samples
low-density polyethylene
low-energy germanium
laboratory information management system
lower limit of detection
Lawrence Livermore National Laboratory (DOE)
low-level radioactive waste
Low Level Radioactive Waste Policy Act
XXV

MARLAP

Acronyms and Abbreviations
LOMI . . . . . . .
LPC . . . . . . . .
LS . . . . . . . . . .
LSC . . . . . . . .
LWL . . . . . . . .

low oxidation-state transition-metal ion
liquid-partition chromatography; liquid-phase chromatography
liquid scintillation
liquid scintillation counter
lower warning limit

MAPEP . . . . .
MARSSIM . . .
MCA . . . . . . .
MCL . . . . . . . .
MDA . . . . . . .
MDC . . . . . . .
MDL . . . . . . . .
MeV . . . . . . . .
MIBK . . . . . . .
min . . . . . . . . .
MPa . . . . . . . .
MQC . . . . . . .
MQO . . . . . . .
MS . . . . . . . . .
MSD . . . . . . . .
MVRM . . . . . .

Mixed Analyte Performance Evaluation Program (DOE)
Multi-Agency Radiation Survey and Site Investigation Manual
multichannel analyzer
maximum contaminant limit
minimum detectable amount; minimum detectable activity
minimum detectable concentration
method detection limit
mega electron volts
methyl isobutyl ketone
minute[s]
megapascals
minimum quantifiable concentration
measurement quality objective
matrix spike; mass spectrometer
matrix spike duplicate
method validation reference material

NAA . . . . . . . .
NaI(Tl) . . . . . .
NCP . . . . . . . .
NCRP . . . . . . .
NELAC . . . . .
NESHAP . . . .
NIM . . . . . . . .
NIST . . . . . . . .
NPL . . . . . . . .

neutron activation analysis
thallium-activated sodium iodide [detector]
National Oil and Hazardous Substances Pollution Contingency Plan
National Council on Radiation Protection and Measurement
National Environmental Laboratory Accreditation Conference
National Emission Standards for Hazardous Air Pollutants (EPA)
nuclear instrumentation module
National Institute of Standards and Technology
National Physics Laboratory (United Kingdom); National Priorities List (United
States)
U.S. Nuclear Regulatory Commission
NIST Radiochemistry Intercomparison Program
nitrilotriacetate
nephelometric turbidity units
National Voluntary Laboratory Accreditation Program (NIST)

NRC . . . . . . . .
NRIP . . . . . . .
NTA (NTTA) .
NTU . . . . . . . .
NVLAP . . . . .

OA . . . . . . . . . observational approach
OFHC . . . . . . . oxygen-free high-conductivity
MARLAP

XXVI

JULY 2004

Acronyms and Abbreviations
OFPP . . . . . . . Office of Federal Procurement Policy
φMR . . . . . . . . .
Pa . . . . . . . . . .
PARCC . . . . .
PBBO . . . . . . .
PCB . . . . . . . .
pCi . . . . . . . . .
pdf . . . . . . . . .
PE . . . . . . . . . .
PERALS . . . . .
PFA . . . . . . . .
PIC . . . . . . . . .
PIPS . . . . . . . .
PM . . . . . . . . .
PMT . . . . . . . .
PT . . . . . . . . . .
PTB . . . . . . . .
PTFE . . . . . . .
PUREX . . . . .
PVC . . . . . . . .

required relative method uncertainty
pascals
precision, accuracy, representativeness, completeness, and comparability
2-(4'-biphenylyl) 6-phenylbenzoxazole
polychlorinated biphenyl
picocurie
probability density function
performance evaluation
Photon Electron Rejecting Alpha Liquid Scintillation®
perfluoroalcoholoxil™
pressurized ionization chamber
planar implanted passivated silicon [detector]
project manager
photomultiplier tube
performance testing
Physikalisch-Technische bundesanstalt (Germany)
polytetrafluoroethylene
plutonium uranium reduction extraction
polyvinyl chloride

QA . . . . . . . . .
QAP . . . . . . . .
QAPP . . . . . . .
QC . . . . . . . . .

quality assurance
Quality Assessment Program (DOE)
quality assurance project plan
quality control

rad . . . . . . . . .
RCRA . . . . . . .
REE . . . . . . . .
REGe . . . . . . .
rem . . . . . . . . .
RFP . . . . . . . .
RFQ . . . . . . . .
RI/FS . . . . . . .
RMDC . . . . . .
ROI . . . . . . . . .
RPD . . . . . . . .
RPM . . . . . . . .
RSD . . . . . . . .
RSO . . . . . . . .

radiation absorbed dose
Resource Conservation and Recovery Act
rare earth elements
reverse-electrode germanium
roentgen equivalent: man
request for proposals
request for quotations
remedial investigation/feasibility study
required minimum detectable concentration
region of interest
relative percent difference
remedial project manager
relative standard deviation
radiation safety officer

JULY 2004

XXVII

MARLAP

Acronyms and Abbreviations
s ...........
SA . . . . . . . . .
SC . . . . . . . . . .
SAFER . . . . . .
SAM . . . . . . . .
SAP . . . . . . . .
SEDD . . . . . . .
SI . . . . . . . . . .
SMO . . . . . . . .
SOP . . . . . . . .
SOW . . . . . . . .
SQC . . . . . . . .
SPE . . . . . . . . .
SR . . . . . . . . . .
SRM . . . . . . . .
SSB . . . . . . . .
SSR . . . . . . . .
Sv . . . . . . . . . .

second[s]
spike activity
critical value
Streamlined Approach for Environmental Restoration Program (DOE)
site assessment manager
sampling and analysis plan
staged electronic data deliverable
international system of units
sample management office[r]
standard operating procedure
statement of work
statistical quality control
solid-phase extraction
unspiked sample result
standard reference material
silicon surface barrier [alpha detector]
spiked sample result
sievert[s]

t½ . . . . . . . . . .
TAT . . . . . . . .
TBP . . . . . . . .
TC . . . . . . . . .
TCLP . . . . . . .
TD . . . . . . . . .
TEC . . . . . . . .
TEDE . . . . . . .
TEC . . . . . . . .
TES . . . . . . . .
TFM . . . . . . . .
TIMS . . . . . . .
TIOA . . . . . . .
TLD . . . . . . . .
TnOA . . . . . . .
TOPO . . . . . . .
TPO . . . . . . . .
TPP . . . . . . . . .
TPU . . . . . . . .
TQM . . . . . . . .
TRUEX . . . . .
TSCA . . . . . . .

half-life
turnaround time
tributylphosphate
to contain
toxicity characteristic leaching procedure
to deliver
technical evaluation committee
total effective dose equivalent
technical evaluation committee (USGS)
technical evaluation sheet (USGS)
tetrafluorometoxil™
thermal ionization mass spectrometry
triisooctylamine
thermoluminescent dosimeter
tri-n-octylamine
trioctylphosphinic oxide
technical project officer
technical project planning
total propagated uncertainty
Total Quality Management
trans-uranium extraction
Toxic Substances Control Act

MARLAP

XXVIII

JULY 2004

Acronyms and Abbreviations
TSDF . . . . . . . treatment, storage, or disposal facility
tSIE . . . . . . . . transfomed spectral index of the external standard
TTA . . . . . . . . thenoyltrifluoroacetone
U...........
uMR . . . . . . . . .
uc(y) . . . . . . . .
UBGR . . . . . .
UCL . . . . . . . .
USACE . . . . .
USGS . . . . . . .
UV . . . . . . . . .
UWL . . . . . . .

expanded uncertainty
required absolute method uncertainty
combined standard uncertainty
upper bound of the gray region
upper control limit
United States Army Corps of Engineers
United States Geological Survey
ultraviolet
upper warning limit

V . . . . . . . . . . . volt[s]
WCP . . . . . . . . waste certification plan
XML . . . . . . . . extensible mark-up language
XtGe® . . . . . . . extended-range germanium
y . . . . . . . . . . . year[s]
Y . . . . . . . . . . . response variable
ZnS(Ag) . . . . . silver-activated zinc sulfide [detector]

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MARLAP

UNIT CONVERSION FACTORS
To Convert
Years (y)

Disintegrations
per second (dps)
Bq
Bq/kg
Bq/m3
Bq/m3
Microcuries per
milliliter
(µCi/mL)
Disintegrations
per minute (dpm)
Gallons (gal)
Gray (Gy)
Roentgen
Equivalent Man
(rem)

JULY 2004

To
Multiply by
Seconds (s)
3.16 × 107
Minutes (min)
5.26 × 105
8.77 × 103
Hours (h)
Becquerels (Bq) 1.0

To Convert
s
min
h
Bq

To
y

dps

Multiply by
3.17 × 10!8
1.90 × 10!6
1.14 × 10!4
1.0

Picocuries (pCi)
pCi/g
pCi/L
Bq/L
pCi/L

27.03
2.7 × 10!2
2.7 × 10!2
103
109

pCi
pCi/g
pCi/L
Bq/L
pCi/L

Bq
Bq/kg
Bq/m3
Bq/m3
µCi/mL

3.7 × 10!2
37
37
10!3
10!9

µCi
pCi
Liters (L)
rad

4.5 × 10!7
4.5 × 10!1
3.78
100

pCi

dpm

2.22

Liters
rad

Gallons
Gy

0.265
10!2

Sievert (Sv)

10!2

Sv

rem

102

XXXI

MARLAP

10 FIELD AND SAMPLING ISSUES THAT AFFECT
LABORATORY MEASUREMENTS
Part A: Generic Issues
10.1 Introduction
This chapter provides guidance to project managers, planners, laboratory personnel, and the
radioanalytical specialists tasked with developing a field sampling plan. It emphasizes those
activities conducted at the time of sample collection and other activities conducted after sample
collection that could affect subsequent laboratory analyses.
A field sampling plan should provide comprehensive guidance for collecting, preparing,
preserving, shipping, and tracking field samples and recording field data. The principal objective
of a well-designed sampling plan is to provide representative samples of the proper size for
analysis. Critical to the sampling plan are outputs of the systematic planning process, which
commonly define the Analytical Protocol Specifications (APSs) and the measurement quality
objectives (MQOs) that must be met. While comprehensive discussions on actual field sample
collection and sampling strategies are beyond the scope of MARLAP, specific aspects of sample
collection methods and the physical preparation and preservation of samples warrant further
discussion because they impact the analytical process and the data quality.
This chapter has two main parts. Part A identifies general elements of a field sampling plan and
provides project planners with general guidance. Part B provides detailed, matrix-specific
guidance and technical data for liquid, solid, airborne, and surface contaminants requiring field
sampling. This information will assist project planners further in the development of standard
operating procedures (SOPs) and training for field personnel engaged in preparation and
preservation of field samples.
Contents

The need to specify sample collection methods,
and to prepare and preserve field samples, is
Part A: Generic Issues . . . . . . . . . . . . . . . . . . . . . . 10-1
10.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . 10-1
commonly dictated by one or more of the
10.2 Field Sampling Plan: Non-Matrix-Specific
following:

Issues . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-3
Part B: Matrix-Specific Issues That Impact Field
• The systematic planning process that
Sample Collection, Processing, and
identifies the type, quality, and quantity of
Preservation . . . . . . . . . . . . . . . . . . . . . . . . . 10-16
data needed to satisfy a decision process;
10.3 Liquid Samples . . . . . . . . . . . . . . . . . . . . . . 10-17
10.4 Solids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10-28
• The potential alteration of field samples by 10.5 Air Sampling . . . . . . . . . . . . . . . . . . . . . . . 10-34
10.6 Wipe Sampling for Assessing Surface
physical, chemical, and biological processes
Contamination . . . . . . . . . . . . . . . . . . . . . . 10-41
during the time between collection and
10.7 References . . . . . . . . . . . . . . . . . . . . . . . . . 10-45

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Field and Sampling Issues That Affect Laboratory Measurements
analysis;
• Requirements specified by the analytical laboratory pertaining to sample analysis;
• Requirements of analytical methods; and
• Requirements of regulators (e.g., Department of Transportation).
10.1.1 The Need for Establishing Channels of Communication
To design an effective sampling plan, it is critical to obtain the input and recommendations of
representatives of (1) the field sampling team, (2) the health physics professional staff, (3) the
analytical laboratory, (4) statistical and data analysts, (5) quality assurance personnel, and (6)
end-users of data.
Beyond the initial input that assist the project planners in the design of the sampling plan, it is
equally important to maintain open channels of communication among key members of the
project team throughout the process. For example, the analytical laboratory should be provided
with contacts within the field sampling team to ensure that modifications, discrepancies, and
changes are addressed and potential problems may be resolved in a timely manner.
Communication among project staff, field personnel, and the laboratory offer a means to
coordinate activities, schedules, and sample receipt. Project planning documents generated from
the systematic planning process, such as APSs and statements of work (SOWs), should be
consulted, but they cannot address all details. Additional communication will be necessary to
convey information about the number and type of samples the laboratory can expect at a certain
time. Documentation with special instructions regarding the samples should be received before
the samples arrive. This information notifies the laboratory of any health and safety concerns so
that laboratory personnel can implement proper contamination management practices. Health and
safety concerns may affect analytical procedures, sample disposition, etc. The analytical
laboratory should have an initial understanding about the relative number of samples that will be
received and the types of analyses that are expected for specific samples. Furthermore, advance
communications allow laboratory staff to adjust to modifications, discrepancies, and changes.
10.1.2 Developing Field Documentation
The field organization must conduct its operations in such a manner as to provide reliable
information that meets the data quality objectives (DQOs). To achieve this goal, all relevant
procedures pertaining to sample collection and processing should be based on documented
standard operating procedures that may include, but are not limited to, the following activities:
• Developing a technical basis for defining the size of individual samples;
MARLAP

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Field and Sampling Issues That Affect Laboratory Measurements
•
•
•
•
•
•
•
•
•
•
•

Selecting field equipment and instrumentation;
Using proper sample containers and preservatives;
Using consistent container labels and sample identification codes;
Documenting field sample conditions and exceptions;
Documenting sample location;
Tracking, accountability, custody, and shipment forms;
Legal accountability, such as chain-of-custody record, when required;
Selecting samples for field quality control (QC) program;
Decontaminating equipment and avoiding sample cross-contamination;
Specifying sample packaging, radiological surveys of samples, shipping, and tracking; and
Documenting the health and safety plan.

10.2 Field Sampling Plan: Non-Matrix-Specific Issues
10.2.1 Determination of Analytical Sample Size
When collecting environmental samples for radiochemical analysis, an important parameter for
field personnel is the mass or volume of an individual sample that must be collected. The
required minimum sample size is best determined through the collective input of project
planners, field technicians, and laboratory personnel who must consider the likely range of the
contaminant concentrations, the type of radiation emitted by constituents or analytes (alpha, beta,
and gamma emitters), field logistics, and the radioanalytical methods that are to be employed. It
is important to have a quantitative understanding of the relationship between sample size and
project specific requirements in order for samples to yield useful data.
10.2.2 Field Equipment and Supply Needs
Before starting field sampling activities, all necessary equipment and supplies should be
identified, checked for proper operation and availability, and—when appropriate—preassembled. Instrumentation and equipment needs will depend not only on the matrix to be
sampled, but also on the accessibility of the matrix and the physical and chemical properties of
radionuclide contaminants under investigation.
In addition to specialized field equipment and instrumentation, field sampling supplies
commonly include, but are not limited to, the following:
• Sampling devices (e.g., trowel, hand auger, soil core sampler, submersible water pump, high
volume air filter, etc.);
• Sampling preparation equipment (e.g., weighing scales, volume measuring devices, soil
screening sieves, water filtering equipment, etc.);
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Field and Sampling Issues That Affect Laboratory Measurements
• Sample preservation equipment and agents (e.g., refrigeration, ice, formaldehyde or acid
additives);
• Personnel protective gear (e.g., respiratory protective devices, protective clothing such as
gloves and booties, life-preservers, etc.);
• Proper writing utensils (e.g., permanent pens and markers);
• Field logbooks and field tracking forms;
• Maps, distance measuring equipment, global positioning systems, or other locationdetermining equipment;
• Field sampling flags or paint;
• Chain-of-custody (COC) forms;
• Sample tags, labels, and documents;
• Appropriately labeled sample containers;
• Shipment containers and packing materials that meet national and international shipping
regulations (see Section 10.2.10);
• Shipment forms;
• Analysis request forms identifying the type of radioanalysis to be performed; and
• Items required by the health and safety plan (medical kit, etc.).
10.2.3 Selection of Sample Containers
There are several physical and chemical characteristics to consider when selecting a suitable
container for shipping and storing samples. These include the container material and its size,
configuration, and method for ensuring a proper seal.
10.2.3.1

Container Material

Sample containers must provide reasonable assurance of maintaining physical integrity (i.e.,
against breakage, rupture, or leakage) during handling, transport, and potentially long periods of
storage. The most important factor to consider in container selection is the chemical
MARLAP

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Field and Sampling Issues That Affect Laboratory Measurements
compatibility between container material and sample. Containers may be made from ordinary
bottle glass, borosilicate glass (such as Pyrex® or Corex®), plastics (e.g., high-density
polyethylene, HDPE), low-density polyethylene, polycarbonate, polyvinyl chloride (PVC),
fluorinated ethylene or propylene (Teflon™), or polymethylpentene. For certain samples, the
choice of containers may require metal construction or be limited to paper envelopes.
10.2.3.2

Container Opening and Closure

A suitable container also should be shaped appropriately for the purpose. For example, a widemouthed container will provide easier access for the introduction and withdrawal of sample
material and eliminate spills or the need for additional tools or equipment (e.g., funnel) that may
become a source of cross contamination among samples.
Equally important is the container’s closure. As a rule, snap-on caps should not be considered for
liquid samples because they do not ensure a proper seal. Even when screw caps are used, it is
frequently prudent to protect against vibration by securing the cap with electrical or duct tape. A
proper seal is important for air samples, such as radon samples. The container cap material, if
different from the container material, must be equally inert with regard to sample constituents.
10.2.3.3

Sealing Containers

Tamper-proof seals offer an additional measure to ensure sample integrity. A simple example
includes placing a narrow strip of paper over a bottle cover and then affixing this to the container
with a wide strip of clear tape (EPA, 1987, Exhibit 5-6 provides examples of custody seals). The
paper strip can be initialed and dated in the field to indicate the staff member who sealed the
sample and the date of the seal. Individually sealing each sample with a custody seal with the
collector’s initials and the date the sample was sealed may be required by the project. The seal
ensures legal defensibility and integrity of the sample at collection. Tamper-proof seals should
only be applied once field processing and preservation steps are completed. Reopening this type
of sealed container in the field might warrant using a new container or collecting another sample.
10.2.3.4

Precleaned and Extra Containers

The reuse of sample containers is discouraged because traces of radionuclides might persist from
initial container use to subsequent use. The use of new containers for each collection removes
doubts concerning radionuclides from previous sampling. New containers might also require
cleaning (ASTM D5245) to remove any plasticizer used in production or to pretreat glass
surfaces. Retaining extra empty containers from a new lot or a special batch of precleaned and
treated containers can provide the laboratory container blanks for use as part of quality control.
Extra containers are also useful for taking additional samples as needed during field collection
and to replace broken or leaking containers.
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Field and Sampling Issues That Affect Laboratory Measurements
10.2.4 Container Label and Sample Identification Code
Each sample can only be identified over the life of a study if a form of permanent identification is
provided with or affixed to the container or available in sample log. The most useful form of
identification utilizes a unique identifier for each sample. Such unique identification codes
ensure the project’s ability to track individual samples. The standard operating procedure (SOP)
that addresses sample identification should describe the method to be used to assure that samples
are properly identified and controlled in a consistent manner. Containers sometimes may be prelabeled with identification numbers already in place.
Any identification recorded on a container or a label affixed to the container should remain with
the container throughout sample processing and storage. The identification information should be
written with a permanent marker—especially if the labels are exposed to liquids. Information can
be recorded directly on the container or on plastic or paper tags securely fixed to the container.
However, tags are more likely to become separated from containers than are properly secured
labels.
Labels, tags, and bar codes should be durable enough so no information is lost or compromised
during field work, sample transport, or laboratory processing. Transparent tape can be used to
cover the label once it is completed. The tape protects the label, adds moisture resistance,
prevents tampering with the sample information, and helps secure the label to the container.
The project manager needs to determine if a field-sample identification (ID) scheme may
introduce bias into the analysis process, such as allowing the laboratory to become aware of
trends or locations from the sample identification. This could influence their judgment about the
anticipated result and thereby introduce actions on the part of laboratory personnel that they
would not otherwise take (such as reanalyzing the sample). The project manager needs to
determine the applicability of electronic field data recorders and the issue of electronic signatures
for the project.
A unique identifier can include a code for a site, the sample location at the site, or a series of
digits identifying the year and day of year (e.g., “1997-127” uses the Julian date, and “062296”
describes a month, day, and year). Alternatively, a series of digits can be assigned sequentially by
site, date, and laboratory destination. The use of compass headings and grid locations also
provides additional unique information (e.g., “NW fence, sampled at grid points: A1 through
C25, 072196, soil”). With this approach, samples arriving at a laboratory are then unique in two
ways. First, each sample can be discriminated from materials collected at other sites. Second, if
repeat samples are made at a single site, then subsequent samples from the same location are
unique only by date. Labeling samples sequentially might not be appropriate for all studies. Bar
coding may reduce transcription errors and should be evaluated for a specific project.

MARLAP

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Field and Sampling Issues That Affect Laboratory Measurements
10.2.5 Field Data Documentation
All information pertinent to field sampling is documented in a log book or on a data form. The
log book should be bound and the pages numbered consecutively, and forms should be pagenumbered and dated. Where the same information is requested routinely, preprinted log books or
data sheets will minimize the effort and will standardize the presentation of data. Even when
standardized preprinted forms are used, all information recorded should be in indelible ink, with
all entry errors crossed out with a single line and initialed. The color of ink used should be
compatible with the need to copy that information. All entries should be dated and signed on the
date of entry. Initials should be legible and traceable, so that it is clear who made the entry.
Whenever appropriate, log or data form entries should contain—but are not limited to—the
following:
• Identification of Project Plan or Sampling Plan;
• Location of sampling (e.g., reference to grid location, maps, photographs, location in a
room);
• Date and time of sample collection;
• Sample matrix (e.g., surface water, soil, sediment, sludge, etc.);
• Suspected radionuclide constituents;
• Sample-specific ID;
• Sample volume, weight, depth;
• Sample type (e.g., grab, composite);
• Sample preparation used (e.g., removal of extraneous matter);
• Sample preservation used;
• Requested analyses to be performed (e.g., gross beta/gamma, gamma spectroscopy for a
specific radionuclide, radiochemical analysis);
• Sample destination, including name and address of analytical laboratory;
• Names of field people responsible for collecting sample;
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Field and Sampling Issues That Affect Laboratory Measurements
• Physical and meteorological conditions at time of sample collection;
• Special handling or safety precautions;
• Results of field radiation measurements, including surveys of sample containers; and
• Signatures or initials of appropriate field personnel. When using initials, ensure that they can
be uniquely identified with an individual.
Labels affixed to individual sample containers should contain key information that forms an
abstract of log book data sheets. When this is not practical, a copy of individual sample data
sheets may be included along with the appropriately ID-labeled sample.
10.2.6 Field Tracking, Custody, and Shipment Forms
A sample tracking procedure must be in place for all projects in order that the proper location and
identification of samples is maintained throughout the process from collection through handling,
preservation, storage, transfer to laboratory, and disposal. The term “tracking” means an
accountability process that meets generally acceptable laboratory practices as described by
accrediting bodies, but is less stringent than a formal chain-of-custody process. Tracking also
develops a record of all individuals responsible for the custody and transfer of the samples.
Chapter 4 (Project Plan Documents) discusses the process of tracking and accountability. Also,
Chapter 11 (Sample Receipt, Inspection, and Tracking) discusses the laboratory process of
tracking.
When transferring the possession of samples, the individuals relinquishing and the individuals
receiving the samples should sign, date, and note the time on the form. A standardized form
should be designed for recording tracking or formal chain-of-custody information related to
tracking sample possession. An example of a COC form is shown in Figure 10.1. Additional
information and examples of custody forms are illustrated by EPA (1987 and 1994). If samples
are to be split and distributed to more than one analytical laboratory, multiple forms will be
needed to accompany sample sets. The sample collector is responsible for initiating the sample
tracking record. The following information is considered minimal for sample tracking:
•
•
•
•
•
•
•
•

Name of project;
Sampler’s signature;
Sample ID;
Sample location
Date and time sampled;
Sample type;
Preservatives;
Number of containers;

MARLAP

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Field and Sampling Issues That Affect Laboratory Measurements
•
•
•
•

Analysis required;
Signatures of persons relinquishing, receiving, and transporting the samples;
Signature for laboratory receipt;
Method of shipment or carrier and air bill when shipped or shipping manifest identification
upon receipt; and
• Comments regarding the integrity of shipping container and individual samples.

10.2.7 Chain of Custody
The legal portion of the tracking and handling process that ensures legal defensibility from
sample collection to data reporting has become relatively standardized and is referred to as the
CHAIN-OF-CUSTODY RECORD
SAMPLED BY:
FIELD
IDENTIFICATION
NUMBER

SAMPLE MATRIX
FIELD
LOCATION

DATE

TIME

Water

Soil

SEQ.
No. of
No. Containers

Analysis
Required

Other

Relinquished by:

Date/Time Received by:
/

Date/Time
/

Relinquished by:

Date/Time Received by:
/

Date/Time
/

Relinquished by:

Date/Time Received by:
/

Date/Time
/

Relinquished by:

Date/Time Received by:
/

Date/Time
/

Relinquished by:

Date/Time Received by laboratory for field analysis:
/

Date/Time
/

Method of Shipment:
Distribution: Orig. - Accompany Shipment
1 Copy – Survey Coordinator Field Files
FIGURE 10.1—Example of chain-of-custody record

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COC process (APHA, 1998). Guidance is provided in ASTM D4840 and NIOSH (1983). The
level of security required to maintain an adequate chain of custody is that necessary to establish a
“reasonable probability” that the sample has not been tampered with. For court proceedings, the
requirements are established in law. COC procedures are important in demonstrating sample
control when litigation is involved. In many cases, federal, state or local agencies may require
that COC be maintained for specific projects. COC is usually not required for samples that are
generated and immediately tested within a facility or continuous (rather than discrete or integrated) samples that are subject to real- or near-real-time analysis (e.g., continuous screening).
When COC is required, the custody information is recorded on a COC form. Chain-of-custody
documents vary by organization and by project. Communication between field and laboratory
personnel is critical to the successful use of COC. Any error made on a custody form is crossed
out with a single line and dated and initialed. Use of correction ink or obliteration of data is not
acceptable. Inform the laboratory when COC is required before the samples are received (see
Section 11.2.4, “Sample Chain-of-Custody,” for further information). The COC documents are
signed by personnel who collect the samples. A COC record accompanies the shipment and one
or more copies are distributed to the project coordinator or other office(s) where field and
laboratory records are maintained.
10.2.8 Field Quality Control
A project plan should have been developed to ensure that all data are accurate and that decisions
based on these data are technically sound and defensible. The implementation of a project plan
requires QC procedures. QC procedures, therefore, represent specific tools for measuring the
degree to which quality assurance objectives are met. Field QC measures are discussed
comprehensively in ASTM D5283.
While some types of QC samples are used to assess analytical process, field QC samples are used
to assess the actual sampling process. The type and frequency of these field QC samples must be
specified by the project planning process along with being included in the project planning
documents and identified in the sampling plan. Definitions for certain types of field QC samples
can be found in ASTM D5283 and MARSSIM (2000).
10.2.9 Decontamination of Field Equipment
Sampling SOPs must describe the recommended procedure for cleaning field equipment before
and during the sample collection process, as well as any pretreatment of sample containers. The
SOPs should include the cleaning materials and solvents used, the purity of rinsing solution or
water, the order of washing and rinsing, associated personnel safety precautions, and the disposal
of cleaning agents.
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activity soils, soil gas, sludges, surface water, and ground water are given in ASTM D5608.
10.2.10 Packing and Shipping
The final responsibility of field sampling personnel is to prepare and package samples properly
for transport or shipment by a commercial carrier. All applicable state and federal shipping
requirements, discussed later in this section, must be followed. When samples must be shipped
by commercial carrier or the U.S. Postal Service, containers must be designed to protect samples
against crushing forces, impacts, and severe temperature fluctuations. Within each shipping
container, the cushioning material (sawdust, rubber, polystyrene, urethane foam, or material with
similar resiliency) should encase each sample completely. The cushioning between the samples
and walls of the shipping containers should have a minimum thickness of 2.5 cm. A minimum
thickness of five centimeters should be provided on the container floor.
Samples should also be protected from the potentially adverse impacts of temperature fluctuations. When appropriate, protection from freezing, thawing, sublimation, evaporation, or extreme
temperature variation may require that the entire interior surface of the shipping container be
lined with an adequate layer of insulation. In many instances, the insulating material also may
serve as the cushioning material.
The requirements for container security, cushioning, and insulation apply regardless of container
material. For smaller volume and low-weight samples, properly lined containers constructed
from laminated fiberboard, plastic, or reinforced cardboard outer walls also may be used.
When samples are shipped as liquids in glass or other breakable sample containers, additional
packaging precautions may have to be taken. Additional protection is obtained when sample
containers are shipped in nested containers, in which several smaller containers (i.e., inner
containers) are packed inside a second larger container (i.e., the outer pack or overpack). To
contain any spills of sample material within the shipping container, it is advisable either to wrap
individual samples or to line the shipping container with absorbent material, such as asbestosfree vermiculite or pearlite.
For proper packaging of liquid samples, additional guidance has been given by EPA (1987) and
includes the following:
• All sample bottles are taped closed;
• Each sample bottle is placed in a plastic bag and the bag is sealed;
• Each sample bottle may be placed in a separate metal can filled with vermiculite or other
packing material, and the lid taped to the can;
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• The cans are placed upright in a cooler that has its drain plug taped closed, inside and out,
and lined with a plastic bag; and
• The cooler is filled with packing material—“bubble wrap” or cardboard separators may be
used—and closed with sealing tape.
Field screening measurements are made for compliance with U.S. Department of Transportation
regulations, 49 CFR Parts 170 through 189, as well as compliance with the laboratory’s license
from the U.S. Nuclear Regulatory Commission (NRC; 10 CFR Part 71) and Agreement State (if
applicable). International requirements may also apply. See the International Air Transport
Association’s Dangerous Goods Regulations for additional guidance. These regulations not only
set contamination and radiation levels for shipping containers, but also describe the types of
containers and associated materials that are to be used based on the total activity and quantity of
materials shipped. When the samples are screened in the field with survey instrumentation, the
results should be provided to the laboratory. This information should also state the distance used
from the probe to the packing container wall. Measurements normally are made in contact or at
one meter. The readings in contact are most appropriate for laboratory use. The screening
measurements in the field are mainly for compliance with transportation requirements and are
usually in units of exposure. Laboratory license requirements are usually by isotope and activity.
Project planning and communication are essential to ensure that a specific set of samples can be
transported, received, and analyzed safely while complying with applicable rules and regulations.
The external surface of each shipping container must be labeled clearly, contain information
regarding the sender and receiver, and should include the respective name and telephone number
of a contact. When required, proper handling instructions and precautions should be clearly
marked on shipping containers. Copies of instructions, shipping manifest or container inventory,
chain of custody, and any other paperwork that are enclosed within a shipping container should
be safeguarded by placing documents within a sealed protected envelope.
10.2.11 Worker Health and Safety Plan
In some cases, field samples will be collected where hazardous agents or site conditions might
pose health and safety considerations for field personnel. These can include chemical, biological,
and radiological agents, as well as common industrial hazards associated with machinery, noise
levels, and heat stress. The health and safety plan established in the planning process should be
followed. For the U.S. Department of Defense, these plans may include imminent threats to life,
such as unexploded ordnance, land mines, hostile forces, chemical agents, etc. A few of the
hazards particular to field sampling are discussed in the following sections, but these should not
be construed as a comprehensive occupational health and safety program. The Occupational
Safety and Health Administration’s (OSHA) regulations governing laboratory chemical hygiene
plans are located at 29 CFR 1910.1450. These requirements should apply as well to field
sampling.
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10.2.11.1 Physical Hazards
MECHANICAL EQUIPMENT
Personnel working with hand-held tools (e.g., sledge hammers used for near-surface coring) or
power tools and equipment are subject to a variety of hazards. For example, personnel drilling
monitoring wells are exposed to a variety of potential mechanical hazards, including moving
machinery, high-pressure lines (e.g., hydraulic lines), falling objects, drilling through underground utilities, flying machinery parts, and unsafe walking and working surfaces. The
consequences of accidents involving these physical hazards can range from minor to fatal injury.
At a minimum, workers should be required to wear protective clothing, which includes hard hats,
gloves, safety glasses, coveralls (as an option) and steel-toed safety shoes. Workers required to
climb (e.g., ladders, drilling masts) must be trained according to OSHA standards in the proper
use of devices to prevent falls.
For sampling operations that require drilling, open boreholes and wells must be covered or
secured when unattended, including during crew breaks.
ELECTRICAL HAZARDS
Electric power often is supplied by gasoline or diesel engine generators. Working conditions may
be wet, and electrical shock with possibly fatal consequences may occur. In addition, drilling
operations may encounter overhead or buried electrical utilities, potentially resulting in exposure
to very high voltages, which could be fatal or initiate fires.
All electrical systems used during field operations should be checked for proper grounding
during the initial installation. Temporary electrical power provided to the drill site shall be
protected by ground-fault circuit interrupters.
NOISE HAZARDS
Power equipment is capable of producing sound levels in excess of 85 dB(A), the eight-hour
threshold limit value recommended by the American Conference of Governmental Industrial
Hygienists. Exposure to noise levels in excess of 85 dB(A) for long periods of time can cause
irreversible hearing loss. If noise levels exceed
CAUTION
85dB(A), a controlled area must be maintained
NOISE HAZARD
at this distance with a posting at each entrance
Hearing
Protection
Required Beyond This Point
to the controlled area to read:

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HEAT STRESS
The use of protective clothing during summer months significantly increases the potential for
personnel to experience heat stress. Adverse effects from heat stress include heat cramps,
dehydration, skin rash, heat edema, heat exhaustion, heat stroke, or death. When heat stress
conditions exist, the following ought to be available:
•
•
•
•

A cool and shaded rest area;
Regular rest breaks;
An adequate supply of drinking water; and
Cotton coveralls rather than impermeable Tyvek® coveralls.

CHEMICAL AND RADIOLOGICAL HAZARDS
The health and safety plan should contain information about a site’s potential radionuclides and
hazards that might be encountered during implementation of field sampling and survey
procedures. All field personnel should read the health and safety plan and acknowledge an
understanding of the radiological hazards associated with a site. Site specific training must be
provided that addresses the chemical and radiological hazards likely to be associated with a site.
Field procedures should include either information relating to these hazards or should reference
appropriate sections of the health and safety plan. References related to the use of protective
clothing are given in EPA (1987), DOE (1987, Appendix J), and in 29 CFR 1910, Subpart I.
When procuring environmental solid and liquid samples, unusual characteristics such as color,
suspended material, or number of phases and unusual odors should be noted and a description
should be provided to the on-site safety officer as well as the analytical laboratory. Additional
information concerning field methods for rapid screening of hazardous materials is presented in
EPA (1987). This source primarily addresses the appearance and presence of organic compounds
that might be present on occasions when one is collecting materials to detect radioactivity.
Checking samples for chemical or radiological hazards can be as simple as visual inspection or
using a hand-held radiation meter to detect radiation levels. Adjustments to laboratory procedures, particularly those involving sample handling and preparation, can only be made when
pertinent field information is recorded and relayed to the project planner and to the laboratory. In
some cases, a laboratory might not have clearance to receive certain types of samples (such as
explosives or chemical agents) because of their content, and it will be necessary to divert these
samples to an alternate laboratory. It might be necessary to reduce the volume sampled in order
to meet shipping regulations if high concentrations of radioactivity are present in the samples. In
some cases, the activity of one radionuclide might be much higher than others in the same
sample. Adjustments made on the basis of the radionuclide of higher activity might result in
collection of too little of another radionuclide to provide adequate detection and thus prevent
identification of these radionuclides because of their relatively low minimum detectable
concentrations. These situations should be considered during planning and documented in the
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appropriate sampling plan document.
10.2.11.2 Biohazards
Precautions should be taken when handling unknown samples in the field. Some examples are
wearing gloves, coveralls or disposable garments, plastic booties, dust masks or other respiratory
protection. Some biohazards may be snakes, ticks, spiders, and rodents (Hanta virus). Prevention
of potential exposure is the goal of a safety program. The type of protective equipment in the
field should be discussed in the planning process and specified in the appropriate plan document.
Since there are many specifics that are site dependent, it is difficult to create a comprehensive
list. But the information is discussed to provide an awareness and starting point for additional
discussion.
PERSONNEL TRAINING AND QUALIFICATION
All field operations that could lead to injury for sample collectors should be performed by
personnel trained to documented procedures. When sampling is conducted in radiologically
controlled areas (RCAs) as defined in regulatory standards (i.e., 10 CFR 20, 10 CFR 835).
Formal training and qualification of field personnel may be required.
Training may require both classroom and practical applications in order to familiarize personnel
with the basic theory of radiation and radioactivity and the basic rules for minimizing external
exposures through time, distance, shielding, and avoidance of internal exposure (by complying
with rules regarding smoking, drinking, eating, and washing of hands). Other topics to cover
include common routes of exposure (e.g., inhalation, ingestion, skin contact); proper use of
equipment and the safe handling of samples; proper use of safety equipment such as protective
clothing, respirators, portable shielding, etc.
Guidance for the training and qualification of workers handling radioactive material has been
issued by the Nuclear Regulatory Commission (see appropriate NRC NUREGs and Regulatory
Guides on training of radiation workers), Department of Energy (1994a–d), and the Institute of
Nuclear Power Operations (INPO 88-010). These and other documents should be consulted for
the purpose of training and qualifying field personnel.
PERSONNEL MONITORING AND BIOASSAY SAMPLING
When conditions dictate the need for personnel monitoring, various methods are commonly
employed to assess external and internal exposure that might have resulted from the inhalation or
ingestion of a radionuclide.
Thermoluminescent dosimeters, film badges, or other personnel dosimeters may be used to
monitor and document a worker’s external exposures to the whole body or extremities. For
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internal exposures, assessment of dose may be based on: (1) air monitoring of the work area or
the worker’s breathing zone; (2) in vivo bioassay (whole-body counting); or (3) in vitro bioassays
that normally involve urinalysis but also may include fecal analysis and nasal smears. For in vitro
bioassays (i.e., urine or fecal), the standard method involves a 24-hour sample collection in a
sealable container. Samples may be kept under refrigeration until laboratory analysis can be
performed to retard bacterial action. (Bioassay sample collection is normally not performed in the
“field.”)
The following guidance documents may be used for personnel monitoring and the collection and
preservation of bioassay samples:
• ANSI/ANS HPS N13.30 (1996), Performance Criteria for Radiobioassay;
• ANSI/ANS HPS N13.14 (1994), Internal Dosimetry Programs for Tritium Exposure—
Minimum Requirements;
• ANSI/ANS HPS 13.22 (1995), Bioassay Programs for Uranium;
• ANSI/ANS HPS 13.42 (1997), Internal Dosimetry for Mixed Fission Activation Products;
• DOE Implementation Guide, Internal Dosimetry Program, G-10 CFR 835/C1—Rev. 1 Dec.
1994a;
• DOE Implementation Guide, External Dosimetry Program, G-10 CFR 835/C2—Rev. 1 Dec.
1994b;
• DOE Implementation Guide, Workplace Air Monitoring, G-10 CFR 835/E2-Rev. 1 Dec.
1994c;
• DOE Radiological Control Manual, DOE/EH-0256T, Rev. 1, 1994d;
• NRC Regulatory Guide 8.9, Acceptable Concepts, Models, Equations, and Assumptions for a
Bioassay Program (September 1993);
• NRC Regulatory Guide 8.11, Applications of Bioassay for Uranium (Revision 1, July 1993);
• NRC Regulatory Guide 8.20, Applications of Bioassay for 125I and 131I (June 1974);
• NRC Regulatory Guide 8.22, Bioassays at Uranium Mills (Revision 1, August 1988);
• NRC Regulatory Guide 8.26, Applications of Bioassay for Fission and Activation Products
(September 1980);
• NRC Regulatory Guide 8.32, Criteria for Establishing a Tritium Bioassay Program (July
1988);
• NCRP (1987), Use of Bioassay Procedures for Assessment of Internal Radionuclides
Deposition; and
• INPO (1988), Guidelines for Radiological Protection at Nuclear Power Stations.

Part B: Matrix-Specific Issues That Impact Field Sample Collection,
Processing, and Preservation
Field processing should be planned in advance so that all necessary materials are available during
field work. Preparing checklists of processing equipment, instruments, and expendable
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materials—exemplified in part by lists accompanying sampling procedures described by EPA
(1994)—helps this planning effort and serves to organize field methods. Field personnel who
communicate problems should prevent loss of time, effort, and improper sample collection, as
well as documents exactly what equipment, instruments, etc. were used.
The initial steps taken in the field frequently are critical to laboratory analysis performed hours,
days, or even weeks after a sample is obtained. Various sample preparation steps may be required
before samples are packaged and shipped for laboratory analysis. The need for sample processing
and preservation is commonly determined by the sample matrix, the DQOs of the analysis, the
nature of the radionuclide, and the analytical method.
The goal of sample preservation is to maintain the integrity of the sample between the time the
sample is collected and the time it is analyzed, thus assuring that the analysis is performed on a
sample representative of the matrix collected. Sample preservation should limit biological and
chemical actions that might alter the concentration or physical state of the radionuclide
constituents or analytes. For example, cations at very low concentrations can be lost from
solution (e.g., cesium can exchange with potassium in the glass container, and radionuclides can
be absorbed by algae or slime growths in samples or containers that remain in the field for
extended periods). Requirements for sample preservation should be determined during project
planning when analytical protocols are selected. Sample preservation in the field typically
follows or accompanies processing activities. Sample preservatives may be added to sample
collection containers before they are sent to the field.
This section provides matrix-specific guidance that focuses on the preparation and processing of
field samples. In order to assist project planners in developing a sampling plan, a limited
discussion is also provided that describes matrix-specific methods commonly employed for the
collection of field samples. Guidance is presented for only the most common materials or
environmental media, which are generically classified as liquids, solids, and air. In some
instances, a solid material to be analyzed involves particulate matter filtered from a liquid or air
suspension. Because filter media can affect analytical protocols, a separate discussion is provided
that addresses sample materials contained on filter materials, including surface contamination
associated with wipe samples.

10.3 Liquid Samples
Liquid samples typically are classified as aqueous, nonaqueous, or mixtures. Aqueous samples
requiring analysis are likely to represent surface water, ground water, drinking water,
precipitation, tanks and lagoons, and runoff. Nonaqueous liquids may include a variety of
solvents, oils and other organic liquids. Mixtures of liquids represent a combination of aqueous
and nonaqueous liquids or a solid suspended in either aqueous and nonaqueous liquids.
Standardized water sampling procedures are described in numerous documents (APHA, 1998;
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EPA, 1985; EPA, 1987; DOE, 1997; ASTM D3370). Important decisions include the choice of
instrument or tool used to obtain the sample, the sample container material, the need for sample
filtration, and the use of sample preservatives.
10.3.1 Liquid Sampling Methods
The effect of the sample collection process on the sample integrity needs to be understood and
managed. Two examples are dissolved gases and cross-contamination. It may be necessary to
minimize dissolved oxygen and carbon dioxide, which can cause some dissolved metals to
undergo reaction or precipitation.
Sampling is discussed in NAVSEA (1997) and USACE (1995). The latter reference has been
superseded, but the revision does not include sampling. The sampling references listed in
USACE (1995) are:
• U.S. Environmental Protection Agency (EPA). 1984. Characterization of Hazardous Waste
Sites—A Method Manual, Vol. II, Available Sampling Methods, Second Edition, EPA 600-484-076.
• U.S. Environmental Protection Agency (EPA). 1982. Handbook for Sampling and Sample
Preservation of Water and Wastewater, EPA 600-4-82-029.
• U.S. Environmental Protection Agency (EPA). 1986. Compendium of Methods for
Determination of Superfund Field Operation Methods, EPA 600-4-87/006.
• U.S. Environmental Protection Agency (EPA). 1987. A Compendium of Methods for
Determination of Superfund Field Operation Methods, EPA 540-P-87-001a, OSWER
Directive 9355.0-14.
• U.S. Department of the Interior (DOI). 1980. National Handbook of Recommended Methods
for Water for Water-Data Acquisition, Volume I and II.
10.3.2 Liquid Sample Preparation: Filtration
Filtration of a water sample may be a key analytical planning issue and is discussed in Section
3.4.3, “Filters and Wipes.” A decision needs to be made during project planning whether or not
to filter the sample in the field. Filtration of water or other liquids may be required to determine
contaminant concentrations in solubilized form, suspended particulates, or sediment. The method
of filtration will depend on the required sample volume, the amount and size of suspended
particulates, and the availability of portable equipment and resources (e.g., electricity).
The potential need to filter a water sample principally depends on the source of water and the
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objectives of the project investigation. If, for example, the intent is to assess human exposure
from ingestion of drinking water “at-the-spigot,” unfiltered tap water samples are likely to be
required. Conversely, filtration may be required for water taken from an unlined field monitor
well that is likely to contain significant amounts of particulate matter. These solids are of little
relevance but may interfere with radioanalytical protocols (e.g., sample absorption may occur
during gross alpha or beta counting where the analytical procedure involves the simple
evaporation of a water aliquant on a planchet).
For remote sampling sites, sample processing may be restricted to gravity filtration that requires a
minimum of equipment and resources. Drawing samples through filters by pressure or suction
that is created by syringe, vacuum pump, or aspiration are alternative options. If filter papers or
membranes capture materials that will be retained for analysis, they should be handled with clean
rubber or plastic gloves, forceps, or other instruments to prevent sample contamination.
Each federal agency may have unique guidance to determine the need and process for filtering
samples. One performance-based example is that of EPA, discussed in the next section. This
guidance applies to either the field or laboratory filtration.
10.3.2.1

Example of Guidance for Ground-Water Sample Filtration

After considering whether or not to filter ground-water samples when analyzing for metals, the
Environmental Engineering Committee of EPA’s Science Advisory Board (EPA, 1997)
recommended:
• Several factors could introduce errors in the sampling and analysis of ground water for metals
or metallic radionuclides. Well construction, development, sampling, and field filtering are
among the steps that could influence the metals measured in the ground-water samples. Field
filtering is often a smaller source of variability and bias compared to these other factors.
Therefore, the Agency should emphasize in its guidance the importance of proper well
construction, development, purging, and water pumping rates so that the field filtering
decisions can also be made accurately.
• Under ideal conditions, field-filtered ground-water samples should yield identical metals
concentrations when compared to unfiltered samples. However, under non-ideal conditions,
the sampling process may introduce geological materials into the sample and would require
field filtration. Under such conditions, filtering to remove the geological artifacts has the
potential of removing colloids (small particles that may have migrated as suspended materials
that are mobile in the aquifer). Available scientific evidence indicates that when wells have
been properly constructed, developed, and purged, and when the sample has been collected
without stirring or agitating the aquifer materials (turbidity less than 5 nephelometric
turbidity units, NTU), then field filtering should not be necessary. For Superfund site
assessments, the low-flow sampling technique without filtration is the preferred sampling
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approach for subsequent metal analysis when well construction, well maintenance, and
hydrogeological conditions such as flow rate allow. Under such conditions, the collected
samples should be representative of the dissolved and particulate metals that are mobile in
ground-water systems. The Agency’s proposal to rely on low flow sampling and unfiltered
samples is a conservative approach that favors false positives over false negatives.
• When the turbidity of the sample is high, the situation is different. In-line filtering provides
samples that retain their chemical integrity. Therefore, field filtering of properly collected
ground-water samples should be done when turbidity in the samples is higher than 5 NTU,
even after slow pumping has been utilized to obtain the sample.
They acknowledged, however, that differences in the way wells are installed, their packing
materials, and the techniques used to collect ground-water samples can lead to variability in
analytical results between wells and between individual samples. Filtering a sample can be a way
to remove suspended particles and some colloids that contain metals that would not normally be
in the ground water if the material were not disturbed during sampling. Here, a colloid is defined
as a particle that ranges in size from 0.003 to 10 µm (Puls et al., 1990; Puls and Powell, 1992).
The literature indicates that colloids as large as 2 µm can be mobile in porous media (Puls and
Powell, 1992). Saar (1997) presents a review of the industry practice of filtration of ground-water
samples. For some sites with low hydraulic conductivity the presence of an excess of colloids
presents numerous monitoring challenges and field filtration might be necessary.
The desire to disturb the aquifer as little as possible has led to the use of low-flow sampling of
wells—low-flow purging and sampling occurs typically at 0.1 to 0.3 L/min (Saar, 1997). The
low-flow technique maximizes representativeness by (EPA, 1997):
• Minimizing disturbances that might suspend geochemical materials that are not usually
mobile;
• Minimizing disturbances that might expose new reactive sites that could result in leaching or
adsorption of inorganic constituents of ground water;
• Minimizing exposure of the ground water to the atmosphere or negative pressures, ensuring
that the rate of purging and sampling does not remove ground water from the well at a rate
much greater than the natural ground-water influx; and
• Monitoring indicator parameters to identify when stagnant waters have been purged and the
optimum time for sample collection.
In summary, based on the ability of the low-flow sampling technique to collect representative
samples, EPA suggests that filtering of ground-water samples prior to metals analysis is usually
not required (EPA, 1997).
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10.3.2.2

Filters

The removal of suspended particles is commonly achieved by filtration. When filtration is
required, it should be done in the field or as soon as practicable. Field filtration permits acid
preservatives to be added soon after collection, which minimizes the adsorption of soluble
contaminants on the container walls and avoids the dissolution of particulate matter which may
not be part of the sample to be analyzed.
An arbitrary size of 0.45 µm has gained acceptance as the boundary between soluble and
insoluble matter (particularly for water in power plant boilers (ASTM D6301). It is the filter pore
size that is commonly recommended by laboratory protocols. Material that may be present in
colloidal form (a second phase in a liquid that is not in solution), can have particles that range
from 0.001 to 2 µm. Such particles may be problematic since they may or may not be filterable
(Maron and Lando, 1974). Thus, there can be no single standard for filter type or pore size, and
every project should establish its own filtration protocol based upon its needs.
The fact that small particles pass through membrane filters has been recognized for some time
(Kennedy et al., 1974). Conversely, as the filters clog, particles an order of magnitude smaller are
retained by these filters (Sheldon and Sutcliffe, 1969). It should be noted, however, that
manufacturers of filters usually specify only what will not pass through the filter; they make no
claims concerning what actually does pass through the filter. Laxen and Chandler (1982) present
a comprehensive discussion of some effects of different filter types. They refer to thin (5 to 10
µm) polycarbonate filters as “screen types,” and thick (100 to 150 µm) cellulose nitrate and
acetate filters as “depth type.” The screen-type filters (e.g., polycarbonate) clog much more
rapidly than the depth type (e.g., cellulose nitrate and acetate) filters. Once the filtration rate
drops, particles that would normally pass through the filter are trapped in the material already
retained. Also, filtering through screen-type filters may take considerable time and may require
suction or pressure to accomplish in a reasonable time. Hence, the use of screen-type filters,
because of their increased propensity to clog, generally is not recommended.
In addition to the difficulty of contending with clogging, Silva and Yee (1982) report adsorption
of dissolved radionuclides on membrane filters. Although these drawbacks cannot be completely
overcome, they are still less than the potential difficulties that arise from not filtering.
Finally, good laboratory practices must be used for field sampling. The most likely sources of
contamination for the filters are improperly cleaned tubing and filter holders and handling the
filters with contaminated fingers. Tubing and holders should be thoroughly cleaned and rinsed
between samples and the entire system should be rinsed several times with the water to be
sampled. Filters should be handled with clean rubber gloves.

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10.3.3 Field Preservation of Liquid Samples
Sample degradation may occur between the time of collection and analysis due to microbial
contaminants or chemical interactions. Although sample degradation cannot destroy or alter the
radiological properties of a contaminant, it can alter the radionuclide’s chemical properties and
its potential distribution within a sample. For example, microbial processes are known to affect
both the chemical state and the distribution of radioelements due to oxidation-reduction
reactions, complexation and solubilization by metabolic compounds, bioaccumulation,
biomylation, and production of gaseous substances such as CO 2, H 2, CH 4, and H2S (Francis,
1985; Pignolet et al., 1989).
The selected field preservation method also should take into account compatibility with the
radionuclides, analytical methods, analytical requirements, and container properties (see Section
10.2.3, “Selection of Sample Containers”). One example that illustrates compatibility with the
analytical method is the addition of HCl to water samples as a preservative for gross alpha and
gross beta analyses. The HCl will corrode stainless steel planchets used in the method. If
laboratory personnel are aware of this, they can include steps to prevent the corrosion. Other
preservation issues for liquid samples are discussed in Table 10.1 (page 10-25). Compatibility
issues should be evaluated during the planning phase and included in the field sampling plan.
10.3.3.1

Sample Acidification

Acidification is the method of choice for preserving most types of water samples. The principal
benefit of acidification is that it keeps many radionuclides in solution and minimizes their
potential for removal by chemical and physical adsorption or by ion exchange. The mode by
which a radionuclide is potentially removed from solution is strongly affected by the radionuclide
and the container material. For example, studies conducted by Bernabee et al. (1980) and Milkey
(1954) demonstrated that the removal of metal ions from solution is dominated by physical (i.e.,
van der Waals) adsorption. Milkey’s conclusion is based on: (1) the observation that the loss of
uranium, lead, and thorium ions from solution was significantly greater for containers made of
polyethylene than of borosilicate glass; and (2) the fact that while adsorption by glass may
potentially involve all three adsorption processes; with polyethylene plastic, there are no valencetype attractive forces or ions to exchange, and only physical van der Waals adsorption is possible.
Similar observations were reported by: (1) Dyck (1968), who compared long-term adsorption of
silver ions by molded plastic to glass containers; (2) Jackson (1962), who showed that
polyethylene containers absorbed about five times as much 90Sr as glass containers at pH of about
seven; and (3) Martin and Hylko (1987a; 1987b), who reported that greater than 50 percent of
99
Tc was adsorbed by polyethylene containers from non-acidified samples.
For sample acidification, either nitric or hydrochloric acid is commonly added until a pH of less
than two (APHA, 1998, Table 7010.1; EPA, 1980, Method 900.0). Other guidance for sample
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preservation by acidification is summarized below.
In instances of very low-activity samples where container adsorption poses a significant concern,
but where acidification of the sample interferes with the radioanalytical method, the choice of
sample container may be limited to glass or require alternative methods. For example, the use of
acids as a preservative is not recommended for the analysis of tritium (3H), carbon-14 (14C), or
radon in water, and precautions must be taken for the following reasons:
• For radon, sample preservation offers no benefit and is therefore not required for analytical
accuracy. Adding acid also may cause the generation of CO2 in the sample, which could
purge radon gas.
• The addition of acid to a sample containing 14C may result in the production of 14CO 2 and the
loss of 14C from the sample.
• Acid does not have a direct effect on tritium. However, it may affect the cocktail used in
liquid scintillation analysis, or as with HCl, may add significant quench to the cocktail (see
Section 15.5.3, “Liquid Scintillation”).
Although acidification has been shown to effectively reduce the adsorption of technetium by
polyethylene, technetium in the TcO 4!4 state has been observed to volatilize in strong acid
solutions during evaporation while preparing water samples for gross beta analysis (NAS, 1960).
To hasten evaporation, the planchet is commonly flamed. This dilemma can be resolved by either
precoating planchets with a film of detergent prior to the addition of the acidified water sample
or by passive evaporation of the acidified water sample that avoids the higher temperature
associated with flaming (Blanchard et al., 1993).
10.3.3.2

Non-Acid Preservation Techniques

If a sample contains significant organics, or if contaminants under investigation react with acids
that interfere with the radioanalytical methods, other methods of sample preparation should be
considered.
REFRIGERATION AND FREEZING
The effect of refrigeration or freezing temperatures to arrest microbial activity is a fundamental
concept. Temperatures near the freezing mark or below not only retard or block bacterial growth
but arrest essentially all other metabolic activity. It should, however, be noted that most bacteria
can survive even in extreme temperatures. (Indeed, if a suspension of bacterial cells is frozen
rapidly with no appreciable formation of ice crystals, it can be kept at temperatures as low as
-194 EC for indefinite periods of time with little loss of viability.)
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The choice between refrigeration and freezing is dictated by the potential impacts of ice
formation on sample constituents. Besides physical changes of organic constituents, the initial
formation of ice crystals and the exclusion of any solutes may concentrate the solutes to the point
of precipitation. Quick freezing methods that minimize ice crystal formation are beneficial for
preserving some organic constituents. Quick freezing is commonly done by packing sealed
samples in liquid nitrogen or dry ice. Care must be taken, however, to avoid container breakage
due to sample volume expansion. An air space of a least 10 percent and a container made of
plastic provide reasonable assurance for container integrity.
When refrigeration is employed, attempts should be made to avoid temperatures that could result
in slow freezing and the formation of ice crystals. Optimum refrigeration temperatures for sample
preservation at 4 ± 2 EC can be achieved by packing samples in ice or freeze packs within a
thermally insulated leak-proof container (ASTM D3856; ASTM D3370).
PAPER PULP
The addition of paper pulp, with its adsorptive property and large surface area, can avoid the
adsorption and loss of easily hydrolyzed radionuclides to the container wall over time (Bernabee
et al., 1980). About two grams of finely ground paper pulp are added per liter of acidified sample
at time of collection. The pH should be adjusted to one or less and vigorously shaken. The
sample may be stored in this condition for an extended period of time. To prepare for analysis,
the pulp is removed from solution by filtration and subjected to wet ashing using strong acids
(Chapter 12, Laboratory Sample Preparation). This ashed solution is commonly added to the
original filtrate to make a reconstituted sample solution.
The use of paper pulp and the need for wet ashing, however, pose problems for certain
radioanalytical laboratory protocols and must therefore be thoroughly evaluated.
SULFITE
To prevent the loss of radioiodine from solution, sodium bisulfite (NaHSO3), sodium thiosulfate
(Na2S2O3), or sodium metabisulfite (Na2S2O5) may be used. These compounds are strong
reducing agents and will convert volatile iodine (I2) to nonvolatile iodine (I-). If acid is also
employed to preserve samples for analysis of other radionuclides, it is important to note that acid
will counteract the effectiveness of the reducing agent. For this reason, samples collected for
iodine analyses typically are collected and preserved in a separate container. It should also be
noted that the reducing environment produced by the sulfite-type preservatives may convert iron,
uranium, and other reducible ions or their compounds to a different oxidation state. The
inadvertent change in oxidation state of other radionuclides will have an obvious adverse impact
on radioanalytical measurements that require chemical separation. Section 14.9 has additional
information on carriers and tracers.
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OTHERS
Other methods that have been used to preserve liquid samples containing organics and biological
materials include chemical preservatives (e.g., formaldehyde and methanol). Table 10.1
summarizes the advantages and disadvantages of these and previously described preservation
methods.
TABLE 10.1—Summary of sample preservation techniques.
Preservation Technique

Advantages

Disadvantages

Addition of HNO3

Reduces pH and inhibits plating of
metals on container walls.

Addition of HCl

Reduces pH and inhibits plating of
metals on container walls.
Chloride forms strong anionic
complexes with Iron and Uranium.

Addition of Sulfite

Forms a reducing environment to
prevent the volatilization of iodine.
Addition of
Preserves organic samples.
Formaldehyde
Prevents further biological activity.
Cooling
Preserves organic samples (i.e.,
water, foods).
(Ice at approximately 0
EC)
Reduces dehydration and retains
moisture.
Reduces biological activity.
Preserves organic samples (i.e.,
Freezing
(Dry Ice at approximately water, plant, animal).
-78 EC)
Suspends biological activity.
Addition of Paper Pulp
Provides large surface area for
adsorption of metals, thus minimizing adsorption on container walls.

Strong oxidizer that might react with organic
compounds, such as liquid scintillation
cocktails.
14
C might be lost as 14CO2.
Causes quench in liquid scintillation cocktails.
14
C might be lost as 14CO2.
Might cause corrosion of stainless steel
planchets on gross analyses.
May produce undesirable oxidation states of
iron or uranium.
May create disposal problems.
Ice melts, requiring replacement over time.

Dry ice sublimates and requires replacement.
May crack sample container if frozen too
quickly.
Requires pH to be one or less.
Requires filtration and wet ashing of paper pulp
and combining liquids to make a new solution.

10.3.4 Liquid Samples: Special Cases
In some cases, liquid samples require special handling in order to preserve or retain a volatile or
gaseous radionuclide. The following are examples of specific methods used to recover or
preserve such samples of interest.
10.3.4.1

Radon-222 in Water

Waterborne radon is analyzed most commonly by liquid scintillation methods, although gammaray spectrometry and other methods have been employed or proposed. Liquid scintillation has the
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obvious advantage of being designed for automated sample processing and is, therefore, less
labor intensive or costly. A key to consistency in analytical results is the zero headspace sampling
protocol such as the one described below.
Since radon is inert and nonpolar, it diffuses through plastic more rapidly than glass. The use of
plastic scintillation vials, therefore, leads to significant loss of radon in water (Whittaker, 1989;
Hess and Beasley, 1990). For this reason, it is recommended that the water sample is collected in
a 23 mL glass scintillation vial, capped with a Teflon™ or foil-lined cap.
Samples are collected from a nonaerated faucet or spigot, which has been allowed to flow for
sufficient time so that the sample is representative of the water in the distribution system or well.
The time will vary depending on the source.
10.3.4.1

Milk

Milk commonly is viewed as the food product of greatest potential dose significance for airborne
releases of radionuclides. Due to the animals’ metabolic discrimination, however, only a few
radionuclides have a significant dose impact via the milk pathway, notably 90Sr, 131I, and 137Cs.
To prevent milk from souring or curdling, samples should be refrigerated. Preservation of milk
may also be achieved through the addition of formaldehyde or methanol (DOE, 1987),
methimazole (Harrington et al., 1980), or Thimerosal (EPA, 1994). Analytical procedures for
select radionuclides in milk are well established and should be considered when deciding on a
sample preservation method. Adding formaldehyde to milk samples may require them to be
disposed of as hazardous or mixed wastes.
Due to the volatility and potential loss of 131I (as I2), a known amount of NaI dissolved in water
may be added to the milk sample at time of collection if iodine analysis is required. The NaI not
only serves as a carrier for the chemical separation of radioiodine, but also provides a
quantitative tool for determining any loss prior to analysis (DOE, 1990).
10.3.5 Nonaqueous Liquids and Mixtures
Nonaqueous liquids and mixtures include a wide range of organic fluids or solvents, organic
materials dissolved in water, oils, lubricants, etc. These liquids are not likely to represent
contaminated environmental media or matrices, but most likely represent waste streams that must
be sampled. Nonaqueous waste streams are generated as part of normal operations by nuclear
utilities, medical facilities, academic and research facilities, state and federal agencies, radiopharmaceutical manufacturers, DOE weapons complexes, mining and fuel fabrication facilities,
etc. Examples of these nonaqueous liquids and mixtures include waste oils and other lubricants
that are generated routinely from maintenance of equipment associated with nuclear power plant
operations or the production of nuclear fuel and nuclear weapon components; and organic and
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inorganic solvents, acids, and bases that are used in a variety of medical, research, and industrial
applications.
In addition to the production of nonaqueous liquid wastes from routine operations by these
facilities, large quantities of nonaqueous liquids containing radionuclide contaminants are also
generated by routine facility decontamination efforts and final decontamination associated with
facility decommissioning. For decontamination and decommissioning activities, a wide range of
processes have been developed that employ halogenated organic compounds, such as Freon®,
chloroform, or trichloroethane. Other aggressive chemical decontamination processes involve
dissolution and removal of metal and oxide layers from surfaces using acid solutions (e.g.,
sulfuric acid, nitric acid, phosphoric acids, and oxalic acid). Chemical decontamination also may
use chelating agents in concentrated processes (5 to 25 percent by weight chemical in solution)
and dilute processes (one percent wt. or less chemicals in solution). Examples of chemical
processes that can be used in both concentrated and dilute forms include the low oxidation-state
transition-metal ion (LOMI) and LOMI-nitric permanganate, developed by Dow Chemical
Company and AP/Citron. The reagents used in both the concentrated and dilute processes include
chelating and complexing agents such as ethylene diamine tetraacetic acid (EDTA), diethylene
triamine pentaacetic acid (DTPA), citric acid, oxalic acid, picolinic acid, and formic acid.
Chelating agents and organic acids are used in decontamination formulas because they form
strong complexes with actinides, lanthanides, heavy metals, and transition metals and assist in
keeping these elements in solution.
Generally, these chemical decontamination solutions, once used, are treated with ion-exchange
resins to extract the soluble activity. The ion-exchange decontamination solutions must be
sampled, nevertheless, to assess the amount of residual radioactivity.
The radionuclides that may be encountered with nonaqueous liquids and mixtures depend on
both the nature of the liquid and its usage. The following listing of radionuclides and liquids are
based on published data collected by NRC (1992) and the State of Illinois (Klebe 1998; IDNS
1993-1997), but are not intended to represent a comprehensive list:
• Toluene/xylene/scintillation fluids used by research and clinical institutions: 3H, 14C, 32/33P,
35
S, 45Ca, 63Ni, 67Ga, 125/131I, 99Tc, 90Sr, 111In, 123/125I, 147Pm, 201/202Tl, 226/228Ra, 228/230/232Th,
232/234/235/238
U, 238/239/241/242Pu, 241Am.
• Waste oils and lubricants from operation of motors, pumps, and other equipment: 3H, 54Mn,
65
Zn, 60Co, 134/137Cs, 228/230/232Th.
• Halogenated organic and solvents from refrigeration, degreasing, and decontamination: 3H,
14
C, 32/33P, 35S, 54Mn, 58/60Co, 63Ni, 90Sr, 125/129I, 134/137Cs, 226/228Ra, 228/230/232Th, 232/234/235/238U,
238/239/241
Pu.
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• Other organic solvents from laboratory and industrial operations and cleaning: 3H, 32/33P, 35S,
45
Ca, 125I, U-natural.
• Inorganic and organic acids and bases from extraction processes and decontamination: 3H,
14
C, 32/33P, 35S, 54Mn, 67Ga, 125/131I, 60Co, 137Cs, and U-natural.
Due to the large number of potential nonaqueous liquids and the complex mixtures of radionuclide contaminants that may require radiochemical analysis, a comprehensive discussion of sample
preparation and preservation is beyond the scope of this discussion. In most instances, however,
these samples are not likely to require refrigeration or chemical preservatives that protect against
sample degradation.
Some organic solvents and highly acidic or basic liquids may react with plastic containers,
causing brittleness or breakage. In selecting sample containers for these nonaqueous samples, it
is important to assess the manufacturers product specifications, which typically provide
information regarding the container’s resistance to chemical and physical agents. When
nonaqueous samples are stored for long periods of time, containers should be checked routinely.

10.4 Solids
Solid samples consist of a wide variety of materials that include soil and sediment, plant and
animal tissue, metal, concrete, asphalt, trash, etc. In general, most solid samples do not require
preservation, but require specific processing in the field before transporting to the laboratory for
analysis. For example, soil sample field processing may require sieving in order to establish
sample homogeneity. These and other specific handling requirements are described below in the
section on each type of solid sample.
The most critical aspect is the collection of a sufficient amount of a representative sample. One
purpose of soil processing is to bring back only that sample needed for the laboratory. Unless
instructed otherwise, samples received by the laboratory are typically analyzed exactly as they are
received. This means that extraneous material should be removed at the time of sample
collection, if indicated in the appropriate plan document.
In many instances, sample moisture content at the time of collection is an important factor. Thus,
the weights of solid samples should be recorded at the time a sample is collected. This allows one
to track changes in wet weight from field to laboratory. Dry and ash weights generally are
determined at the laboratory.
Unlike liquid samples that may be introduced or removed from a container by simple pouring,
solid samples may require a container that is designed for easy sample placement and removal.
For this reason, large-mouth plastic containers with screw caps or individual boxes with sealable
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plastic liners are commonly used. The containers also minimize the risk for breakage and sample
cross-contamination.
10.4.1 Soils
ASTM D653 defines soil as: “Sediments or other unconsolidated accumulations of solid particles
produced by the physical and chemical degradation of rocks, and that might or might not contain
organic matter.” ASTM C999 provides generic guidance for soil sample preparation for the
determination of radionuclides. ASTM D4914 and D4943 provide additional information on soil
and rock.
The distribution of radionuclides in soil should be assumed to be heterogeneous. The degree of
heterogeneity is dictated by the radionuclide’s mode of entry into the environment and soil, the
chemical characteristics of the radionuclide contaminant, soil composition, meteorological and
environmental conditions, and land use. For example, soil contamination from an airborne
release of a radionuclide with strong affinity for clay or other mineral constituents of soil likely
will exhibit a gradient with rapidly diminishing concentrations as a function of soil depth (the
parameter associated with this affinity is KD, which is the concentration of the solid phase
divided by the concentration of the liquid phase). Moreover, contamination may be differentially
distributed among soil particles of different sizes. In most cases, because the contaminant is
adsorbed at the surface of soil particles and since the surface-to-volume ratio favors smaller
particles, smaller soil particles will exhibit a higher specific activity when compared to larger
particles. If land areas include areas of farming, tilling of soil will clearly impact the distribution
of surface contamination.
10.4.1.1

Soil Sample Preparation

Extraneous material should be removed at the time of sample collection, if indicated in the
appropriate plan document. The material may have to be saved and analyzed separately,
depending on the project requirements and MQOs. If rocks, debris, and roots are removed from a
soil sample after it arrives at the laboratory, there may be insufficient material to complete all the
requested analyses (see Section 12.3.1.1 “Exclusion of Material”). A sufficient amount of sample
should be collected to provide the net quantity necessary for the analysis. Subsequent drying at
the laboratory may remove a large percentage of the sample weight that is available for analysis.
Field-portable balances or scales may be used to weigh samples as they are collected, further
ensuring sufficient sample weights are obtained. For certain types of samples, the project DQOs
may require maintaining the configuration of the sample, such as core samples where
concentration verses depth will be analyzed.
The project plan should address the impact of heterogeneity of radionuclide distribution in soil.
Some factors to consider that may impact radionuclide distribution are: determining sampling
depth, the need for removal of vegetative matter, rocks, and debris, and the homogenation of soil
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particulates. For example, soil sampling of the top 5 cm is recommended for soils contaminated
by recent airborne releases (ASTM C998); soil depth to 15 cm may be appropriate when
exposure involves the need to monitor the root zone of food crops (MARSSIM, 2000; NRC,
1990). The need for sample field QC, such as splitting, should be evaluated. Some types of field
QC can be used to evaluate the extent of radionuclide homogeneity. In general, no special
preservation measures are required for soil samples; however, preliminary soil sample
preparation involving drying, sieving, homogenizing, and splitting may be performed by a field
laboratory prior to sample shipment to the analytical laboratory.
If volatile elements are suspected to be present with other nonvolatile contaminants, samples
must be split before drying to avoid loss of the contaminant of interest. Dried samples are
homogenized by mortar and pestle, jaw crusher, ball mill, parallel plate grinder, blender, or a
combination of these techniques and sieved to obtain a uniform sample. Sieve sizes from 35 to
200 mesh generally are recommended for wet chemistry procedures. ASTM C999 correlates
various mesh sizes with alternative designations, inclusive of physical dimensions expressed in
inches or in the metric system. In addition, samples for chemical separations are usually ashed in
a muffle furnace to remove any remaining organic materials that may interfere with the
procedures.
10.4.1.2

Sample Ashing

Soil samples that require chemical separation for radionuclide analysis may also be ashed by the
field laboratory. The use of the term “field laboratory” can cause confusion, since no single
definition is possible. It is used here to define a laboratory that is close to the point of sample
collection. It does not imply that there is a distinction in requirements or specifications that
impact quality. For soil samples, ashing is performed in a muffle furnace to remove any organic
materials that may interfere with radiochemical procedures.
10.4.2 Sediments
Sediments of lakes, reservoirs, cooling ponds, settling basins, and flowing bodies of surface
water may become contaminated as a result of direct liquid discharges, wet surface deposition, or
from runoffs associated with contaminated soils. Because of various chemically and physically
binding interactions with radionuclides, sediments serve as integrating media that are important
to environmental monitoring. An understanding of the behavior of radionuclides in the aquatic
environment is critical to designing a sampling plan, because their behavior dictates their
distribution and sampling locations.
In most cases, sediment is separated from water by simple decanting, but samples also may be
obtained by filtering a slurry or through passive evaporation. As noted previously, care must be
taken to avoid cross contamination from sampling by decontaminating or replacing tools and also
from avoiding contact between successive samples. Suitable sample containers include glass or
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plastic jars with screw caps. The presence of volatile or semi-volatile organic and microorganisms may impact the radionuclide concentration, therefore, samples should be kept on ice
while in the field and refrigerated while awaiting radioanalysis. Sediment cores may be sampled,
frozen, and then sectioned.
10.4.3 Other Solids
10.4.3.1

Structural Materials

In some cases, a project plan requires sample analysis of structural materials such as concrete or
steel. Concrete from floors, walls, sidewalks or road surfaces is typically collected by scabbling,
coring, drilling, or chiseling. Depending on the radionuclides of interest and detection methods,
these sample preparations may require crushing, pulverization, and sieving.
Metal associated with structures (e.g., I-beams, rebar) or machines may be contaminated on
exterior or interior surfaces or through activation may become volumetrically contaminated.
Surface contamination may be assessed by swipe samples that provide a measure of removable
contamination (Section 10.6) or by scraping, sandblasting, or other abrasive techniques.
Volumetric contamination is frequently assessed by nondestructive field measurements that rely
on gamma-emitting activation products. However, drill shavings or pieces cut by means of a
plasma arc torch may be collected for further analysis in a laboratory where they can be analyzed
in a low-background environment. In general, these materials require no preservation but, based
on activity/dose-rate levels and sample size and weight, may require proper shielding, engineered
packaging, and shipping by a licensed carrier.
10.4.3.2

Biota: Samples of Plant and Animal Products

The release of radionuclides to the environment from normal facility operations or as the result of
an accident requires the sampling of a wide variety of terrestrial and aquatic biota. For most
biota, sample preservation usually is achieved by icing samples in the field and refrigeration until
receipt by the analytical laboratory. The field sampling plan should describe the type of
processing and preservation required.
Foods may be categorized according to the U.S. Department of Agriculture scheme as leafy
vegetables, grains, tree-grown fruits, etc., and representative samples from each group may be
selected for analysis.
MEAT, PRODUCE, AND DAIRY PRODUCTS
Samples of meat, poultry, eggs, fresh produce, and other food should be placed in sealed plastic
bags and appropriately labeled and preserved by means of ice in the field and refrigeration during
interim storage prior to delivery to the analytical laboratory. All food samples may be reduced to
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edible portions (depending on study objective) for analysis in a manner similar to that for human
consumption (i.e., remove cores, bones, seeds, other nonedible parts) and weighed as received
from the field (i.e., wet weight) within 24 hours. Wet weights are desired, since consumption
data are generally on this basis.
ANIMAL FEED AND VEGETATION
Animal feeds also provide important data for determining radionuclide concentrations in the food
chain. Crops raised for animal feed and vegetation consumed by grazing farm animals may be
sampled. Depending upon radionuclides under investigation and their associated MQOs,
kilogram quantities of vegetative matter may be needed.
As in all terrestrial samples, naturally occurring 40K and the uranium and thorium series
radionuclides contribute to the radiation observed. Deposition of such cosmic-ray-produced
nuclides as 7Be and fallout from nuclear tests also may be present. Properly selected processed
items from commercial sources may be helpful in providing natural and anthropogenic
background data.
TERRESTRIAL WILDLIFE
Wild animals that are hunted and eaten may be of interest for potential dose estimates and
therefore may require sampling. Examples of wildlife that have been used are deer, rabbits, and
rodents that may feed or live in a contaminated site. An estimate of the radionuclide intake of the
animal just before its death may be provided by analyzing the stomach content, especially the
rumen in deer.
AQUATIC ENVIRONMENTAL SAMPLES
In addition to natural radionuclides and natural radionuclides enhanced by human activity, there
are numerous man-made radionuclides that have the potential for contaminating surface and
ground water. The most common of these are fission and activation products associated with
reactor operation and fuel cycle facilities. Radioanalysis of aquatic samples may therefore
include 54Mn, 58Co, 60Co, 65Zn, 95Zr, 90Sr, 134Cs, 137Cs, and transuranics, such as 239Pu.
When surface and ground waters are contaminated, radionuclides may be transferred through a
complex food web consisting of aquatic plants and animals. Aquatic plants and animals, as
discussed here, are any species which derive all or substantial portions of their nourishment from
the aquatic ecosystem, are part of the human food chain, and show significant accumulation of a
radionuclide relative to its concentration in water. Although fish, aquatic mammals, and
waterfowl provide a direct link to human exposure, lower members of the food chain also may be
sampled.
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FLORA
Aquatic biota such as algae, seaweed, and benthic organisms are indicators and concentrators of
radionuclides—especially 59Fe, 60Co, 65Zn, 90Sr, 137Cs, and the actinides—and can be vectors in
the water-fish-human food chain. As such, they may be sampled upstream and downstream at
locations similar to those described for sediment. Because of their high water content, several
kilograms (wet weight) should be collected per sample. The wet weight of the sample should be
recorded. Enough of the wet sample should be processed so that sufficient sample remains
following the drying process. Both algae (obtained by filtering water or by scraping submerged
substrates) and rooted aquatic plants should be sampled.
FISH AND SHELLFISH
Several kilograms of each fish sample are usually required; this may be one large fish, but
preferably a composite of a number of small ones. Analysis of the edible portions of food fish as
prepared for human consumption is of major interest. Fish may be de-boned, if specified in the
sampling plan. The whole fish is analyzed if it is used for the preparation of a fish meal for
consumption or if only trend indication is required. In a program where fish are the critical
pathway, fish are analyzed by species; if less detail is required, several species with similar
feeding habits (such as bottom feeders, insectivores, or predators) may be collected and the data
grouped. Some species of commercial fish, though purchased locally, may have been caught
elsewhere. Thus, the presence or absence of a radionuclide in a specific fish may not permit any
definite conclusion concerning the presence of the radionuclide in water at that location.
Shellfish, such as clams, oysters, and crabs, are collected for the same reasons as fish, but have
the advantage as indicators of being relatively stationary. Their restricted mobility contributes
substantially to the interpretation and application of analytical results to environmental
surveillance. Edible and inedible portions of these organisms can be prepared separately.
WATERFOWL
Waterfowl, such as ducks and geese, may also concentrate radionuclides from their food sources
in the aquatic environment and serve as important food sources to humans. The migratory
patterns and feeding habits of waterfowl vary widely. Some species are bottom feeders and, as
such, tend to concentrate those radionuclides associated with sediments such as 60Co, 65Zn, and
137
Cs. Others feed predominantly on surface plants, insects, or fish.
An important consideration in obtaining a sample from waterfowl is that their exterior surfaces,
especially feathers, may be contaminated. It is important to avoid contaminating the “flesh”
sample during handling. As with other biota samples, analyses may be limited to the edible
portions and should be reported on a wet weight basis.
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10.5 Air Sampling
The measurement of airborne radionuclides as gases or particulates provides a means of
evaluating internal exposure through the inhalation pathways. The types of airborne radioactivity
that may require air sampling are normally categorized as: (1) airborne particulates; (2) noble
gases; (3) volatilized halogens (principally radioiodines); and (4) tritiated water. Depending upon
the source term and the objectives of the investigation, air sampling may be conducted outdoors
as well as indoors on behalf of a variety of human receptors. For example, routine outdoor air
samples may be taken for large population groups living within a specified radius of a nuclear
facility. On the other end of the spectrum, air samples may be taken for a single person or small
group of persons exposed occupationally to a highly localized source of airborne radioactivity.
The purpose of the samples being collected must, therefore, be well defined in terms of sampling
location, field sampling equipment, and required sample volumes. Due to the wide range of
conditions that may mandate air sampling, and the limited scope of this section, only generic
topics of air sampling will be discussed.
10.5.1 Sampler Components and Operation
Common components of air sampling equipment include a sample collector (i.e., filter), a sample
collector holder, an air mover, and a flow-rate measuring device.
The sample holder should provide adequate structural support while not damaging the filter,
should prevent sampled air from bypassing the filter, should facilitate changing the filter, and
should facilitate decontamination. A backup support that produces negligible pressure drop
should be used behind the filter to prevent filter distortion or deterioration. If rubber gaskets are
used to seal the filter to the backing plate, the gasket should be in contact with the filter along the
entire circumference to ensure a good fit.
Air movers or vacuum systems should provide the required flow through the filter and minimize
air flow reduction due to filter loading. Consideration should be given to the use of air movers
that compensate for pressure drop. Other factors to consider should include size, power
consumption, noise, durability, and maintenance requirements.
Each air sampler should be equipped with a calibrated air-flow measuring device with specified
accuracy. To calculate the concentrations of any radionuclide in air collected, it is necessary to
determine the total volume of air sampled and the associated uncertainties. The planning
documents should state who is responsible for making volume corrections. Also, the information
needed for half-life corrections for short-lived radionuclides needs to be recorded. If the mean
flow during a collection period can be determined, the total volume of air sampled can be readily
calculated.
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Accurate flow measurements and the total integrated sample volume of air can be obtained using
a mass flow meter and a totalizer. This direct technique of air flow measurement becomes
impractical at remote field locations, due to cost and exposure of the flow meter to harsh
environments. Other procedures for the measurement of air flow in sampling systems are
reviewed by Lippmann (1989a). The sample parameters (flow rate, volume, associated
uncertainties, etc.) should be recorded by the sample collector.
The collection medium or filter used depends on the physical and chemical properties of the
materials to be collected and counted. A variety of particulate filters (cellulose, celluloseasbestos, glass fiber, membrane, polypropylene, etc.) is available. The type of filter is selected
according to needs, such as high collection efficiency, particle-size selectivity, retention of alpha
emitters on the filter surface, and the compatibility with radiochemical analysis. The criteria for
filter selection are good collection efficiency for submicron particles at the range of face
velocities used, high particle and mass loading capacity, low-flow resistance, low cost, high
mechanical strength, low-background activity, compressibility, low-ash content, solubility in
organic solvents, nonhygroscopicity, temperature stability, and availability in a variety of sizes
and in large quantities. The manufacturer’s specifications and literature should provide a source
for filter collection efficiency. In the selection of a filter material, a compromise must be made
among the above-cited criteria that best satisfies the sampling requirements. An excellent review
of air filter material used to monitor radioactivity was published by Lockhart and Anderson
(1964). Lippmann (1989b) also provides information on the selection of filter materials for
sampling aerosols by filtration. See ANSI HPS N13.1, Annex D and Table D.1, for criteria for
the selection of filters for sampling airborne radioactive particles.
In order to select a filter medium with adequate collection efficiency, it may be necessary to first
determine the distribution of size of airborne particulates. Several methods, including impactors
(e.g., multistage cascade impactor) and electrostatic precipitators, can be used to classify particle
size. Waite and Nees (1973) and Kotrappa et al. (1974) discuss techniques for particle sizing
based on the flow discharge perturbation method and the HASL cyclone, respectively. These
techniques are not recommended for routine environmental surveillance of airborne particulates,
although their use for special studies or for the evaluation of effluent releases should not be
overlooked. Specific data on various filter materials, especially retention efficiencies, have been
reported by several authors (Lockhart and Anderson, 1964; Denham, 1972; Stafford, 1973;
ASTM STP555) and additional information is available from manufacturers.
10.5.2 Filter Selection Based on Destructive Versus Nondestructive Analysis
Pure cellulose papers are useful for samples to be dissolved and analyzed radiochemically, but
the analytical filter papers used to filter solutions are inefficient collectors for aerosols and clog
easily. Cellulose-asbestos filter papers combine fairly high efficiency, high flow rates, high
mechanical strength, and low pressure drops when loaded. They are very useful for collecting
large samples but present difficulties in dissolution, and their manufacture is diminishing because
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of the asbestos. Fiberglass filters can function efficiently at high flow rates, but require fluoride
treatment for dissolution and generally contain sufficient radioactive nuclides to complicate lowactivity analysis. Polystyrene filters are efficient and capable of sustaining high air flow rates
without clogging. They are readily destroyed for analysis by ignition (300 EC) or by wet washing
with oxidizing agents, and also are soluble in many organic liquids. They have the disadvantage
of low mechanical and tensile strength, and they must be handled carefully. Membrane filters are
excellent for surface collection efficiency and can be used for direct alpha spectrometry on the
filter. However, they are fragile and suffer from environmental dust loading. An alternative
choice for radionuclides in the environment is the polypropylene fiber filter. Teflon ™ fiber filters
can be efficient, but they should be used with care because of their high ashing temperatures and
difficulties with digestion.
10.5.3 Sample Preservation and Storage
Since particulate air samples are generally dry samples that are chemically and physically stable,
they require no preservation. However, care must be exercised to avoid loss of sample from the
filter medium and the cross contamination among individual samples. Two common methods are
to fold filters symmetrically so that the two halves of the collection surface are in contact, or to
insert the filter into glassine envelopes. Filters should be stored in individual envelopes that have
been properly labeled. Filters may also be stored in special holders that attach on the filter’s edge
outside of the collection surface.
Since background levels of 222Rn and 220Rn progeny interfere with evaluating alpha air samples, a
holdup time of several hours to several days may be required before samples are counted.
Corrections or determinations can also be made for the contribution of radon or thoron progeny
present on a filter (Setter and Coats, 1961).
10.5.4 Special Cases: Collection of Gaseous and Volatile Air Contaminants
Prominent radionuclides that may exist in gaseous states include noble gases (e.g., 131/133Xe, 85Kr),
14
C as carbon dioxide or methane, 3H as water vapor, gaseous hydrogen, or combined in volatile
organic compounds and volatilized radioiodines.
10.5.4.1

Radioiodines

The monitoring of airborne iodine, such as 129I and 131I, may be complicated by the probable
existence of several species, including particulate iodine or iodine bound to foreign particles,
gaseous elemental iodine, and gaseous non-elemental compounds of iodine. A well-designed
sampling program should be capable of distinguishing all possible iodine forms. While it may
not always be necessary to differentiate between the various species, care should be taken so that
no bias can result by missing one or more of the possible species. See ANSI HPS N13.1 (Annex
C.3) for information on collection media for radioiodine.
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In addition to the problems noted above, charcoal cartridges (canisters) for the collection of
radioiodine in air are subject to channeling. Several should be mounted in series to prevent loss
of iodine. Too high a sampling rate reduces both the collection efficiency and retention time of
charcoal filters, especially for the non-elemental forms of iodine (Keller et al., 1973; Bellamy,
1974). The retention of iodine in charcoal is dependent not only on charcoal volume, but also the
length of the charcoal bed. Typical air flow rates for particulate sampling of 30 to 90 L/min (1 to
3 ft3/min) are normally acceptable for environmental concentrations of radioiodine. The method
proposed by the Intersociety Committee (APHA, 1972) for 131I concentrations in the atmosphere
involves collecting iodine in its solid and gaseous states with an “absolute” particulate filter in
series with an activated charcoal cartridge followed by gamma spectrometric analysis of the filter
and cartridge. The Intersociety-recommended charcoal cartridges are e inch (16 mm) diameter
by 1½ inch (38 mm) deep containing 3 g of 12-to-30-mesh KI-activated charcoal. The minimum
detectable level using the Intersociety method is 3.7×10-3 Bq/m3 (0.1 pCi/m3). Larger cartridges
will improve retention, permitting longer sampling periods. A more sensitive system has been
described by Baratta et al. (1968), in which concentrations as low as 0.037 Bq/m3 (0.01 pCi/mL)
of air are attainable.
For the short-lived radioiodines (mass numbers 132, 133, 135), environmental sampling is
complicated by the need to obtain a sufficient volume for analysis, while at the same time,
retrieving the sample soon enough to minimize decay (with half-lives ranging from two hours to
21 hours). Short-period (grab) sampling with charcoal cartridges is possible, with direct counting
of the charcoal as soon as possible for gamma emissions.
Because of the extremely long half-life and normally low environmental concentrations, 129I
determinations must usually be performed by neutron activation or mass spectrometry analysis
after chemical isolation of the iodine. For concentrations of about 0.11 Bq/L (3×10-10 µCi/mL),
liquid scintillation counting can be used after solvent extraction (Gabay et al., 1974).
10.5.4.2

Gases

Sampling for radioactive gases is either done by a grab sample that employs an evacuated
chamber or by airflow through a medium, such as charcoal, water, or a variety of chemical
absorbers. For example, radioactive CO2 is most commonly extracted by passing a known
volume of air through columns filled with 3 M NaOH solution. After the NaOH is neutralized
with sulfuric acid, the CO2 is precipitated in the form of BaCO3, which then can be analyzed in a
liquid scintillation counter (NCRP, 1985). An alternative method for collecting noble gases by
compression into high-pressure canisters is described in Section 15.3.5.1, “Radioactive Gases.”
Because noble gases have no metabolic significance, and concern is principally limited to
external exposure, surveillance for noble gases is commonly performed by ambient dose rate
measurements. However, the noble gases xenon and krypton may be extracted from air by
adsorption on activated charcoal (Scarpitta and Harley, 1990). However, depending upon the
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analytical method and instrumentation employed, significant interference may result from the
presence of naturally occurring radioactive gases of 222Rn and 220Rn.
10.5.4.3

Tritium Air Sampling

In air, tritium occurs primarily in two forms: as water vapor (HOT) and as hydrogen gas (HT).
However, if tritiated hydrogen (HT) is a suspected component of an air sample (e.g., from a vent
or stack), the sampling must take place in the emission point of the gas. This is because the high
escape velocity of hydrogen gas causes rapid, isotropic dispersion immediately beyond the
discharge point. Tritiated organic compounds in the vapor phase or attached to particulate matter
occur only occasionally. To measure tritium as HT or in tritiated organic, the gas phase can be
oxidized, converting the tritium to HOT before desiccation and counting. For dosimetric
purposes, the fraction present as HT can usually be neglected, since the relative dose for a given
activity concentration of HOT is 400 times that for HT (NCRP, 1978). However, if HT analysis
is required, it can be removed from the atmosphere by oxidation to water (HOT) using
CuO/MnO2 at 600 EC (Pelto et al., 1975), or with air passed over platinum alumina catalyst
(Bixel and Kershner 1974). These methods also oxidize volatile tritiated organic compounds to
yield tritiated water (ANSI HPS N13.1, Annex H).
A basic system for sampling HOT consists of a pump, a sample collector, and a flow-measuring
or flow-recording device. Air is drawn through the collector for a measured time period at a
monitored flow rate to determine the total volume of air sampled. The total amount of HOT
recovered from the collector is divided by the total volume of air sampled to determine the
average HOT-in-air concentration of the air sampled. In some sampler types, the specific activity
of the water collected is measured and the air concentration is determined from the known or
measured humidity. Some common collectors are cold traps, tritium-free water, and solid
desiccants, such as silica gel, DRIERITE™, or molecular sieve.
Cold traps are usually made of glass and consist of cooled collection traps through which sample
air flows. The trap is cooled well below the freezing point of water, usually with liquid nitrogen.
The water vapor collected is then prepared for analysis, usually by liquid scintillation counting.
Phillips and Easterly (1982) have shown that more than 95 percent HOT collection efficiency can
be obtained using a single cold trap. Often a pair of cold traps is used in series, resulting in a
collection efficiency in excess of 99 percent.
Gas-washing bottles (i.e., “bubblers”) filled with an appropriate collecting liquid (usually tritiumfree water) are used quite extensively for collecting HOT from air. HOT in the sample gas stream
“dissolves” in the collecting liquid. For the effective collection rate to remain the same as the
sample flow rate, the specific activity of the bubbler water must be negligible with respect to the
specific activity of the water vapor. Thus, the volume of air that can be sampled is ultimately
limited by the volume of water in the bubbler. However, except when sampling under conditions
of very high humidity, sample loss (dryout) from the bubbler usually limits collection time rather
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than the attainment of specific-activity equilibrium. Osborne (1973) carried out a thorough
theoretical and experimental evaluation of the HOT collection efficiency of water bubblers over a
wide range of conditions.
The use of silica gel as a desiccant to remove moisture from air is a common technique for
extracting HOT. The advantage of using silica gel is that lower HOT-in-air concentrations can be
measured, since the sample to be analyzed is not significantly diluted by an initial water volume,
which occurs when a liquid-sampling sink is used. Correcting for dilution is discussed in Rosson
et al. (2000).
10.5.4.4

Radon Sampling in Air

There are three isotopes of radon in nature: 222Rn is a member of the 238U decay chain; 220Rn is a
member of the 232Th decay chain; and 219Rn is a member of the 235U decay chain. Because of the
small relative abundance of the parent nuclides and the short half-lives of 220Rn (55 seconds) and
219
Rn (4 seconds), the term “radon” generally refers to the isotope 222Rn. Owing to its ubiquitous
presence in soils, uranium mill tailings, underground mines, etc., and the health risks to large
populations and occupational groups, radon is perhaps the most studied radionuclide.
Consequently, many reports and articles have been published in the scientific literature dealing
with the detection methods and health risks from radon exposures. Many of them appear in
publications issued by the EPA, DOE, NCRP, NAS, and in radiation-related journals, such as
Health Physics and Radiation Research. Given the voluminous amount of existing information,
only a brief overview of the sampling issues that impact laboratory measurements can be
presented here.
Quantitative measurements of radon gas and its short-lived decay products can be obtained by
several techniques that are broadly categorized as grab sampling, continuous radon monitoring,
and integrative sampling. Each method imposes unique requirements that should be followed
carefully. Continuous monitors are not discussed further, since they are less likely to be used by
laboratory analysts. Guidance for radon sample collection was published by EPA’s Radon
Proficiency Program, which was discontinued in October 1998 (EPA 1992; 1993). Additional
sampling methods and materials are also presented in EPA (1994) and Cohen (1989).
In general, EPA’s protocols specify that radon sampling and measurements be made under
standardized conditions when radon and its progeny are likely to be at their highest concentrations and maximum equilibrium. For indoor radon measurement, this implies minimum building
ventilation through restrictions on doors, windows, HVAC systems, etc. Also sampling should
not take place during radical changes in weather conditions. Both high winds and rapid changes
in barometric pressure can dramatically alter a building’s natural ventilation rate. Although
recommended measurements are likely to generate higher than actual average concentrations, the
benefit of a standardized sampling condition is that it is reproducible, least variable, and
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moderately conservative.
The choice among sampling methods depends on whether the measurement is intended as a
short-term, quick-screening measurement or as a long-term measurement that determines average
exposure or integration. In practice, the choice of a measurement system often is dictated by
availability. If alternative systems are available, the cost or duration of the measurement may
become the deciding factor. Each system has its own advantages and disadvantages, and the
investigator must exercise some judgment in selecting the system best suited to the objectives of
the investigation. Brief descriptions of several basic techniques used to sample air for radon and
its progeny are provided below.
GRAB SAMPLING
The term “grab sampling” refers to very short-term sampling. This method consists of evaluating
a small volume of air for either radon or radon decay product concentration. In the radon grab
sampling method, a sample of air is drawn into and subsequently sealed in a flask or cell that has
a zinc sulfide phosphor coating on its interior surfaces. One surface of the cell is fitted with a
window that is put in contact with a photomultiplier tube to count light pulses (scintillations)
caused by alpha disintegrations from the sample interacting with the zinc sulfide coating. The
general terms “flask” or “cell” are used in this discussion. Sometimes they are referred to as
“Lucas cells” (Lucas, 1982). The Lucas cell—or alpha scintillation counter—has specific
attributes, and not all radon cells are Lucas cells.
Several methods for performing such measurements have been developed. However, two
procedures that have been most widely used with good results are the Kusnetz procedure and the
modified Tsivogiou procedure. In brief, the Kusnetz procedure (Kusnetz, 1956; ANSI N13.8)
may be used to obtain results in working levels when the concentration of individual decay
products is not important. Decay products in up to 100 liters of air are collected on a filter in a
five-minute sampling period. The total alpha activity on the filter is counted any time between 40
and 90 minutes after sampling is completed. Counting can be done using a scintillation-type
counter to obtain gross alpha counts for a selected counting time. Counts from the filter are
converted to disintegrations using the appropriate counter efficiency. The disintegrations from
the decay products may be converted into working levels using the appropriate “Kusnetz factor”
for the counting time used.
The Tsivogiou procedure may be used to determine both working level and the concentration of
the individual radon decay products. Sampling is the same as in the Kusnetz procedure.
However, the filter is counted three separate times following collection. The filter is counted
between 2 and 5 minutes, 6 and 20 minutes, and 21 and 30 minutes after sampling is complete.
Count results are interpreted by a series of equations that calculate concentrations of the three
radon decay products and working levels.
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INTEGRATING SAMPLING DEVICES
By far, the most common technique for measuring radon is by means of integrating devices.
Integrating devices, like charcoal canister and the Electret-Passive Environmental Radon Monitor
(E-PERM®), are commonly employed as short-term integrating devices (two to seven days), while
alpha-track detectors are commonly used to provide measurements of average radon levels over
periods of weeks to months. Only charcoal canisters are discussed below, since they are more
likely to be used by laboratory analysts than electrets and alpha-track detectors.
CHARCOAL CANISTERS
Charcoal canisters are passive devices requiring no power to function. The passive nature of the
activated charcoal allows continual adsorption and desorption of radon. During the measurement
period, the adsorbed radon undergoes radioactive decay. Therefore, the technique does not
uniformly integrate radon concentrations during the exposure period. As with all devices that
store radon, the average concentration calculated using the mid-exposure time is subject to error
if the ambient radon concentration adsorbed during the first half of the sampling period is
substantially higher or lower than the average over the period. The ability of charcoal canisters to
concentrate noble gases or other materials may be affected by the presence of moisture,
temperature, or other gaseous or particulate materials that may foul the adsorption surface of the
charcoal.

10.6 Wipe Sampling for Assessing Surface Contamination
Surface contamination falls into two categories: fixed and loose. The wipe test (also referred to
as “swipes” or “smears”) is the universally accepted technique for detecting removable
radioactive contamination on surfaces (Section 12.5, “Wipe Samples”). It is often a stipulation of
radioactive materials licenses and is widely used by laboratory personnel to monitor their work
areas, especially for low-energy radionuclides that are otherwise difficult to detect with handheld survey instruments.” Frame and Abelquist (1999) provide a comprehensive history of using
smears for assessing removable contamination.
The purpose of the wipe test, organizational requirements or regulations, the nature of the
contamination, the surface characteristics, and the radionuclide all influence the conditions for
the actual wipe-test process. The wipe-test process should be standardized to ensure that the
sampling process is consistent. Since surfaces and wipe materials vary considerably, wipe-test
results provide qualitative indication of removable contamination. Fixed contamination will, by
definition, not be removed. Therefore, direct measurements may be necessary to determine the
extent on contamination.
The U.S. Nuclear Regulatory Commission (NRC, 1981) suggests that 100 cm2 areas be wiped
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and lists acceptable levels for surface contamination. However, NRC neither recommends the
collection device nor the manner in which to conduct such surveys, relying instead on
suggestions by the National Committee on Radiation Protection (1964) and the National Council
on Radiation Protection and Measurements (1978).
To maintain constant geometry in an automatic proportional counter, it is important that the wipe
remain flat during counting. Additionally, material that will curl can jam the automatic counter
and cause cross contamination or even destroy the instrument window. When it is necessary to do
destructive analysis on the wipe, it is critical that the wipe can easily be destroyed during the
sample preparation step, and that the residue not cause interference problems.
When wipes are put directly into liquid scintillation cocktail, it is important that the wipe not add
color or react with the cocktail. For maximum counting efficiency, as well as reproducibility, the
wipe either should dissolve in the cocktail or become transparent to the counting system.
10.6.1 Sample Collection Methods
10.6.1.1

Dry Wipes

Dry wipes (smears) for removable surface activity usually are obtained by wiping an area of 100
cm2 using a dry filter paper of medium hardness while applying moderate pressure. A 47 mm
diameter filter typically is used. This filter can be placed into a proportional counter for direct
counting. Smaller filters may be advantageous when the wipe is to be counted using liquid
scintillation counter for low energy beta-emitting radionuclides, such as tritium, 14C, and 63Ni.
The choice of wipe-test media and cocktail is critical when counting low-energy beta-emitting
radionuclides in liquid scintillation counters, because the liquid scintillation counting process
depends on the detection of light produced by the interaction of the radiation with the cocktail.
The filter may absorb energy from the radiation (see “Quench” under Section 15.5.3.3). A filter
that is in the cocktail can prevent light from being seen by both detectors at the same time. If
light is produced and seen by only one of the two detectors typical in liquid scintillation counting
systems, then the count will be rejected as noise. A filter/cocktail combination that produces a
sample that is transparent to the counting system is the best combination for liquid scintillation
counting. Background produced by the filter may also be a consideration.
For surveys of small penetrations, such as cracks or anchor-bolt holes, cotton swabs are used to
wipe the area of concern. The choice of material for wipe-testing for special applications is
critical (Hogue, 2002), and the material selected can significantly affect the efficiency of the
removal of surface radioactivity. Usually, switching wipe test material should be avoided during
a project, when possible. Samples (dry wipes or swabs) are placed into envelopes or other
individual containers to prevent cross-contamination while awaiting analysis. Dry wipes for
alpha and medium- or high-energy beta activity can be evaluated in the field by counting them on
an integrating scaler unit with appropriate detectors; the same detectors utilized for direct
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measurements may be used for this purpose. However, the more common practice is to return the
dry wipes to the laboratory, where analysis can be conducted using more sensitive techniques.
The most common method for analyzing wipe samples is to use a proportional counter. For very
low-energy beta emissions, wipe samples are commonly analyzed by liquid scintillation
counting.
Additional information on wipe-test counting can be found in ISO (7503-1; 7503-2; 7503-3),
which apply to surfaces of equipment and facilities, containers of radioactive materials, and
sealed sources. Abelquist (1998) discusses using smears to assess the quantity of removable
contamination as it applies to radiological surveys in support of decommissioning, compliance
with DOT shipping criteria, and operational radiological protection programs.
10.6.1.2

Wet Wipes

Although dry wipes are more convenient to handle, and there are fewer chances of cross
contamination, a general limitation of dry wipes is their low recovery of surface contamination.
The low recovery using dry wipes is due to the higher affinity for the surface by the contaminant
than for the filter paper. Several studies have shown that for maximum sensitivity, a wipe
material moistened with a suitable solvent may be indicated. For example, Ho and Shearer
(1992) found that alcohol-saturated swabs were 100 times more efficient at removing
radioactivity than dry swabs.
In another study, Kline et al. (1992) assessed the collection efficiency of wipes from various
surfaces that included vinyl floor tile, plate glass, and lead foil. Two different collection devices,
cotton swabs and 2.5 cm diameter glass fiber filter disks, were evaluated under various collection
conditions. Dry wipes were compared to collections made with the devices dampened with
different amounts of either distilled H2O, 70 percent ethanol, or a working-strength solution of a
multipurpose laboratory detergent known to be effective for removing contaminants from
laboratory glassware (Manske et al., 1990).
The entire area of each square was manually wiped in a circular, inwardly-moving motion with
consistent force. The collection capacity of each device was estimated by wiping progressively
larger areas (multiple grids) and comparing the measured amounts of radioactivity with the
amounts placed on the grids.
Collection efficiency varied with both the wipe method and the surface wipe. Contamination was
removed most readily from unwaxed floor tile and glass; lead foil released only about one-half
the radioactivity. Stainless steel, another common laboratory surface, has contamination retention
properties similar to those of glass.
In most cases, collection was enhanced by at least a factor of two after dampening either the
swabs or filter disks with water. Dampening with ethanol or the detergent produced removals that
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were statistically indistinguishable from samples dampened with an equal amount of water.
The filter disks had a higher collection capacity for removable contaminants than cotton swabs,
nearly doubling the radioactivity removed for each doubling of surface area wiped. Variability
within all methods was high, with coefficients of variation ranging from 2 to 30 percent.
For the moistened wipes, wipe efficiency depended on three factors, including the polarity of the
solvent, the polarity of the contaminant being measured, and the affinity of the compound for the
contaminated surface. For a solvent to readily dissolve a compound (i.e., remove it from the
surface), the solvent and the compound must have similar polarities. Nonpolar solvents include
ethyl acetate and petroleum ether; for polar solvents, water or methanol may be used (Campbell
et al., 1993). There are other factors that influence the affinity of a compound for a surface,
including porosity of the surface and available binding sites on the surface. One important factor
that influences binding capacity is the type of treatment that a surface has received. When
working with a surface treated with a nonpolar wax, such as that used on floor tile, a nonpolar
compound will be adsorbed to the surface, which further limits recovery. Recovery from
absorbent surfaces, such as laboratory bench paper or untreated wood, also may be poor due to
the porous nature of the surface.
10.6.2 Sample Handling
Filter paper or other materials used for wipe tests in the field should be placed in separate
containers that prevent cross contamination during transport and allow for labeling of each
sample. Plastic bags, paper or glassine envelopes, and disposable plastic petri dishes are typically
used to store and transport wipe samples. Field workers can use plastic or rubber gloves and
forceps when applying the wipe material to a surface and during handling as each wipe is placed
into a container. Protection of the sample wipe surface is the main concern when a wipe must be
placed in a container for transport. If a scintillation vial or planchet will be used in the laboratory,
then a field worker may put wipes directly into them. Planchets containing loose or self-sticking
wipes can also be put into self-sealing plastic bags to separate and protect the integrity of the
sample’s surface. Excessive dust and dirt can cause self adsorption or quenching, and therefore
should be minimized.
10.6.3 Analytical Considerations for Wipe Material Selection
Some analytical considerations for selecting wipe materials are included here, because field
sample collection and subsequent sample counting usually occur without such intervening steps
as sample preparation, sample dissolution, or separation. It is critical, therefore, to ensure that the
wipe material used for collection and the actual counting process are compatible. The following
paragraphs offer some general guidance for proportional and liquid scintillation counting. The
final paragraph discusses some key issues that impact dissolution of wipes.
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The wipe should remain flat during counting in order to maintain optimum counting geometry in
an automatic proportional counter. Wipe material that can curl may jam an automatic counter and
destroy the detector window of the counter, become a source of cross-contamination of samples,
or contaminate the counting system. Most proportional counting systems use two-inch (5 cm)
planchets, and the wipe should fit into the planchet. If not, a subsample will need to be taken, and
subsampling adds additional uncertainty due to sample homogeneity considerations.
When wipes are put directly into a liquid scintillation cocktail, the wipe should not add color or
react with the cocktail. For maximum counting efficiency and reproducibility, the wipe either
should dissolve or become transparent to the counting system. When wipes that have an adhesive
backing are put directly in a liquid scintillation cocktail, the adhesive may not dissolve
completely. Compatibility should be checked before use to prevent problems during actual
sample analysis. Special cocktails are available to dissolve filters, but they may cause a wastedisposal problem. Since the possible combination of cocktails and filters is large, only general
guidance is provided here. Consult the manufacturer’s specifications for specific guidance.
When it is necessary to do destructive analysis on a wipe, select a wipe that can be destroyed
easily or dissolved during the sample preparation steps, and the residue will not cause
interference problems in the subsequent counting. Some wipes have adhesive backing; the wipe
materials may dissolve easily but the adhesive backing may not. Additional steps would then be
necessary to destroy the adhesive backing. Dissolving glass-fiber wipes may require the use of
hydrofluoric acid. These extra processes can add time or cost to the analysis. See Section 10.5.2
(“Filter Selection Based on Destructive Versus Nondestructive Analysis”), Section 12.5 (“Wipe
Samples”) and Chapter 13 (Sample Dissolution) for additional information.

10.7 References
Abelquist, E.W. 1998. “Use of Smears for Assessing Removable Contamination,” Health
Physics Newsletter, Ops Center, July, pp. 18-19.
American National Standards Institute (ANSI) HPS N13.1. Sampling and Monitoring Releases of
Airborne Radioactive Substances from the Stacks and Ducts of Nuclear Facilities. 1999.
American National Standards Institute/American Nuclear Society (ANSI/ANS) HPS N13.14.
Internal Dosimetry Programs for Tritium Exposure - Minimum Requirements. 1994.
American National Standards Institute/American Nuclear Society (ANSI/ANS) HPS N13.22.
Bioassay Programs for Uranium. 1995.
American National Standards Institute/American Nuclear Society (ANSI/ANS) HPS N13.30.
Performance Criteria for Radiobioassay. 1996.
JULY 2004

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Field and Sampling Issues That Affect Laboratory Measurements
American National Standards Institute/American Nuclear Society (ANSI/ANS) HPS N13.42.
Internal Dosimetry for Mixed Fission Activation Products, 1997.
American National Standards Institute (ANSI). N13.8. American National Standard for
Radiation Protection in Uranium Mines. 1973.
American Public Health Association (APHA). 1972. Intersociety Committee for a Manual of
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American Public Health Association (APHA). 1998. Standard Methods for the Examination of
Water and Waste Water, 20th Edition. Washington, DC. Available at: www.standardmethods.
org.
American Society for Testing and Materials (ASTM) STP 555. Instrumentation for Monitoring
Air Quality, 1974. West Conshohocken, Pennsylvania.
American Society for Testing and Materials (ASTM) C998. Sampling Surface Soil for
Radionuclides, 1995. West Conshohocken, Pennsylvania.
American Society for Testing and Materials (ASTM) C999. Soil Sample Preparation for the
Determination of Radionuclides, 1995. West Conshohocken, Pennsylvania.
American Society for Testing and Materials (ASTM) D420. Site Characterization for
Engineering, Design, and Construction Purposes, 1998. West Conshohocken, Pennsylvania.
American Society for Testing and Materials (ASTM) D653. Terminology Relating to Soil, Rock,
and Contained Fluids, 1997. West Conshohocken, Pennsylvania.
American Society for Testing and Materials (ASTM) D3370, Standard Practices for Sampling
Water from Closed Conduits. ASTM, West Conshohocken, Pennsylvania.
American Society for Testing and Materials (ASTM) D3856. Good Laboratory Practices in
Laboratories Engaged in Sampling and Analysis of Water, 1995. West Conshohocken,
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American Society for Testing and Materials (ASTM) D3977. Determining Sediment
Concentration in Water Samples, 1997. West Conshohocken, Pennsylvania.
American Society for Testing and Materials (ASTM) D4840. Sampling Chain-of-Custody
Procedures, 1999. West Conshohocken, Pennsylvania.
MARLAP

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Field and Sampling Issues That Affect Laboratory Measurements
American Society for Testing and Materials (ASTM) D4914. Density of Soil and Rock in Place
by the Sand Replacement Method in a Test Pit, 1999. West Conshohocken, Pennsylvania.
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American Society for Testing and Materials (ASTM) D5245. Cleaning Laboratory Glassware,
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American Society for Testing and Materials (ASTM) D5283. Generation of Environmental Data
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American Society for Testing and Materials (ASTM) D5608. Decontamination of Field
Equipment Used at Low Level Radioactive Waste Sites, 1994. West Conshohocken,
Pennsylvania.
American Society for Testing and Materials (ASTM) D6301. Standard Practice for the
Collection of Samples of Filterable and Nonfilterable Matter in Water. West Conshohocken,
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Blanchard, R.L., R. Leiberman, W.S. Richardson III, and C.L. Wakamo. 1993. “Considerations
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Campbell, J.L., C.R. Santerre, P.C. Farina, and L.A. Muse. 1993. “Wipe Testing for Surface
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Cohen, B. S. 1989. “Sampling Airborne Radioactivity,” in Air Sampling Instruments for
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Department of Energy (DOE). 1994d. Radiological Control Manual. DOE/EH-0256T, Rev. 1.
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Dyck, W. 1968. “Adsorption of Silver on Borosilicate Glass,” Anal. Chem. 40:454-455.
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Radioactivity in Drinking Water. EPA-600/4-80-032, EPA, Environmental Monitoring and
Support Laboratory, Cincinnati, Ohio.
U.S. Environmental Protection Agency (EPA). 1982. Handbook for Sampling and Sample
Preservation of Water and Wastewater. EPA-600/4-82-029, EPA, Washington, DC. (PB83124503)
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U.S. Environmental Protection Agency (EPA). 1984. Characterization of Hazardous Waste
Sites–A Method Manual, Vol. II, Available Sampling Methods. EPA-600-4-84-076, Second
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U.S. Environmental Protection Agency (EPA). 1985. Sediment Sampling Quality Assurance
User’s Guide. EPA/600/4-85/048, Environmental Monitoring Systems Laboratory, Las
Vegas, NV. (PB85-233542).
U.S. Environmental Protection Agency (EPA). 1986. Compendium of Methods for Determination
of Superfund Field Operation Methods, EPA 600-4-87/006. Office of Emergency and
Remedial Response, Washington, DC.
U.S. Environmental Protection Agency (EPA). 1987. A Compendium of Superfund Field
Operations Methods. EPA/540/P-87/001. Office of Emergency and Remedial Response,
Washington, DC. (PB88-181557).
U.S. Environmental Protection Agency (EPA). 1989. Indoor Radon and Radon Decay Product
Measurement Protocols. Office of Air and Radiation, Washington, DC.
U.S. Environmental Protection Agency (EPA). 1992. Indoor Radon and Radon Decay Product
Measurement Device Protocols. EPA 402-R-92-004, EPA, Office of Air and Radiation,
Washington, DC. Available at www.epa.gov/iaq/radon/rpp_docs.htm.
U.S. Environmental Protection Agency (EPA). 1993. Protocols for Radon and Radon Decay
Product Measurements in Homes. EPA 402-R-92-003, EPA, Office of Air and Radiation,
Washington, DC. Available at www.epa.gov/iaq/radon/rpp_docs.htm.
U.S. Environmental Protection Agency (EPA). 1994. Routine Environmental Sampling
Procedures Manual For Radionuclides. EPA, Office of Radiation and Indoor Air and
National Air and Radiation Environmental Laboratory, Montgomery, AL.
U.S. Environmental Protection Agency (EPA). 1996. Radon Proficiency Program - Handbook.
EPA 402-R-95-013, EPA, Office of Radiation and Indoor Air, Washington, DC.
U.S. Environmental Protection Agency (EPA). 1997. To Filter or Not to Filter, That is the
Question. EPA Science Advisory Board (SAB), Environmental Engineering Committee,
Special Topics Subcommittee, July 11, 1997. EPA-SAB-EEC-LTR-97-011.
Frame, P.W. and E.W. Abelquist. 1999. Use of Smears for Assessing Removable Contamination.
Operational Radiation Safety, May, 76(5):S57-S66. Available at: www.hps1.org/sections/
rso/ophpinfo/papers.htm.
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Francis, A.J. 1985. Low-Level Radioactive Wastes in Subsurface Soils. Soil reclamation
Processes: Microbiological Analyses and Applications, NY.
Friend, A.G., A.H Story, C.R. Henderson, and K.A. Busch. 1965. Behavior of Certain
Radionuclides Released into Fresh-Water Environments. U.S. Public Health Service
Publication 999-RH-13.
Gabay, J.J., C.J. Paperiello, S. Goodyear, J.C. Daly, and J.M. Matuszek. 1974. “A Method for
Determining Iodine-129 in Milk and Water,” Health Physics 26, p. 89.
Harrington, C.L., R.A. Mellor, R.E. Lockwood, and K.G. Dagenais 1980. “Advantages and
Limitations of Chemical Preservatives for Use in the Radiological Analysis of I-131 in
Environmental Milk Samples, “ Health Physics 40:6, p. 907.
Hess, C.T. and S.M. Beasley. 1990. Setting Up a Laboratory for Radon in Water Measurments.
Radon, Radium and Uranium in Drinking Water, Lewis Publishers, Chelsea, MI.
Ho, S.Y. and D.R. Shearer. 1992. “Radioactive Contamination in Hospitals from Nuclear
Medicine Patients,” Health Physics 62, pp. 462-466.
Hogue, M.G. 2002. “Field Comparison of the Sampling Efficacy of Two Smear Media: Cotton
Fiber and Kraft Paper,” Operational Radiation Safety, 83:2, pp. S45-S47
Illinois Department of Nuclear Safety (IDNS). 1993. 1992 Annual Survey Report. Springfield,
Illinois.
Illinois Department of Nuclear Safety (IDNS). 1994. 1993 Annual Survey Report. Springfield,
Illinois.
Illinois Department of Nuclear Safety (IDNS). 1995. 1994 Annual Survey Report. Springfield,
Illinois.
Illinois Department of Nuclear Safety (IDNS). 1996. 1995 Annual Survey Report. Springfield,
Illinois.
Illinois Department of Nuclear Safety (IDNS). 1997. 1996 Annual Survey Report. Springfield,
Illinois.
Institute of Nuclear Power Operations (INPO). 1988. Guidelines for Radiological Protection at
Nuclear Power Stations. INPO 88-010, Atlanta, Georgia.
International Standards Organization (ISO) 7503-1. Evaluation of Surface Contamination – Part
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1: Beta-emitters (Maximum Beta Energy Greater than 0.15 MeV) and Alpha-Emitters. 1988,
Geneva, Switzerland.
International Standards Organization (ISO) 7503-2. Evaluation of Surface Contamination – Part
2: Tritium Surface Contamination. 1988, Geneva, Switzerland.
International Standards Organization (ISO) 7503-3. Evaluation of Surface Contamination – Part
3: Isomeric Transition and Electron Capture Emitters, Low Energy Beta-emitters (E βmax
<0.15 MeV). 1996, Geneva, Switzerland.
Jackson, E.W. 1962. “Prevention of Uptake of Strontium Ions on Glass,” Nature 194:672.
Johnson, B.H. 1980. A Review of Numerical reservoir Hydrodynamic Modeling. U.S. Army
Corps of Engineers, Waterways Experiment Station, Vicksburg, Mississippi.
Keller, J.H., T.R. Thomas, D.T. Pence, and W.J. Maeck. 1973. “An Evaluation of Materials and
Techniques Used for Monitoring Airborne Radioiodine Species,” in Proceedings of the 12th
AEC Air Cleaning Conference. U.S. Atomic Energy Commission, Washington, DC.
Kennedy, V.C., G.W. Zellweger, and B.F. Jones. 1974. “Filter Pore Size Effects on the Analysis
of Al, Fe, Mn, and Ti in Water,” Water Resources Research 10:4, pp. 785-790.
Klebe, M. 1998. Illinois Department of Nuclear Safety. Correspondence of June 12, 1998 to Mr.
J.C. Dehmel, SC&A, Inc., with copies of Tables 4 and 5 from survey questionnaires for the
years of 1994 to 1997.
Kline, R.C, I. Linins, E.L. Gershey. 1992. “Detecting Removable Surface Contamination,”
Health Phys. 62, pp. 186-189.
Kotrappa, P., S.K. Dua, D.P. Bhanti, and P.P. Joshi. 1974. “HASL Cyclone as an Instrument for
Measuring Aerosol Parameters for New Lung Model,” in Proceedings of the 3rd International
Congressional Radiation Protection Association, September 9-14, 1973.
Kusnetz, H.L. 1956. “Radon Daughers in Mine Atmospheres–A Field Method for Determining
Concentrations,” Am. Ind. Hyg. Assoc. Quarterly Vol. 17.
Laxen, D.P.H. and I.M. Chandler. 1982. “Comparison of Filtration Techniques for Size
Distribution in Freshwaters,” Analytical Chemistry, 54:8, pp. 1350-1355.
Lippmann, M. 1989a. “Calibration of Air Sampling Instruments,” in Air Sampling Instruments,
7th Ed., American Conference of Governmental Industrial Hygienists, Cincinnati, OH, pp. 73100.
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Lippmann, M. 1989b. “Sampling Aerosols by Filtration,” in Air Sampling Instruments, 7th Ed.,
American Conference of Governmental Industrial Hygienists, Cincinnati, OH, pp. 305-336.
Lockhart, L., R. Patterson and W. Anderson. 1964. Characteristics of Air Filter Media Used for
Monitoring Airborne Radioactivity. Naval Research Laboratory Report NRL-6054,
Washington, DC.
Lucas, H.F. 1982. What is the “Lucas Emanation Method for 226Ra”? Health Physics, 43:2, pp
278-279, [Letters].
Manske, P., T. Stimpfel, and E.L. Gershey. 1990. “A Less Hazardous Chromic Acid Substitute
for Cleaning Glassware,” J. Chem. Educ. 67:A280-A282.
Maron, S.H. and J. B. Lando. 1974. Fundamentals of Physical Chemistry. New York: Macmillan
Publishing Company.
MARSSIM. 2000. Multi-Agency Radiation Survey and Site Investigation Manual, Revision 1.
NUREG-1575 Rev 1, EPA 402-R-97-016 Rev1, DOE/EH-0624 Rev1. August. Available
from www.epa.gov/radiation/marssim/.
Martin, J.E. and J.M. Hylko. 1987a. “Formation of Tc-99 in Low-Level Radioactive Waste
Samples from Nuclear Plants,” Radiation Protection Management, 4:6, pp. 67-71.
Martin, J.E. and J.M. Hylko. 1987b. “Measurement of 99Tc in Low-Level Radioactive Waste
from Reactors Using 99Tc as a Tracer,” Applied Radiation and Isotopes, 38:6, pp. 447-450.
Milkey, R.G. 1954. “Stability of Dilute Solutions of Uranium, Lead, and Thorium Ions,” Anal.
Chem. 26:11, pp. 1800-1803.
National Academy of Sciences (NAS). 1960. The Radiochemistry of Technetium. Office of
Technical Services, Washington, DC.
National Committee on Radiation Protection. 1964. Safe Handling of Radioactive Materials.
NCRP Report 30, Washington, DC.
National Council on Radiation Protection and Measurements (NCRP). 1978. Instrumentation
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National Council on Radiation Protection and Measurements (NCRP). 1985. A Handbook of
Radioactivity Measurements Procedures. NCRP Report 81.
National Council on Radiation Protection and Measurements (NCRP). 1987. Use of Bioassay
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Procedures for Assessment of Internal Radionuclides Deposition. NCRP Report No. 87.
National Institute for Occupational Safety and Health (NIOSH). 1983. Industrial Hygiene
Laboratory Quality Control-587. NIOSH, Cincinnati, Ohio.
Naval Sea Systems Command (NAVSEA), 1997. Navy Environmental Compliance Sampling
and Field Testing Procedures Manual, NAVSEA T0300-AZ-PRO-010, 10 June 1997
U.S. Nuclear Regulatory Commission (NRC). Acceptable Concepts, Models, Equations, and
Assumptions for a Bioassay Program. NRC Regulatory Guide 8.9. Revision 1, September
1993.
U.S. Nuclear Regulatory Commission (NRC). Applications of Bioassay for Uranium. NRC
Regulatory Guide 8.11. June 1974.
U.S. Nuclear Regulatory Commission (NRC). 1977. Estimating Aquatic dispersion of Effluents
from Accidental and Routine Reactor Releases for the Purpose of Implementing Appendix I.
NRC Regulatory Guide 1.113.
U.S. Nuclear Regulatory Commission (NRC). Applications of Bioassay for I-125 and I-131.
NRC Regulatory Guide 8.20. Revision 1, September 1979.
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Activation Products. NRC Regulatory Guide 8.26. September 1980.
U.S. Nuclear Regulatory Commission (NRC). Bioassays at Uranium Mills. NRC Regulatory
Guide 8.22. Revisoin 1, August 1988.
U.S. Nuclear Regulatory Commission (NRC). Criteria for Establishing a Tritium Bioassay
Program. NRC Regulatory Guide 8.32. July 1988.
U.S. Nuclear Regulatory Commission (NRC). 1981. Radiation Safety Surveys at Medical
Institutions. NRC Regulatory Guide 8.23.
U.S. Nuclear Regulatory Commission (NRC). 1990. Model Feasibility Study of Radioactive
Pathways From Atmosphere to Surface Water. NUREG/CR-5475, Washington, DC.
U.S. Nuclear Regulatory Commission (NRC). 1992. National Profile on Commercially
generated Low-Level Radioactive Mixed Waste. NUREG/CR-5938, Washington, DC.
Osborne, R.V. 1973. “Sampling for Tritiated Water Vapor,” in Proceedings of the 3rd
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1973:1428-1433.
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14:62-71.
Pelto, R.H., C.J. Wierdak, and V.A. Maroni. 1975. “Tritium Trapping Kinetics in Inert Gas
Streams,” in Liquid Metals Chemistry and Tritium Control Technology Annual Report ANL75-50, p. 35 (Argonne National Laboratory, Lemont, IL).
Phillips, J.E. and C.E. Easterly. 1982. “Cold Trapping Efficiencies for Collecting Tritiated Water
Entrained in a Gaseous Stream,” Rev. Sci. Instrum., 53:1.
Pignolet, L., F. Auvray, K. Fonsny, F. Capot, and Z. Moureau. 1989. “Role of Various
Microorganisms on Tc Behavior in Sediments,” Health Physics, 57:5, pp. 791-800.
Puls, R.W., J.H. Eychaner, and R.M. Powell. 1990. Colloidal-Facilitated Transport of Inorganic
Contaminants in Ground Water: Part I. Sampling Considerations. EPA/600/M-90/023, NTIS
PB 91-168419.
Puls, R. W., and R. M. Powell. 1992. “Transport of Inorganic Colloids Through Natural Aquifer
Material: Implications for Contaminant Transport,” Environmental Science & Technology
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R. Rosson, R. Jakiel, S. Klima, B. Kahn, P.D. Fledderman. 2000. “Correcting Tritium
Concentrations in Water Vapor with Silica Gel,” Health Physics, 78:1, p 68-73. Available at:
www.srs.gov/general/sci-tech/fulltext/ms9800901/ms9800901.html.
Saar, R. A. 1997. “Filtration of Ground Water Samples: A Review of Industry Practice,” GWMR,
Winter, 1997, pp. 56-62.
Scarpitta, S.C. and N.H. Harley. 1990. “Adsorption and Desorption of Noble Gases on Activated
Charcoal. I. 133X Studies in a Monolayer and Packed Bed,” Health Physics 59:4, pp. 383-392.
Setter, L.R. and G.I. Coats. 1961. “The Determination of Airborne Radioactivity,” Industrial
Hygiene Journal, February, pp 64-69.
Sheldon, R.W. and W.H. Sutcliffe. 1969. “Retention of Marine Particles by Screens and Filters,”
Limn. & Ocean 14, pp 441-444.
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Report, Testing of Methods for the Separation of Soil and Aqueous Phases. Lawrence
Berkeley Laboratory, Report LBL-14696, UC-70.
Stafford, R.B. 1973. Comparative Evaluation of Several Glass-Fiber Filter Media. Los Alamos
Scientific Laboratory, LA-5297.
U.S. Army Corps of Engineers (USACE). 1995. Technical Project Planning Guidance for
Hazardous, Toxic and Radioactive Waste (HTRW) Data Quality Design. Engineer Manual
EM-200-1-2, Appendix H, Sampling Methods.
Waite, D.A. and W.L. Nees. 1973. A Novel Particle Sizing Technique for Health Physics
Application. Battelle, Pacific Northwest Laboratories, BNWL-SA-4658.
Whittaker, E.L., J.D. Akridge, and J. Giovino. 1989. Two Test Procedures for Radon in Drinking
Water: Interlaboratory Collaborative Study. EPA 600/2-87/082, Environmental Monitoring
Systems Laboratory.

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11 SAMPLE RECEIPT, INSPECTION, AND TRACKING
11.1 Introduction
This chapter provides guidance on laboratory sample receiving and surveying, inspecting,
documenting custody, and assigning laboratory tracking identifiers (IDs). These topics are
presented sequentially in this chapter, but they may be performed in a different order. The
chapter is directed primarily at laboratory personnel (as are all of the Part II chapters), although
the project manager and field personnel need to be aware of the steps involved in sample receipt,
inspection, and tracking. Within MARLAP, the “sample receipt” process includes the surveying
of the package and sample containers for radiological contamination and radiation levels.
“Sample inspection” means checking the physical integrity of the package and samples,
confirming the identity of the sample, confirming field preservation (if necessary), and recording
and communicating the presence of hazardous materials. “Laboratory sample tracking” is a
process starting with logging in the sample and assigning a unique laboratory tracking identifier
(numbers and/or letters) to be used to account for the sample through analyses, storage, and
shipment. Laboratory tracking continues the tracking that was initiated in the field during sample
collection (see Section 10.2, “Field Sampling Plan: Non-Matrix-Specific Issues”).
This chapter focuses on sample receipt, inspection, and tracking of samples in the laboratory
because these are the three modes of initial control and accountability (Figure 11.1). Sample
receipt and inspection activities need to be done in a timely manner to allow the laboratory and
field personnel to resolve any problems (e.g., insufficient material collected, lack of field
preservation, etc.) with the samples received by the laboratory as soon as is practical. Effective
communications between field personnel and the laboratory not only facilitates problem
resolution but also prevents unnecessary delays in the analytical process.
Other relevant issues, including the laboratory’s radioactive materials license conditions and
proper operating procedures, are also discussed because these topics are linked to receipt,
inspection, and tracking activities. The result of the sample receipt and inspection activities is to
accept the samples as received or to perform the necessary corrective action (which may include
rejecting samples). Health and safety information on radiological issues can be found in NRC
(1998a; 1998b).

11.2 General Considerations
Contents

11.2.1 Communication Before Sample
Receipt
Before the samples are received, the laboratory
should know the approximate number of
samples that will be received within a specific
JULY 2004

11.1
11.2
11.3
11.4
11.5
11.6

11-1

Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . 11-1
General Considerations . . . . . . . . . . . . . . . . . 11-1
Sample Receipt . . . . . . . . . . . . . . . . . . . . . . . 11-5
Sample Inspection . . . . . . . . . . . . . . . . . . . . . 11-8
Laboratory Sample Tracking . . . . . . . . . . . . 11-11
References . . . . . . . . . . . . . . . . . . . . . . . . . . 11-13
MARLAP

Sample Receipt, Inspection, and Tracking

Field Sample Shipment
(see Chapter 10)
Tracking Documents

• Number and type of samples along
with field sample number
• Field processing and preservation
• Analysis requested

Sample Receipt

In designated rad receiving/prep area:
• Check container labels against sample
• Check radionuclides requested
against tracking documents
• Check tamper seals
• Verify preservation against tracking
documents
• Check field preparation against
tracking documents

Laboratory
Tracking
Sample Inspection

SAMPLE
REJECTED

Check with client
relative to sample
disposition

Sample received in designated area
• Authorized user notified for radiological
screening of package
• Check for evidence of breakage or
leakage of exterior of shipping
package, then shipping containers.
If found, radiologically survey and
decontaminate if necessary
• Radiological survey
• License requirements
• Chain-of-custody procedures if
required

SAMPLE
ACCEPTED

Any discrepancies in the following will
result in corrective action:
• Survey limits
• Expected radionuclides
• Number and type of samples
• Tracking Documents

Short-term
sample storage
or sample
prep/analysis
laboratory

FIGURE 11.1 — Overview of sample receipt, inspection, and tracking

period of time and the types of analyses that are expected for the samples. Laboratory personnel
should be provided with a contact in the field and with means of contacting the person
(telephone, FAX, e-mail). The information about the client, points of contact, number of samples,
and types of analyses can be entered into the laboratory information management system (LIMS)
to facilitate communication between the laboratory—in both the sample receipt area and the
project management area—and the project manager. Communication between laboratory
personnel and project staff in the field allows the parties to coordinate activities, schedules, and
sample receipt. In particular, the project manager should provide to the laboratory any special
instructions regarding the samples before shipment of samples. This information serves to notify
MARLAP

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the laboratory of health and safety concerns and provides details that will affect analytical
procedures, sample disposition, etc. For example, without this communication, a laboratory
might receive a partial shipment and not realize that samples are missing. Furthermore, advance
communications allow laboratory staff to arrange for special handling or extra storage space
should the need arise.
Planning for the samples to be received at the laboratory starts during the development of the
appropriate plan document and the statement of work (SOW) and continues through the
communication between the project staff in the field and the laboratory. For example, the
laboratory could use its LIMS to generate labels and bar-codes for the appropriate containers to
be used in the field. This process would assist in assigning appropriate sample IDs for the
laboratory tracking system, which starts with sample receipt. The laboratory should instruct the
field staff to place the tracking documents on the inside of the cooler lid for easy access and to
include any other pertinent information (field documentation, field surveying information, etc.).
11.2.2 Standard Operating Procedures
A laboratory should have standard operating procedures (SOPs) for activities related to sample
receipt, inspection, and tracking. Some typical topics that might be addressed in laboratory SOPs
are presented in Table 11.1. For example, the laboratory should have an SOP that describes what
information should be included in the laboratory sample tracking system. Laboratory SOPs
should describe chain-of-custody procedures giving a comprehensive list of the elements in the
program such as signing the appropriate custody forms, storing samples in a secure area, etc.
(ASTM D4840; ASTM D5172; EPA, 1995).
TABLE 11.1 — Typical topics addressed in standard operating procedures related to
sample receipt, inspection, and tracking
Sample
Receipt:

• Order and details for activities associated with receiving shipments of samples
• Surveying methods

Inspection:

•
•
•
•
•

Tracking:

• Maintain chain of custody and document sample handling during transfer from the field to
the laboratory, then within the laboratory
• Ensure proper identification of samples throughout process
• Procedures to quickly determine location and status of samples within laboratory

Custodian:

• Execution of responsibilities of the sample custodian

Forms/Labels:

• Examples of forms and labels used to maintain sample custody and document sample
handling in the laboratory

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Check physical integrity
Confirm sample identification
Identify/manage hazardous materials
pH measurement instructions
Use the laboratory information management system (LIMS) to assign laboratory sample IDs

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The laboratory needs to establish corrective action guidelines (Section 11.3.3) as part of every
SOP for those instances when a nonconformance is noted. Early recognition of a nonconformance will allow the project manager and the laboratory more options for a quick resolution.
11.2.3 Laboratory License
Laboratories that handle radioactive materials are required (with few exceptions, such as certain
U.S. Department of Energy National Laboratories and Department of Defense laboratories) to
have a radioactive materials license issued by the NRC or the Agreement State in which the
laboratory operates. The radioactive materials license lists the radionuclides that the laboratory
can possess, handle, and store. In addition, the license limits the total activity of specific
radionuclides that can be in the possession of the laboratory at a given time.
The client must have a copy of the current radioactive materials license for the facility to which
the samples are being shipped. The laboratory staff and the project manager all need to be aware
of the type of radionuclide(s) in the samples and the total number of samples to be sent to the
laboratory. This information should be included in the appropriate plan document and SOW prior
to sampling.
The laboratory is required by the license to maintain a current inventory of certain radioactive
materials present in the facility. The radioactive materials license also requires the laboratory to
develop and maintain a radiation protection plan (NRC, 1998b) that states how radioactive
samples will be received, stored, and disposed. The laboratory will designate an authorized user
(NRC, 1998b) to receive the samples. A Radiation Safety Officer (RSO) may be an authorized
user, but not always. NRC (1998b) gives procedures for the receipt of radioactive samples during
working hours and non-working hours.
11.2.4 Sample Chain-of-Custody
Sample chain-of-custody (COC) is defined as a process whereby a sample is maintained under
physical possession or control during its entire life cycle, that is, from collection to disposal
(ASTM D4840—see Section 10.2.7). The purpose of COC is to ensure the security of the sample
throughout the process. COC procedures dictate the documentation needed to demonstrate that
COC is maintained. When a sample is accepted by the laboratory it is said to be in the physical
possession or control of the laboratory. ASTM D4840 states that a sample is under “custody” if it
is in possession or under control so as to prevent tampering or alteration of its characteristics.
If the samples are transferred under COC, the relinquisher and the receiver should sign the
appropriate parts of the COC form with the date and time of transfer (see Figure 10.1). After
receipt and inspection the samples should be kept in a locked area or in an area with controlled
access.
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COC is not a requirement for all samples. COC is most often required when the sample data may
be used as legal evidence. The project plan should state whether COC will be required. The
paperwork received with the samples should also indicate whether COC has been maintained
from the time of collection and must be maintained in the laboratory. If the laboratory has been
informed that COC procedures should be followed, but it appears that appropriate COC
procedures have not been followed (before or after sample receipt at the laboratory) or there are
signs of possible sample tampering when the samples arrive, the project manager should be
contacted. The problem and resolution should be documented. Additional information on COC
can be found in EPA (1985).

11.3 Sample Receipt
Laboratory sample receipt occurs when a package containing samples is accepted, the package
and sample containers are surveyed for external surface radiological contamination and radiation
level, and the physical integrity of the package and samples is checked. Packages include the
shipping parcel that holds the smaller sample containers with the individual samples (see Section
11.3.2 on radiological surveying). Also note that topics and activities covered in Section 11.3
appear in a sequence but, in many cases, these activities are performed simultaneously during
initial receiving activities (i.e., package surveying and observation of its physical integrity).
11.3.1 Package Receipt
Some laboratories require arriving samples to go through a security inspection process at a
central receiving area before routing them to the appropriate laboratory area(s). In addition, if
samples are shipped by an air transport carrier, the shipping containers may be subject to airport
security. In these cases, the container housing the samples may be opened and the samples
inspected and reinserted in an order not consistent with the original packaging. In these cases, it
is imperative that each individual sample container have a permanent identifier either in indelible
ink or as a label affixed on the side of the sample container (see Section 10.2.4, “Container Label
and Sample Identification Code”). Within each shipping container, a separate sample packing
slip or tracking documents that lists the samples (by sample ID) for the container should be
included.
Packages should be accepted only at designated receiving areas. Packages brought to any other
location by a carrier should be redirected to the appropriate receiving area. All packages labeled
RADIOACTIVE I, II, or III require immediate notification of the appropriate authorized user (NRC,
1998b).
A sample packing slip or tracking documents is required and must be presented at the time of
receipt, and the approximate activity of the shipment should be compared to a list of acceptable
quantities. If known, the activity of each radionuclide contained in the shipment must be
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reviewed relative to the total amount of that radionuclide currently on site to ensure that the
additional activity will not exceed that authorized by the NRC or Agreement State in the
laboratory’s license.
Surveying measures described in Section 11.3.2 may indicate that the samples are more
radioactive than expected and that the radiation license limit may be exceeded. The laboratory
should take extra precautions with these samples, but the survey results should be verified. The
federal, state, or local agency should be contacted immediately when verified license limits are
exceeded. The laboratory must respond quickly to stay in compliance with its license.
If the package is not accepted by the laboratory, the laboratory should follow corrective-action
procedures prescribed in the radiation materials license, the appropriate plan document (if this is
a reasonable possibility for the project), and the laboratory’s SOPs. The project manager should
be contacted about possible disposition of any samples.
11.3.2 Radiological Surveying
In addition to ensuring compliance with the laboratory’s license and verifying estimates of radionuclide activity (Section 11.3.1), the radiological surveying of packages during sample receipt
serves to identify and prevent the spread of external contamination. All packages containing
samples for analysis received by the laboratory should be surveyed for external contamination
using a wipe (sometimes referred to as a “swipe”) and for surface exposure rate using the appropriate radiation survey meter. Exceptions may include known materials intended for analysis as:
well-characterized samples, bioassays, or radon and associated decay products in charcoal media
(exceptions should be listed in the laboratory SOP). Surveying of packages and sample
containers received in the laboratory should be conducted in accordance with the laboratory’s
established, documented procedures and the laboratory radiation protection and health and safety
plan. The exterior of the package is surveyed first; if there is no evidence of contamination or that
the laboratory licence would be exceeded, the package is opened up and the sample containers
surveyed individually. These procedures should include the action level and appropriate action as
established by the facility. Personnel performing surveying procedures should be proficient in the
use of portable radiation surveying instruments and knowledgeable in radiological contamination
control procedures. Health and safety considerations are affected by the suspected or known
concentrations of radionuclides in a sample or the total activity of a sample.
Radiation surveying is normally conducted using Geiger-Mueller (GM) detectors, ionization
chambers, micro-R meters, or alpha scintillation probes, as appropriate. The laboratory should
refer to any information they obtained before receipt of samples or with the samples, especially
concerning the identity and concentration of radioactive and chemical constituents in the
samples. Radiological surveying needs to be performed as soon as practical after receipt of the
package, but not later than three hours (10 CFR 20.1906) after the package is received at the
licensee’s facility for packages received during normal working hours. For packages received
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outside of normal working hours, the surveying must be performed no later than three hours from
the beginning of the next workday.
Survey the exterior of a labeled package for radioactive contamination (10 CFR 20.1906). If the
package is small (less than 100 cm2), the whole package should be wiped (swiped). Wipes are not
always used, but if there is reason to believe that something has leaked, then wipes should be
used. This survey is performed to detect possible violations of Department of Transportation
(DOT) packaging and labeling regulations, as well as to determine the possible presence of
gamma- and some beta-emitting radionuclides that may require special handling. Also, such a
survey can help to avoid introducing a high-activity sample into a low-activity area. NRC
(1998b) gives the following sample model for opening packages containing radioactive material:
• Wear gloves to prevent hand contamination.
• Visually inspect the package for any sign of damage (e.g. crushed, punctured). If damage is
noted, stop and notify the RSO.
• Check DOT White I, Yellow II, or Yellow III label or packing slip for activity of contents, so
shipment does not exceed license possession limits.
• Monitor the external surfaces of a labeled package according to specifications in Table 8.4,
Section 13.14, Item 10 [of NRC, 1998b].
• Open the outer package (following supplier’s directions if provided) and remove packing
slip. Open inner package to verify contents (compare requisition, packing slip and label on
the bottle or other container). Check integrity of the final source container (e.g., inspecting
for breakage of seals or vials, loss of liquid, discoloration of packaging material, high count
rate on smear). Again check that the shipment does not exceed license possession limits. If
you find anything other than expected, stop and notify the RSO.
• Survey the packing material and packages for contamination before discarding. If contamination is found, treat them as radioactive waste. If no contamination is found, obliterate the
radiation labels prior to discarding in the regular trash.
• Maintain records of receipt, package survey, and wipe test results.
• Notify the final carrier and by telephone, telegram, mailgram, or facsimile, the administrator
of the appropriate NRC Regional Office listed in 10 CFR 20, Appendix D when removable
radioactive surface contamination exceeds the limits of 10 CFR 71.87(i); or external radiation
levels exceed the limits of 10 CFR 71.47.
In addition to these, laboratories may have additional internal notifications or procedures.
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11.3.3 Corrective Action
The laboratory’s SOPs should specify corrective actions for routine and non-routine sample
problems, including deficiency in sample volume, leaking samples, and labeling errors. The
appropriate corrective action may require consulting the project manager and other laboratory
personnel. Timely response can allow for a broader range of options and minimize the impact of
the sample problem on the project. The laboratory should document the problem, the cause (if
known), the corrective action taken, and the resolution of each problem that requires corrective
action. The documentation should be included in the project files.

11.4 Sample Inspection
After sample receipt, the next steps are to confirm that the correct sample has been sent, to check
that the appropriate field preservation and processing have been performed, and to identify any
hazardous chemicals.
Documents accompanying the samples should be reviewed upon receipt of the samples at the
laboratory. If the proper paperwork is not present, the project manager should be notified. Data
recorded on the paperwork, such as collection dates, sample descriptions, requested analyses, and
field staff personnel, should be compared to data on the sample containers and other documentation. Any deficiencies or discrepancies should be recorded by the laboratory and reported to the
project manager. The documents can provide data useful for health and safety surveying,
tracking, and handling or processing of critical short-lived radionuclides.
11.4.1 Physical Integrity of Package and Sample Containers
Sample containers should be thoroughly inspected for evidence of sample leakage. Leakage can
result from a loose lid, sample container puncture, or container breakage. Packages suspected to
contain leaking sample containers should be placed in plastic bags. The authorized user or alternate authorized user must be notified immediately for assistance. If leakage has occurred, appropriate radiological and chemical contamination controls should be implemented. Sample materials
that have leaked or spilled are normally not suitable for analysis and should be properly disposed.
In all cases, the laboratory’s management and project manager should be notified of leaks,
breakage, spills, and the condition of sample materials that remain in the original containers.
Sample containers that have leaked (from a loose lid or puncture) may still hold enough sample
for the requested analyses, so the laboratory should first determine whether sufficient representative sample remains. The sample is not usually analyzed if its integrity was compromised or is in
doubt. Unless appropriate information is provided in the project plan or SOW, the project
manager should determine whether or not the sample materials can be used for analysis or if new
samples are required.
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Packages, cooler chests, or individual sample containers may arrive at the laboratory bearing
custody seals. These seals provide a means to detect unauthorized tampering. When packages or
samples arrive with custody seals, they should be closely inspected for evidence of tampering.
Custody seals are made from material that cannot be removed without tearing. If a custody seal is
torn or absent, sample tampering may have occurred. This evidence of possible tampering is
generally sufficient to preclude use of the sample for laboratory analyses. The project manager
should be notified of the condition of the custody seal to determine if new samples are needed.
Observations regarding the condition of the custody seals should be recorded according to the
laboratory’s standard procedures.
11.4.2 Sample Identity Confirmation
Visual inspection is the means to confirm that the correct sample has been received. Verifying
the identity of a sample is a simple process where the appearance, sample container label, and
chain-of-custody record or tracking documents are compared. If all three sources of information
identify the same sample, then the sample is ready for the next step. If the sample label indicates
the sample is a liquid and the container is full of soil, this discrepancy would indicate nonconformance. If the sample label states that there is 1,000 mL of liquid and there only appears to be 200
mL in the container, there may be nonconformance. Visual inspection can be used to:
• Verify identity of samples by matching container label IDs and tracking documents;
• Verify that the samples are as described by matrix and quantity;
• Check the tamper seal (if used);
• Verify field preparation (e.g., filtering, removing extraneous material ), if indicated; and
• Note any changes to samples’ physical characteristics that are different than those in the
tracking documents.
11.4.3 Confirmation of Field Preservation
For those liquid samples requiring acid preservation, pH measurements may be performed on all
or selected representative liquid samples to determine if acid has been added. The temperature of
the sample may also be part of field preservation and the actual measured temperature should be
compared to the specified requirements in the documentation.
11.4.4 Presence of Hazardous Materials
The presence of hazardous materials in a sample typically creates the need for additional health
and safety precautions when handling, preparing, analyzing, and disposing samples. If there is
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documentation on the presence of non-radiological hazardous constituents, the project manager
should notify the laboratory about the presence of these chemicals. These chemical contaminants
should be evaluated by the laboratory to determine the need for special precautions. The
laboratory can also perform preliminary sample surveying for chemical contaminants using
surveying devices such as a photoionization detector for volatile components. The presence of
suspected or known hazardous materials in a sample should be identified, if possible, during
project planning and documented in the plan document and SOW. Visual inspection can also be
used such as checking the color of the sample (e.g., a green-colored water sample may indicate
the presence of high chromium levels). The presence of suspected or known hazardous materials
determined in the field should be communicated to the laboratory prior to the arrival of samples
and noted on documentation accompanying the samples to the laboratory. If no documentation on
non-radiological hazardous constituents is available, the laboratory should review previous
experience concerning samples from the site to assess the likelihood of receiving samples with
chemical contaminants. The laboratory’s chemical hygiene officer and the project manager
should be notified about the presence of potentially hazardous chemical contaminants.
11.4.5 Corrective Action
Visual inspection can also verify whether field sample preparation was performed as stated in
accompanying documentation. Samples that were not filtered in the field or that reacted with the
preservative to form a precipitate may represent a significant problem to the laboratory. If it
appears that the sample was filtered in the field (e.g., there is no corresponding filter or there are
obviously solid particles in a liquid sample), the liquid generally will be analyzed as originally
specified. Laboratory personnel should check the project plan or SOW to see if the filter and
filtered materials require analyses along with the filtered sample. If it appears that the sample was
not filtered in the field (i.e., there is no corresponding filter or there are obviously solid particles
in a liquid sample), sample documentation should be reviewed to determine if a deviation from
the project plan was documented for the sample. It may be appropriate to filter the sample in the
laboratory. The project manager should be notified immediately to discuss possible options such
as filtering the sample at the laboratory or collecting additional samples.
One example of a corrective action for inspection is, if the pH is out of conformance, it may be
possible to obtain a new sample. If it is not possible or practical to obtain a new sample, it may
be possible to acidify the sample in the laboratory.
Visual inspection can serve to check certain aspects of sample collection. For example, if the
SOP states that a soil sample is supposed to have twigs, grass, leaves, and stones larger than a
certain size removed during sample collection and some of this foreign material is still included
as part of the sample, this discrepancy results in a nonconformance.

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11.5 Laboratory Sample Tracking
Sample tracking should be done to ensure that analytical results are reported for the “correct”
sample. Sample tracking is a process by which the location and status of a sample can be
identified and documented. The laboratory is responsible for sample tracking starting with receipt
(at which time a unique laboratory sample ID is assigned), during sample preparation, and after
the performance of analytical procedures until final sample disposition. The process of sample
tracking begins the moment a field worker assigns an identification number (based on the
information provided in the appropriate plan document) and documents how materials are
collected. The way samples are transported from the field to the laboratory should be
documented. The sample receipt procedures and documentation should be consistent when
applicable with 10 CFR Part 20 Subpart J, and the client’s requirements as stated in the
appropriate plan document or statement of work.
11.5.1 Sample Log-In
Laboratory sample IDs should be assigned to each sample in accordance with the laboratory’s
SOP on sample codes. Each sample should receive a unique sample ID by which it can be logged
into the LIMS, scheduled for analysis, tracked, and disposed. Information to be recorded during
sample log-in should include the field sample identification number, laboratory sample ID, date
and time samples were collected and received, reference date for decay calculations, method of
shipment, shipping numbers, condition of samples, requested analyses, number and type of each
sample, quality control requirements, special instructions, and other information relevant to the
analysis (e.g., analytical requirements or MQOs) and tracking of samples at the laboratory.
Laboratory sample tracking is a continuation of field sample tracking. Some of this information
may have been entered into the LIMS during the planning phase.
Documents generated for laboratory sample tracking must be sufficient to verify the sample
identity, that the sample may be reliably located, and that the right sample is analyzed for the
right analyte. The documentation should include sample log-in records, the analysis request form,
names of staff responsible for the work, when procedures are completed, and details concerning
sample disposal. The documentation must conform to the laboratory’s SOPs.
During sample log-in, laboratory quality control (QC) samples may be scheduled for the analyses
requested. The type and frequency of QC samples should be provided by the plan document or
SOW and consistent with the laboratory’s SOPs.
11.5.2 Sample Tracking During Analyses
At this point, samples are introduced into the laboratory’s analytical processing system. The
information gathered during surveying, along with the assigned tracking identification, passes to
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the laboratory where specific preparation and analyses are performed. The sample may be further
subsampled. Each subsample, along with the original sample, requires tracking to account for all
materials handled and processed in the laboratory.
Each set of samples received by the laboratory should be accompanied by documents identifying
the analytes required for each sample. These documents should be reviewed against the project
plan documents or the SOW, which should identify the analytes, matrices, and analytical
requirements and be part of the project documentation prior to the samples being received by the
laboratory. Laboratory management personnel should be notified of any discrepancies. The
requested analyses should be entered into the laboratory’s tracking system. Typically, only one
sample container of sufficient volume or quantity will be provided for a single or multiple set of
different analyses. Each aliquant removed from the original container may require tracking (and
perhaps a different laboratory sample ID).
Aliquants used during the analytical process can be tracked using analysis laboratory notebooks,
forms, or bench sheets that record laboratory sample IDs, analyte, reference date for decay
correction, aliquant size, and designated quality control samples. Bench sheets are loose-leaf or
bound pages used to record information during laboratory work and are used to assist in sample
tracking. Each sheet is helpful for identifying and processing samples in batches that include
designated QC samples. The bench sheet, along with the laboratory log book, can later be used to
record analytical information for use during the data review process. Bench sheets can also be
used to indicate that sample aliquants were in the custody of authorized personnel during the
analytical process.
After receipt, verification of sample information and requested analyses, and assignment of
laboratory sample IDs, the requested analyses can be scheduled for performance in accordance
with laboratory procedures. Using this system, the laboratory can formulate a work schedule, and
completion dates can be projected.
11.5.3 Storage of Samples
If samples are to be stored and analyzed at a later date, they should be placed in a secure area.
Before storage, any special preservation requirements, such as refrigeration or additives, should
be determined.
The laboratory should keep records of the sample identities and the location of the sample
containers. Unused sample aliquants should be returned to the storage area for final disposition.
In addition, for some samples, depending on the level of radioactivity or hazardous constituents
present, the laboratory should record when the sample was disposed and the location of the
disposal facility. These records are necessary to ensure compliance with the laboratory’s license
for radioactive materials and other environmental regulations.
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Areas where samples are stored should be designated and posted as radioactive materials storage
areas. Depending on the activity level of the samples, storage areas may require special posting.
If additional storage space or shielding is needed, arrangements that are consistent with the
license should be made with the authorized user. See Chapter 17 for more information on waste
disposal.

11.6 References
American Society for Testing and Materials (ASTM) D4840. Standard Guide for Sampling
Chain-of-Custody Procedures. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) D5172. Standard Guide for Documenting
the Standard Operating Procedures Used for the Analysis of Water. West Conshohocken,
PA.
U.S. Environmental Protection Agency (EPA). 1985. NEIC Policies and Procedures. National
Enforcement Information Center. EPA-300/9-78DDI-R, June.
U.S. Environmental Protection Agency (EPA). 2001. Guidance for the Preparation of Standard
Operating Procedures (SOPS) for Quality-Related Documents (QA/G-6). EPA/240/B01/004, March. Available at: www.epa.gov/quality/qa_docs.html.
U.S. Nuclear Regulatory Commission (NRC). 1998a. Procedures for Receiving and Opening
Packages. 10 CFR Part 20.
U.S. Nuclear Regulatory Commission (NRC). 1998b. Consolidated Guidance About Materials
Licenses, Volume 7. (NRC91). NUREG 1556.

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12 LABORATORY SAMPLE PREPARATION
12.1 Introduction
On first impression, sample preparation may seem the most routine aspect of an analytical
protocol. However, it is critical that analysts realize and remember that a measurement is only as
good as the sample preparation that has preceded it. If an aliquant taken for analysis does not
represent the original sample accurately, the results of this analysis are questionable. As a general
rule, the error in sampling and the sample preparation portion of an analytical procedure is
considerably higher than that in the methodology itself, as illustrated in Figure 12.1.

Sampling
Sample
Preparation
Concentration, Separation,
Isolation, etc. Steps
Measurement
(After Scwedt, 1997)
FIGURE 12.1—Degree of error in laboratory sample preparation relative to other activities

One goal of laboratory sample preparation is to provide, without sample loss, representative
aliquants that are free of laboratory contamination that will be used in the next steps of the
protocol. Samples are prepared in accordance with applicable standard operating procedures
(SOPs) and laboratory SOPs using information provided by field sample preparation (Chapter 10,
Field and Sampling Issues that Affect Laboratory Measurements), sample screening activities,
and objectives given in the appropriate planning documents. The laboratory sample preparation
techniques presented in this chapter include the
Contents
physical manipulation of the sample (heating,
screening, grinding, mixing, etc.) up to the
12.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . 12-1
12.2 General Guidance for Sample Preparation . . 12-2
point of dissolution. Steps such as adding
carriers and tracers, followed by wet ashing or 12.3 Solid Samples . . . . . . . . . . . . . . . . . . . . . . . 12-12
12.4 Filters . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-30
fusion, are discussed in Chapter 13 (Sample
12.5 Wipe Samples . . . . . . . . . . . . . . . . . . . . . . . 12-31
Dissolution) and Chapter 14 (Separation
12.6 Liquid Samples . . . . . . . . . . . . . . . . . . . . . . 12-32
Techniques).
12.7 Gases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12-36
This chapter presents some general guidance
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12.9 References . . . . . . . . . . . . . . . . . . . . . . . . . . 12-37

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for sample preparation to avoid sample loss and sample contamination. Due to the physical
nature of the matrix, sample preparation for solids requires the most attention, and therefore is
discussed at great length (Section 12.3). General procedures for preparing solid samples (such as
drying, obtaining a constant weight, grinding, sieving, mixing, and subsampling) are discussed.
Some sample preparation procedures then are presented for typical types of solid samples (e.g.,
soil and sediment, biota, food, etc.). This chapter concludes with specific guidance for preparing
samples of filters (Section 12.4), wipes (Section 12.5), liquids (Section 12.6), gases (Section
12.7), and bioassay (Section 12.8).

12.2 General Guidance for Sample Preparation
Some general considerations during sample preparation are to minimize sample losses and to
prevent contamination. Possible mechanisms for sample loss during preparation steps are
discussed in Section 12.2.1, and the contamination of samples from sources in the laboratory is
discussed in Section 12.2.2. Control of contamination through cleaning labware is important and
described in Section 12.2.3, and laboratory contamination control is discussed in Section 12.2.4.
12.2.1 Potential Sample Losses During Preparation
Materials may be lost from a sample during laboratory preparation. The following sections
discuss the potential types of losses and the methods used to control them. The addition of tracers
or carriers (Section 14.9) is encouraged at the earliest possible point and prior to any sample
preparation step where there might be a loss of analyte. Such preparation steps may include
homogenization or sample heating. The addition of tracers or carriers prior to these steps helps to
account for any analyte loss during sample preparation.
12.2.1.1

Losses as Dust or Particulates

When a sample is dry ashed, a fine residue (ash) is often formed. The small particles in the
residue are resuspended readily by any air flow over the sample. Air flows are generated by
changes in temperature (e.g., opening the furnace while it is hot) or by passing a stream of gas
over the sample during heating to assist in combustion. These losses are minimized by ashing
samples at as low a temperature as possible, gradually increasing and decreasing the temperature
during the ashing process, using a slow gas-flow rate, and never opening the door of a hot
furnace (Section 12.3.1). If single samples are heated in a tube furnace with a flow of gas over
the sample, a plug of glass or quartz wool can be used to collect particulates or an absorption
vessel can be used to collect volatile materials. At a minimum, all ash or finely ground samples
should be covered before they are moved.
Solid samples are often ground to a fine particle size before they are fused or wet ashed to
increase the surface area and speed up the reaction between the sample and the fluxing agent or
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acid (see Chapters 13 and 14 on dissolution and separation). Since solid samples are frequently
heterogeneous, a source of error arises from the difference in hardness among the sample
components. The softer materials are converted to smaller particles more rapidly than the harder
ones, and therefore, any loss in the form of dust during the grinding process will alter the
composition of the sample. The finely ground particles are also susceptible to resuspension.
Samples may be moistened carefully with a small amount of water before adding other reagents.
Reagents should be added slowly to prevent losses as spray due to reactions between the sample
and the reagents.
12.2.1.2

Losses Through Volatilization

Some radionuclides are volatile under specific conditions (e.g., heat, grinding, strong oxidizers),
and care should be taken to identify samples requiring analysis for these radionuclides. Special
preparation procedures should be used to prevent the volatilization of the radionuclide of interest.
The loss of volatile elements during heating is minimized by heating without exceeding the
boiling point of the volatile compound. Ashing aids can reduce losses by converting the sample
into less volatile compounds. These reduce losses but can contaminate samples. During the wet
ashing process, losses of volatile elements can be minimized by using a reflux condenser. If the
solution needs to be evaporated, the reflux solution can be collected separately. Volatilization
losses can be prevented when reactions are carried out in a properly constructed sealed vessel.
Table 12.1 lists some commonly analyzed radioisotopes, their volatile chemical form, and the
boiling point of that species at standard pressure. Note that the boiling point may vary depending
upon solution, matrix, etc.
Often the moisture content, and thus, the chemical composition of a solid is altered during
grinding and crushing (Dean, 1995). Decreases in water content are sometimes observed while
grinding solids containing essential water in the form of hydrates, likely as a result of localized
heating. (See Section 12.3.1.2 for a discussion of the types of moisture present in solid samples.)
Moisture loss is also observed when samples containing occluded water are ground and crushed.
The process ruptures some of the cavities, and exposes the water to evaporation. More commonly, the grinding process results in an increase in moisture content due to an increase in
surface area available for absorption of atmospheric water. Both of these conditions will affect
the analysis of 3H since 3H is normally present in environmental samples as 3HOH. Analysis for
tritium in soils should avoid these types of sample preparation prior to analysis. Instead, total
water content should be determined separately. Tritium analysis then could be performed by
adding tritium-free (“dead”) water to an original sample aliquant followed by filtration or
distillation.

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TABLE 12.1 — Examples of volatile radionuclides
Isotope

Chemical Form

Boiling Point (EC) *

Tritium — H

H2O

100E

Carbon — 14C

CO2 (produced from CO3-2 or
oxidation of organic material)

-78.5E

Magnesium, calcium, and sodium
carbonates

Natural ores of these metals decompose
between 825E and 1,330E to yield the
respective metal oxides

I2

185.2E (sublimes readily)

Cs0 (as metal)
Cs2O (as metallic oxide)
(nitrates decompose to oxides)
CsCl (as metallic chloride)

678.4E (melts at 28)
~400E

Tc2O7
TcCl4
TcO2

310.6E
Sublimes above 300E
Sublimes above 900E

3

Iodine — 131I, 129I
134

Cesium — Cs,
136
Cs, 137Cs

135

Cs,

Technetium — 99Tc

1290E

[Most Tc compounds sublime above 300E. Tc(VII) is an oxidant that reacts
with organic solvents forming Tc(IV)]
Polonium —
210
Po

208

Po,

209

Po,

Po0
PoCl4
Po(NO3)4 [as a solid]
PoO2

962E
390E
Decomposes to PoO2 above ~150E
Decomposes to Po metal above 500E

Pb0
1744E
PbCl2
950E
210
212
205
Lead — Pb, Pb, Pb
Decomposes to oxide above 470E
Pb(NO3)2
888E
PbO
* The closer the sample preparation temperature is to the boiling point of the compound, the more significant will be
the loss of the material. However, if the objective is to distill the analyte compound from other nonvolatile
materials, then boiling temperature is needed. Sample preparation near the decomposition temperature should be
avoided for those compounds that have a decomposition temperature listed in the table.
Sources: Greenwood and Earnshaw (1984); Windholz (1976); Schwochau (2000); Sneed and Brasted (1958).

Additional elements that volatilize under specific conditions include arsenic, antimony, tin,
polonium, lead, selenium, mercury, germanium, and boron. Chromium can be volatilized in
oxidizing chloride media. Carbon, phosphorus, and silicon may be volatilized as hydrides, and
chromium is volatilized under oxidizing conditions in the presence of chloride. The elements in
Table 12.1 are susceptible to changing oxidation states during sample preparation. Thus, the
pretreatment should be suited to the analyte. The volatility of radionuclides of tritium, carbon,
phosphorus, and sulfur contained in organic or bio-molecules is based on the chemical properties
of those compounds. If such compounds are present, special precautions will be necessary during
sample preparation to avoid the formation of volatile compounds or to capture the volatilized
materials.
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12.2.1.3

Losses Due to Reactions Between Sample and Container

Specific elements may be lost from sample materials from interaction with a container. Such
losses may be significant, especially for trace analyses used in radioanalytical work. Adsorption
reactions are discussed in Chapter 10 for glass and plastic containers. Losses due to adsorption
may be minimized by using pretreated glassware with an established hydrated layer. Soaking new
glassware overnight in a dilute nitric or hydrochloric acid solution will provide an adequate
hydrated layer. Glassware that is used on a regular basis will already have established an
adequate hydrated layer. The use of strong acids to maintain a pH less than one also helps
minimize losses from adsorption.
Reactions among analytes and other types of containers are described in Table 12.2. Leaving
platinum crucibles uncovered during dry ashing to heat samples will minimize reduction of
samples to base metals that form alloys with platinum. Porcelain should not be used for analysis
of lead, uranium, and thorium because the oxides of these elements react with porcelain glazes.
Increasing the amount of sample for dry ashing increases the amount of ash, minimizing the loss
of the sample’s trace materials to the container surface.
TABLE 12.2 — Properties of sample container materials
Recommended
Material
Properties
Use
Borosilicate General
Transparent; good thermal properties; fragile; attacked by HF, H3PO4, and
alkaline solutions.
Glass
applications
Fused Quartz High temperature
Transparent; excellent thermal properties (up to 1,100 EC); fragile; more
applications
expensive than glass; attacked by HF, H3PO4, and alkaline solutions.
Porcelain
High temperature
Used at temperatures up to 1,100 EC; less expensive than quartz; attacked by
applications and
HF, H3PO4, and alkaline solutions.
pyrosulfate fusion
Nickel
Molten alkali metal Suitable for use with strongly alkaline solutions. Do not use with HCl.
hydroxide and
Na2O2 fusions
Platinum
High temperature
Virtually unaffected by acids, including HF; dissolves readily in mixtures of
or corrosive
HNO3 and HCl, Cl2 water or Br2 water; adequate resistance to H3PO4; very
expensive; forms alloys with Hg, Pb, Sn, Au, Cu, Si, Zn, Cd, As, Al, Bi, and
applications
Fe, which may be formed under reducing conditions; permeable to H2 at red
heat, which serves as a reducing agent; may react with S, Se, Te, P, As, Sb, B,
and C to damage container; soft and easily deformed, often alloyed with Ir,
Au, or Rh for strength. Do not use with Na2CO3 for fusion.
Zirconium
Peroxide fusions
Less expensive alternative to platinum; extremely resistant to HCl; resistant to
HNO3; resistant to 50% H2SO4 and 60% H3PO4 up to 100 EC; resistant to
molten NaOH; attacked by molten nitrate and bisulfate; usually available as
Zircaloy—98% Zr, 1.5% Sn, trace Fe, Cr, and Ni. Do not use with KF or HF.

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Material
Alumina
(Al2O3)
Polyethylene

Recommended
Use
Acids and alkali
melts at low
temperatures
Sample and reagent
storage

Properties
Resistant to acids and alkali melts; rapidly attacked by bisulfate melts; brittle,
requires thick walled containers.
Resistant to many acids; attacked by 16M HNO3 and glacial acetic acid;
begins to soften and lose shape at 60 EC; appreciably porous to Br2, NH3,
H2S, H2O, and HNO3 (aqueous solutions can lose ~1% volume per year when
stored for extended periods of time).
Inert to almost all inorganic and organic compounds except F2; porosity to
gases is significantly less than that of polyethylene; safe to use below 250 EC
but decomposes at 300 EC; difficulty in shaping containers results in high
cost; low thermal conductivity (requires long periods of time to heat samples).

Teflon™

Corrosive
applications

Polystyrene

Sample and reagent Only useful for acid solutions < 0.1 M; brittle
storage

The internal surface area of a container, whether used for sample preparation or storage, may
cause loss of analyte. Scratches and abrasions increase the surface area, and their geometry make
loss of analyte likely. Thus, it is important to discard containers that are scratched or abraded on
their interior surfaces.
12.2.2 Contamination from Sources in the Laboratory
Contamination leads to biased data that misrepresent the concentration or presence of
radionuclides in a specific sample. Therefore, laboratory personnel should take appropriate
measures to prevent the contamination of samples. Such precautions are most important when
multiple samples are processed together. Possible sources of contamination include:
•
•
•
•
•

Airborne;
Reagents (tracers are discussed in Chapter 14);
Glassware/equipment;
Facilities; and
Cross-contamination between high- and low-activity samples.

The laboratory should use techniques that eliminate air particulates or the introduction of any
outside material (such as leaks from aerosols) into samples and that safeguard against using
contaminated glassware or laboratory equipment. Contamination of samples can be controlled by
adhering to established procedures for equipment preparation and decontamination before and
after each sample is prepared. Additionally, the results of blank samples (e.g., sand), which are
run as part of the internal quality assurance program, should be closely monitored, particularly
following the processing of samples with elevated activity.
“Cross-contamination” is the contamination of one sample by another sample that is being
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processed concurrently or that was processed prior to the current sample leaving a residue on the
equipment being used. Simply keeping samples covered whenever practical is one technique to
minimize cross-contamination. Another technique is to order the processing of samples
beginning with the lowest contamination samples first. It is not always possible to know the
exact rank of samples, but historical or field screening data may be useful.
Laboratory personnel should be wary of using the same equipment (gloves, tweezers for filters,
contamination control mats, etc.) for multiple samples. Countertops and other preparation areas
should be routinely monitored for contamination.
12.2.2.1

Airborne Contamination

Airborne contamination is most likely to occur when grinding or pulverizing solid samples. Very
small particles (~10 µm) may be produced, suspended in air, and transported in the air before
settling onto a surface. Other sources of potential airborne contamination include samples that
already consist of very small particles, volatile radionuclides (including tritium), or radionuclides
that decay through a gaseous intermediate (i.e., 226Ra decays to 222Rn gas and eventually decays to
210
Pb). Therefore, the grinding or pulverizing of solid samples or the handling of samples that
could produce airborne contamination should be carried out under a laboratory hood or ventilated
enclosure designed to prevent dispersal or deposition in the laboratory of contaminated air
particulates. These particles easily can contaminate other samples stored in the area. To prevent
such cross-contamination, other samples should be covered or removed from the area while
potential sources of airborne contamination are being processed.
If contamination from the ambient progeny of 222Rn is a concern, it can be avoided by refraining
from the use of suction filtration in chemical procedures, prefiltering of room air (Lucas, 1967),
and use of radon traps (Lucas, 1963; Sedlet, 1966). The laboratory may have background levels
of radon progeny from natural sources in soil or possibly in its construction materials.
12.2.2.2

Contamination of Reagents

Contamination from radiochemical impurities in reagents is especially troublesome in low-level
work (Wang et al., 1975). Care must be taken in obtaining reagents with the lowest contamination possible. Due to the ubiquitous nature of uranium and thorium, they and their progeny are
frequently encountered in analytical reagents. For example, Yamamoto et al. (1989) found
significant 226Ra contamination in common barium and calcium reagents. Other problematic
reagents include the rare earths (especially cerium salts), cesium salts that may contain 40K or
87
Rb, and potassium salts. Precipitating agents such as tetraphenyl borates and chloroplatinates
may also suffer from contamination problems. In certain chemical procedures, it is necessary to
replace stable carriers of the element of interest with isotopes of another element when it is
difficult to obtain the stable carrier in a contamination-free condition. Devoe (1961) has written
an extensive review article on the radiochemical contamination of analytical reagents.
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12.2.2.3

Contamination of Glassware and Equipment

Other general considerations in sample preparation include the cleaning of glassware and
equipment (Section 12.2.3). Criteria established in the planning documents or laboratory SOPs
should give guidance on proper care of glassware and equipment (i.e., scratched glassware
increases the likelihood of sample contamination and losses due to larger surface area).
Glassware should be routinely inspected for scratches, cracks, etc., and discarded if damaged.
Blanks and screening should be used to monitor for contamination of glassware.
Whenever possible, the use of new or disposable containers or labware is recommended. For
example, disposable weigh boats can be used to prevent contamination of a balance. Disposable
plastic centrifuge tubes are often less expensive to use than glass tubes that require cleaning after
every use. If non-disposable containers or labware are used, it may be necessary to use new
materials for each new project to reduce the potential for contamination. Blanks can be used to
detect cross-contamination. Periodic rinsing with a dilute solution of nitric acid can aid in
maintaining clean glassware. However, Bernabee et al. (1980) could not easily remove nuclides
sorbed onto the walls of plastic containers by washing with strong mineral acids. They report that
nuclides can be wiped from the walls, showing the importance of the physical action of a brush
to the cleaning process.
12.2.2.4

Contamination of Facilities

In order to avoid contamination of laboratory facilities and possible contamination of samples or
personnel, good laboratory practices must be constantly followed, and the laboratory must be
kept in clean condition. The laboratory should establish and maintain a Laboratory Contamination Control Program (Section 12.2.4) to avoid contamination of facilities and to deal with it
expeditiously if it occurs. Such a program should address possible samples of varying activity or
characteristics. This minimizes sample cross-contamination through laboratory processing
equipment (e.g. filtering devices, glassware, ovens, etc).
12.2.3 Cleaning of Labware, Glassware, and Equipment
12.2.3.1

Labware and Glassware

Some labware is too expensive to be used only once (e.g., crucibles, Teflon™ beakers, separatory
funnels). Labware that will be used for more than one sample should be subjected to thorough
cleaning between uses. A typical cleaning protocol includes a detergent wash, an acid soak (HCl,
HNO3, or citric acid), and a rinse with deionized or distilled water. As noted in Chapter 10,
scrubbing glassware with a brush aids in removing contaminants.
The Chemical Technician’s Ready Reference Handbook (Shugar and Ballinger, 1996) offers
practical advice on washing and cleaning laboratory glassware:
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• Always clean your apparatus immediately after use. It is much easier to clean the glassware
before the residues become dry and hard. If dirty glassware cannot be washed immediately, it
should be left in water to soak.
• Thoroughly rinse all soap or other cleaning agent residue after washing glassware to prevent
possible contamination. If the surface is clean, the water will wet the surface uniformly; if the
glassware is still soiled, the water will stand in droplets.
• Use brushes carefully and be certain that the brush has no exposed sharp metal points that can
scratch the glass. Scratched glassware increases the likelihood of sample contamination and
losses due to larger surface areas. Moreover, scratched glassware is more easily broken,
especially when heated.
Automatic laboratory dishwashers and ultrasound or ultrasonic cleaners are also used in many
radiochemical laboratories. It is important to note that cleaning labware in an automatic
laboratory dishwasher alone may not provide adequate decontamination. Contaminated glassware
may need to be soaked in acid or detergent to ensure complete decontamination. Ultrasonic
cleaning in an immersion tank is an exceptionally thorough process that rapidly and efficiently
cleans the external, as well as the internal, surfaces of glassware or equipment. Ultrasonic
cleaners generate high-frequency sound waves and work on the principle of cavitation, which is
the formation and collapse of submicron bubbles. These bubbles form and collapse about 25,000
times each second with a violent microscopic intensity that produces a scrubbing action (Shugar
and Ballinger, 1996). This action effectively treats every surface of the labware because it is
immersed in the solution and the sound energy penetrates wherever the solution reaches.
EPA (1992) contains a table of glassware cleaning and drying procedures for the various methods
given in the manual (including methods for the analysis of radionuclides in water). The suggested
procedure for cleaning glassware for metals analysis is to wash with detergent, rinse with tap
water, soak for 4 hours in 20 percent (by volume) HNO3 or dilute HNO3 (8 percent)/HCl (17
percent), rinse with reagent water, then air dry. Shugar and Ballinger (1996) suggest treating
acid-washed glassware by soaking it in a solution containing 2 percent NaOH and 1 percent
disodium ethylenediamine tetraacetate for 2 hours, followed by a number of rinses with distilled
water to remove metal contaminants.
More specifically to radionuclides, in their paper discussing the simultaneous determination of
alpha-emitting nuclides in soil, Sill et al. (1974) examined the decontamination of certain
radionuclides from common labware and glassware:
By far the most serious source of contamination is the cell, electrode, and “O” ring used
in the electrodeposition step. Brief rinsing with a strong solution of hydrochloric acid
containing hydrofluoric acid and peroxide at room temperature was totally ineffective in
producing adequate decontamination. Boiling anode and cell with concentrated nitric acid
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for 10 to 15 minutes removed virtually all of the activity resulting from the analysis of
samples containing less than 500 disintegrations per minute (dpm). When larger
quantities of activity such as the 2.5×104 counts per minute (cpm) used in the material
studies ... had been used, a second boiling with clean acid was generally required.
However, boiling nitric acid precipitates polonium and other procedures have to be used
in its presence. When such high levels of activity have been used, a blank should be run
to ensure that decontamination was adequate before the system is permitted to be used in
the analysis of subsequent low-level samples. Prudence suggests that a separate system
should be reserved for low-level samples and good management exercised over the level
of samples permitted in the low-level system to minimize the number of blanks and fulllength counting times required to determine adequate decontamination.
...Beakers, flasks, and centrifuge tubes in which barium sulfate has been precipitated must
be cleaned by some agent known to dissolve barium sulfate, such as boiling perchloric or
sulfuric acids or boiling alkaline DTPA [diethylenetriaminepentacetate]. This is a
particularly important potential source of contamination, particularly if hot solutions
containing freshly-precipitated barium sulfate are allowed to cool without stirring. Some
barium sulfate post-precipitates after cooling and adheres to the walls so tenaciously that
chemical removal is required. Obviously, the barium sulfate will contain whichever
actinide is present, and will not dissolve even in solutions containing hydrofluoric acid.
Beakers or flasks in which radionuclides have been evaporated to dryness will invariably
contain residual activity which generally requires a pyrosulfate fusion to clean completely
and reliably. Separatory funnels can generally be cleaned adequately by rinsing them with
ethanol and water to remove the organic solvent, and then with hydrochloric-hydrofluoric
acids and water to remove traces of hydrolyzed radionuclides...
However, one should note that current laboratory safety guidelines discourage the use of
perchloric acid (Schilt, 1979).
12.2.3.2

Equipment

In order to avoid cross-contamination, grinders, sieves, mixers and other equipment should be
cleaned before using them for a new sample. Additional cleaning of equipment prior to use is
only necessary if the equipment has not been used for some time. The procedure can be as simple
or as complicated as the analytical objectives warrant as illustrated by Obenhauf et al. (2001). In
some applications, simply wiping down the equipment with ethanol may suffice. Another
practical approach is to brush out the container, and briefly process an expendable portion of the
next sample and discard it. For more thorough cleaning, one may process one or more batches of
pure quartz sand through the piece of solid processing equipment, and then wash it carefully. The
efficacy of the decontamination is determined by monitoring this sand for radionuclide
contamination.
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An effective cleaning procedure for most grinding containers is to grind pure quartz sand
together with hot water and detergent, then to rinse and dry the container. This approach
incorporates a safety advantage in that it controls respirable airborne dusts. It is important to note
that grinding containers become more difficult to clean with age because of progressive pitting
and scratching of the grinding surface. Hardened steel containers can also rust, and therefore
should be dried thoroughly after cleaning and stored in a plastic bag containing a desiccating
agent. If rust does occur, the iron oxide coating can be removed by a warm dilute oxalic acid
solution or by abrasive cleaning.
12.2.4 Laboratory Contamination Control Program
The laboratory should establish a general program to prevent the contamination of samples.
Included in the program should be ways to detect contamination from any source during the
sample preparation steps if contamination of samples occurs. The laboratory contamination
control program should also provide the means to correct procedures to eliminate or reduce any
source of contamination. Some general aspects of a control program include:
• Appropriate engineering controls, such as ventilation, shielding, etc., should be in place.
• The laboratory should be kept clean and good laboratory practices should be followed.
Personnel should be well-trained in the safe handling of radioactive materials.
• Counter tops and equipment should be cleaned and decontaminated following spills of
liquids or dispersal of finely powdered solids. Plastic-backed absorbent benchtop coverings
or trays help to contain spills.
• There should be an active health physics program that includes frequent monitoring of
facilities and personnel.
• Wastes should be stored properly and not allowed to accumulate in the laboratory working
area. Satellite accumulation areas should be monitored.
• Personnel should be mindful of the use of proper personnel protection equipment and
practices (e.g., habitual use of lab coats, frequent glove changes, routine hand washing).
• Operations should be segregated according to activity level. Separate equipment and facilities
should be used for elevated and low-level samples whenever possible.
• SOPs describing decontamination and monitoring of labware, glassware, and equipment
should be available.

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• Concentrated standard stock solutions should be kept isolated from the general laboratory
working areas.
As an example, Kralian et al. (1990) have published the guidelines for effective low-level
contamination control.

12.3 Solid Samples
This section discusses laboratory preparation procedures for solid samples as illustrated in
Figure 12.2. General procedures such as exclusion of unwanted material in the sample; drying,
charring, and ashing of samples; obtaining a constant weight (if required); and homogenization
are discussed first. Examples of preparative procedures for solid samples are then presented.
Solid samples may consist of a wide variety of materials, including:
• Soil and sediment;
• Biota (plants and animals); and
• Other materials (metal, concrete, asphalt, solid waste, etc.).
Before a solid sample is prepared, the specific procedures given in the planning documents
should be reviewed. This review should result in a decision that indicates whether materials other
than those in the intended matrix should be removed, discarded, or analyzed separately. Any
material removed from the sample should be identified, weighed, and documented.
To ensure that a representative aliquant of a sample is analyzed, the sample should first be dried
or ashed and then blended or ground thoroughly (Section 12.3.1.4 and Appendix F, Laboratory
Subsampling). Homogenization should result in a uniform distribution of analytes and particles
throughout the sample. The size of the particles that make up the sample will have a bearing on
the representativeness of each aliquant.
12.3.1 General Procedures
The following sections discuss the general procedures for exclusion of material, heating solid
samples (drying, charring, and ashing), obtaining a constant weight, mechanical manipulation
grinding, sieving, and mixing), and subsampling. Not every step is done for all solid sample
categories (soil/sediment, biota, and other) but are presented here to illustrate the steps that could
be taken during preparation.

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FIGURE 12.2—Laboratory sample preparation flowchart (for solid samples)

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12.3.1.1

Exclusion of Material

EXCLUSION OF MATERIAL BY SIZE AND COMPOSITION
During solid preparation, some particles may be identified in the sample that are not a part of the
matrix intended for analysis. Examples of such particles are rocks and pebbles or fragments of
glass and plastic. Depending on the specific procedures given in the planning documents on the
constitution of the sample taken, rocks and pebbles can be removed and analyzed separately if
desired. The sample should be weighed before and after any material is removed. Other materials
that are not a part of the required matrix can also be removed and analyzed separately. If analysis
of the material removed is necessary, applicable SOPs should be used to prepare the material for
analysis.
EXCLUSION OF ORGANIC MATERIAL
Leaves, twigs, and grass can easily be collected inadvertently along with samples of soil or
sediment. Because these are not usually intended for analysis, they are often removed and stored
for future analysis, if necessary. The material removed should be identified, if possible, and
weighed.
12.3.1.2

Principles of Heating Techniques for Sample Pretreatment

Applying elevated temperatures during sample preparation is a widely used technique for the
following reasons:
• To remove moisture or evaporate liquids, raise the temperatures to 60 to 110 EC, which will
not significantly alter the physical composition of the sample.
• To prepare a sample containing organic material for subsequent wet ashing or fusion, “char”
the material by heating to medium temperature of 300 to 350 EC (see page 12-19 on
“Charring of Samples”).
• To prepare the sample for subsequent determination of nonvolatile constituents, dry ash at
high temperature of 450 to 750 EC. This may significantly change the physical and chemical
properties of the sample.
Once a decision is made to use elevated temperatures during sample preparation, several
questions should be considered:
• What material should be used for the sample container?
• What should serve as the heat source?
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• How quickly should the temperature be raised? (Rate of stepwise temperature increase)
• What is the maximum temperature to which the sample should be exposed?
• How long should the sample be heated at the maximum temperature?
• How quickly should the sample be cooled afterward?
The following sections provide information related to these questions.
Note that there are times during sample preparation when samples should not be heated. For
example, samples to be prepared for 3H or 14C determination should not be heated. Since 3H is
normally present as tritiated water in environmental samples, heating will remove the 3H.
Similarly, 14C is usually present in environmental samples as carbonates or 14CO2 dissolved in
water, and heating will release 14C as a gas. Samples to be analyzed for iodine, mercury,
antimony, or other volatile elements should be heated only under conditions specified in the
planning documents. If both volatile and nonvolatile elements are determined from the same
sample, aliquants of the original sample should be removed for determination of the volatile
elements.
Ovens, furnaces, heat lamps, and hot plates are the traditional means to achieve elevated
temperatures in the laboratory. However, more recently, microwave ovens have added an
additional tool for elevating temperature during sample preparation. Walter et al. (1997) and
Kingston and Jassie (1988) give an overview of the diverse field of microwave-assisted sample
preparation. A dynamic database of research articles related to this topic can be found at the
SamplePrep Web™ at www.sampleprep.duq.edu/index.html. As microwave sample preparation
has developed, numerous standard methods with microwave assistance have been approved by
the American Society for Testing and Materials (ASTM), Association of Official Analytical
Chemists (AOAC), and the U.S. Environmental Protection Agency (EPA). The majority of the
microwave-assisted methods are for acid-dissolution (Chapter 13), but several are for drying
samples.
Alternatives to heating samples include drying them slowly in a vacuum desiccator, air-drying, or
freeze-drying. ASTM D3974 describes three methods of preparing soils, bottom sediments,
suspended sediments, and waterborne materials: (1) freeze-drying; (2) air-drying at room
temperature; and (3) accelerated air-drying.
DRYING SAMPLES
It must be determined at the start of an analytical procedure if the results are to be reported on an
as-received or dry-weight basis. Most analytical results for solid samples should be reported on a
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dry-weight basis, which denotes material dried at a specified temperature to a constant weight or
corrected through a “moisture” determination made on an aliquant of the sample taken at the
same time as the aliquant taken for sample analysis.
Typically, samples are dried at temperatures of 105 to 110 EC. Sometimes it is difficult to obtain
constant weight at these temperatures, then higher temperatures must carefully be used.
Alternatively, for samples that are extremely heat sensitive and decompose readily, vacuum
desiccation or freeze-drying techniques are applicable.
The presence of water in a sample is a common problem frequently facing the analyst. Water
may be present as a contaminant (i.e., from the atmosphere or from the solution in which the
substance was formed) or be bonded as a chemical compound (i.e., a hydrate). Regardless of its
origin, water plays a role in the composition of the sample. Unfortunately, especially in the case
of solids, water content is variable and depends upon such things as humidity, temperature, and
the state of subdivision. Therefore, the make-up of a sample may change significantly with the
environment and the method of handling.
Traditionally, chemists distinguish several ways in which water is held by a solid (Dean, 1995).
• Essential water is an integral part of the molecular or crystal structure and is present in
stoichiometric quantities, for example, CaC2O4·2H2O.
• Water of constitution is not present as such in the solid, but is formed as a product when the
solid undergoes decomposition, usually as a result of heating. For example, Ca(OH)2 6 CaO
+ H2O.
• Nonessential water is retained by physical forces, is non-stoichiometric, and is not necessary
for the characterization of the chemical composition of the sample.
• Adsorbed water is retained on the surface of solids in contact with a moist environment, and
therefore, is dependent upon the humidity, temperature, and surface area of the solid.
• Sorbed water is encountered with many colloidal substances such as starch, charcoal, zeolite
minerals, and silica gel and may amount to as much as 20 percent or more of the solid.
Sorbed water is held as a condensed phase in the interstices or capillaries of the colloid and it
is greatly dependent upon temperature and humidity.
• Occluded water is entrapped in microscopic pockets spaced irregularly throughout solid
crystals. These cavities frequently occur naturally in minerals and rocks.

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• Water also may be present as a solid solution in which the water molecules are distributed
homogeneously throughout the solid. For example, natural glasses may contain several
percent moisture in this form.
Heat Source. There are several choices when heating to dryness. The heat source is often
determined by the amount of time available for drying and the potential for the sample to spatter
or splash during drying. When time is not a primary concern and there is little or no chance of
sample cross-contamination, samples are heated uncovered in a drying oven at the minimum
temperature needed to remove moisture. If time is of concern, samples with high moisture
content usually can be dried or evaporated faster using a hot plate. Heating on a hot plate
significantly increases the chance of cross-contamination by spattering or splashing during
boiling. However, ribbed watch glasses, which cover the sample yet still allow for evaporation,
can be used to minimize cross-contamination in this approach. Samples may also be placed under
a heat lamp. This method reduces the risk of cross-contamination by applying heat to the surface
where vaporization occurs, minimizing splashing during boiling. However, the elevated
temperature is difficult to measure or control, and spattering still may be a problem when the
sample reaches dryness.
Microwave systems may also be used to dry samples. ASTM E1358 and ASTM D4643 use
microwave energy to dry either wood or soil to a constant weight. In a similar fashion, AOAC
Official Methods 985.14 and 985.26 use microwave energy to dry fat from meat or water from
tomato juice. Other examples include Beary (1988), who has compared microwave drying to
conventional techniques using solid standards from the National Institute of Standards and
Technology (coal, clays, limestone, sediment) and foods and food materials (rice and wheat
flour), and Koh (1980) who discusses microwave drying of biological materials.
Container Material. A sample container’s composition typically poses no problem. Borosilicate
glass is generally recommended because it is inexpensive, transparent, reusable, and has good
thermal properties. Platinum, Teflon™ (polytetrafluoroethylene—PTFE), porcelain, or aluminum
foil containers are acceptable and may be preferable in certain situations. Polyethylene and other
plastics of low melting point are only useful in hot water baths or ovens where the temperature is
closely monitored. Polyethylene is affected by heat applied directly to the container. The
properties of several common materials used for sample containers are presented in Table 12.2
(on page 12-5). Note that the sample containers commonly received from the field will be those
suitable for bulk samples rather than containers used during sample preparation. The plan will
identify the type of container material to be used for field activities for samples to be shipped to
the laboratory and the type of container material to be used during the various steps of sample
preparation.
Heating Rate. The heating rate is generally not considered when removing moisture, because the
maximum temperature typically is very low (60 to 110 EC). Samples simply are placed inside the
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preset oven. Hot plates may be preheated to the desired temperature before heating the sample or
turned on and gradually heated with the sample in place.
Maximum Temperature. The maximum temperature used for drying samples typically is just
above the boiling point of water—105 to 110 EC. Higher temperatures will not dry the samples
significantly faster and may result in accidents or cross-contamination due to uneven heating.
Lower temperatures will not reduce the chance of cross-contamination, but will significantly
increase the drying time. One exception to this rule occurs when the physical form of the sample
needs to be preserved. Many minerals and chemicals have waters of hydration that affect the
structure and may also affect the chemical and physical properties. Samples heated at 60 EC will
retain the waters of hydration in most chemicals and minerals and still provide dry samples in a
reasonable period of time (e.g., 12 to 15 hrs.).
Time. The duration a sample is heated to remove moisture depends on the size of the sample, the
amount of moisture in the sample, the air flow around the sample, and the temperature applied to
the sample. If heating the sample is to provide a constant dry weight, it is more difficult to
determine how long to heat the sample. One convenient approach, especially when working with
numerous samples, is to dry all materials overnight, or occasionally longer. This amount of
heating is usually more than sufficient for drying samples for radiochemical analysis. If time is a
critical factor or if a quantitative assessment of the uncertainty in the sample weight is required
by the planning documents, the sample can be subjected to repeated cycles of drying and
weighing until a series of weights meet the specified requirements (Section 12.3.1.3). For
example, one such requirement might be to obtain three consecutive weights with a standard
deviation less than 5 percent of the mean. While repeated cycles of drying and weighing can
provide a quantitative measure of the uncertainty in the sample weight over time, a single weight
after an overnight drying cycle typically provides a similar qualitative level of confidence with
significantly less working time. Another time-saving step is to use microwave techniques rather
than conventional heating sources during sample preparation (ANL/ACL, 1992; Walter et al.,
1997).
Alternatives to Heating. (1) Vacuum-desiccation. A desiccator is a glass or aluminum container
that is filled with a substance that absorbs water, a “desiccant.” The desiccator provides a dry
atmosphere for objects and substances. Dried materials are stored in desiccators while cooling in
order to minimize the uptake of ambient moisture. The ground-glass or metal rim of the desiccator should be greased lightly with petroleum jelly or silicone grease to improve performance.
Calcium sulfate, sodium hydroxide, potassium hydroxide, and silica gel are a few of the common
desiccants. The desiccant must be renewed frequently to keep it effective. Surface caking is a
signal to renew or replace the desiccant. Some desiccants contain a dye that changes color upon
exhaustion.
Vacuum desiccators are equipped with a side-arm so that they may be connected to a vacuum to
aid in drying. The contents of the sealed evacuated desiccator are maintained in a dry, reducedMARLAP

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pressure atmosphere. Care must be exercised when applying a vacuum as a rapid pressure
reduction, for high water content samples can result in “boiling” with subsequent sample loss and
potential cross-contamination. The release of vacuum should be accomplished by the slow
introduction of dry or ambient-humidity air into the chamber.
(2) Freeze-drying. Certain substances (i.e., biological materials, pharmaceuticals), which are
extremely heat sensitive and cannot be dried at atmospheric conditions, can be freeze-dried
(Cameron and Murgatroyd, 1996). Freeze-drying, also known as “lyophilization,” is the process
by which substances are frozen, then subjected to high vacuum. Under these conditions, ice
(water) sublimes and other volatile liquids are removed. The non-sublimable material is left
behind in a dry state.
To freeze-dry effectively, dilute solutions are used. In order to increase the surface area, the
material is spread out on the inner surface of the container as it is frozen. Once the solution or
substance to be dried is frozen solid, the primary drying stage begins in which a high vacuum is
applied, and the ice sublimes, desorbing the free ice and some of the bound moisture. During
secondary drying, a prolonged drying stage, the sorbed water that was bound strongly to the
solids is converted to vapor. This can be a slow process, because the remaining bound water has
a lower pressure than the free liquid at the same temperature, making it more difficult to remove.
Secondary drying actually begins during the primary drying phase, but it must be extended after
the total removal of free ice to achieve low levels of residual moisture.
Commercial freeze-drying units are self-contained. Simple units consist of a vacuum pump,
adequate vapor traps, and a receptacle for the material to be dried. More sophisticated models
include refrigeration units to chill the solutions, instrumentation to designate temperature and
pressure, heat and cold controls, and vacuum-release valves. The vacuum pump should be
protected from water with a dry-ice trap and from corrosive gases with chemical gas-washing
towers.
CHARRING OF SAMPLES TO PARTIALLY OXIDIZE ORGANIC MATERIAL
Heating samples at a moderate temperature (300 to 350 EC) is sometimes used as a method of
preparing a sample for subsequent decomposition using wet ashing or fusion techniques. Large
amounts of organic material can react violently or even explosively during decomposition.
Heating the sample to partially oxidize—or “char”—the organic material may limit reactivity
during subsequent preparation.
Heat Source. Heat lamps, muffle furnaces, or hot plates may be used as a heat source for charring
samples. Heat lamps are often selected because they can also be used to dry the sample before
charring. Once dried, the sample can be moved closer to the lamp to raise the temperature and
char the sample (confirmed by visual inspection). Heat lamps also reduce the potential for crosscontamination by minimizing spattering and splashing. Hot plates can be used similarly to heat
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lamps. The sample is dried and the temperature is raised to char the sample; however, hot plates
increase the probability of spattering and splashing. Muffle furnaces can be used when the
charring is performed as part of dry ashing instead of part of the drying process. In this case, the
muffle furnace temperature is first raised slowly.
Sample Container. The choice of sample container depends primarily on the next step in the
sample preparation process. When dry ashing or fusing, the sample container will usually be a
platinum or porcelain crucible. Zirconium or nickel crucibles may also be used. If the sample will
be dissolved using wet ashing techniques, the container may be borosilicate glass or a platinum
crucible. Care should be taken to prevent ignition of samples in glass containers. Ignited samples
may burn at temperatures high enough to cause damage to the container and loss of sample.
Polyethylene and Teflon™ generally are not acceptable because of the increased temperature and
risk of melting the container.
Heating Rate. Heating rate becomes a concern when charring samples because of the increased
temperatures. The general rule is to raise the temperature slowly to heat the sample evenly and
prevent large increases in temperature within the sample, which could lead to ignition. Typically,
a rate of 50 to 100 EC per hour is considered appropriate. Samples containing large quantities of
organic material may require slower heating rates.
Maximum Temperature. One of the primary goals of charring a sample is to oxidize the materials
slowly and gently. Gentle oxidation is accomplished by slowly raising the temperature close to
the ignition point and letting the sample smolder. Most organic compounds will char and
decompose in the range of 300 to 350 EC, so this is usually the range of temperatures where
charring takes place. Ignition results in rapid oxidation accompanied by large volumes of
released gases and potential sample loss. This reaction can raise the temperature of the sample to
several hundred degrees above the desired maximum and result in significant losses during offgassing. The progress of the reaction can be monitored visually by observing the volume of gas
or smoke released. Thin wisps of smoke are usually allowable; clouds of smoke and flames are
not. Visual inspection is easily accomplished when hot plates or heat lamps are used as heat
sources. Some muffle furnaces are fitted with viewing windows to allow visual inspection. Never
open a muffle furnace just to check on the progress of a reaction. This will cause a sudden
change in temperature, increase the oxygen level and possibly ignite the sample, and disrupt air
currents within the furnace to increase potential sample loss.
Time. The duration required to char a sample depends on the sample size, the amount of organic
material in the sample, the ignition point of the organic material, the temperature of the sample,
and the oxygen supply. Samples usually are heated until smoke begins to appear and allowed to
remain at that temperature until no more smoke is evident. This process is repeated until the
temperature is increased and no more smoke appears. Charring samples may require a significant
amount of time and effort to complete. The duration may be reduced by improving the flow of air
to the sample or mixing HNO3 or nitrate salts with the sample before drying. However, this
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approach is recommended only for well-characterized samples, those previously evaluated for the
applicability of this technique, because nitrated organic compounds can oxidize in a violent or
explosive manner.
DRY ASHING SAMPLES
The object of dry ashing is to combust all of the organic material and to prepare the sample for
subsequent treatment using wet ashing or fusion techniques. This procedure involves heating a
sample in an open dish or crucible in air, usually in a muffle furnace to control the temperature
and flow of air. Microwave techniques are also available for dry ashing samples.
Dry ashing is used to determine ash weight as well as nonvolatile constituents. The associated
chemistry is very complex, with oxidizing and reducing conditions varying throughout the
sample and over time. During the combustion process, temperatures in the sample may reach
several hundred degrees above the desired temperature, particularly if there is good air flow at
the beginning of the ashing process (Bock, 1979). Covering samples during heating is not
recommended, especially when using platinum crucibles. The lack of air produces a reducing
atmosphere that results in reduction of metals that alloy with the crucible (Table 12.2 on page 125). This reaction results in loss of sample and potential for contamination of subsequent samples
when using the same crucible.
Heat Source. The traditional heat sources for dry ashing are muffle furnaces or burner flames.
Electronic muffle furnaces are recommended for all heating of platinum crucibles because
burners produce significant levels of hydrogen gas during combustion, and platinum is permeable
to hydrogen gas at elevated temperatures. Hydrogen gas acts as a reducing agent that can result in
trace metals becoming alloyed to the platinum.
Microwave ovens have also proved to be quick and efficient when dry ashing plant tissue
samples, with results comparable to conventional resistance muffle furnaces (Zhang and Dotson,
1998). The microwave units are fitted with ashing blocks (a ceramic insert) that absorb
microwave energy and quickly heats to high temperatures. This, in combination with the
microwave energy absorbed directly by the sample, allows for rapid dry ashing of most materials.
The units are designed for increased air flow that further accelerates combustion of the samples.
Sample Container. Platinum, zirconium, or porcelain are usually used to form crucibles for dry
ashing. Nickel may also be appropriate for some applications (Table 12.2). Platinum generally is
recommended when available and is essentially inert and virtually unaffected by most acids.
Zirconium and porcelain crucibles are resistant to most acids, are more resistant to HCl, and are
significantly less expensive than platinum. Glass and plastic containers should not be used for
dry ashing because the elevated temperatures exceed the melting point of these materials.

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Crucibles fabricated from ceramic, graphite, and platinum can be used in microwave applications. Quartz fiber crucibles can accelerate the ashing process since this material rapidly cools
and allows many sample types to be reweighed in 60 seconds or less after removal from the
microwave unit.
Heating Rate. Samples should be dried before dry ashing and placed in an unheated furnace;
then, the furnace temperature is gradually increased. The sample should be spread as thinly and
evenly as possible on the bottom of the container to allow for its equal heating. To ensure even
heating of the sample and to minimize the chance of ignition, the temperature of the furnace is
raised slowly. If the sample was previously charred, a rate of approximately 100 EC per hour is
typical. This rate is slow enough that small amounts of organic material or water can be removed
from the sample without violent reactions. If the sample is not charred and contains a significant
amount of organic material, a slower rate may be necessary to control the oxidation of organic
material.
Maximum Temperature. The maximum temperature is determined by the sample matrix and the
volatility of the elements to be analyzed. Generally, the temperature should be as low as possible
to reduce the loss of volatile compounds, but high enough to ensure complete combustion of the
sample. A minimum temperature of 450 EC is often used to ensure complete combustion (Bock,
1979). The upper limit for dry ashing is usually determined by the sample container and the
elements being analyzed and is generally considered to be 750 EC, but sample-specific conditions
may use temperatures up to 1,100 EC. However, in practice, some components that are normally
considered to be nonvolatile may be lost at temperatures above 650 EC (Bock, 1979). Ashing
aids may be added to samples to accelerate oxidation, prevent volatilization of specific elements,
and prevent reaction between the sample and the container. Examples include adding nitrate
before drying to assist oxidation and loosen the ash during combustion, adding sulfate to prevent
volatilization of chlorides (e.g., PbCl2, CdCl2, NaCl) by converting them to the higher boiling
sulfates, and adding alkaline earth hydroxides or carbonates to prevent losses of anions (e.g., Cl-,
As-3, P-3, B). Table 12.3 lists dry ashing procedures using a platinum container material for
several elements commonly determined by radiochemical techniques.
Time. The duration required to completely combust a sample depends on the size of the sample,
the chemical and physical form of the sample before and after ashing, and the maximum
temperature required to ash the sample. In many cases, it is convenient to place the sample in an
unheated furnace and gradually raise the temperature during the day until the maximum
temperature is achieved. The furnace is then left at the maximum temperature overnight (12
hours). The furnace is allowed to cool during the next day, and samples are removed from a cold
oven. This procedure helps prevent sudden changes in temperature that could cause air currents
that may potentially disturb the ash. An alternative is to leave the sample at maximum
temperature for 24 hours and let the sample cool in the oven the second night to ensure complete
combustion of the sample.
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The elapsed time for dry ashing samples can be significant (greater than 36 hours), but the actual
time required by laboratory personnel is minimal.
TABLE 12.3 — Examples of dry-ashing temperatures (platinum container)
Element

Temperature/Matrix

Cobalt

450–600 EC for biological material; some losses reported due to reactions with crucible; increased
volume of sample increases volume of ash and limits loss of sample.

Cesium

400–450 EC for food and biological material; CsCl and CsNO3 begin to volatilize when held at
temperatures above 500 EC for any length of time.

Iodine

450–500 EC with an alkaline ashing aid to prevent volatilization; losses reported for temperatures as
low as 450 EC even with alkaline ashing aids added; total volatilization >600 EC.

Lead

450–500 EC acceptable for most samples; bone or coal (lead phosphate) may be ashed as high as
900 EC without significant losses; PbO2 reacts with silica in porcelain glaze at low temperatures;
PbCl2 is relatively volatile and nitrate or sulfate ashing aids have been used to good effect.

Plutonium

450 EC with nitric acid ashing aid for biological material, 550 EC for dust on air filters, 700 EC for
soil; high temperature leads to adsorption onto carbon particles and incomplete dissolution of ash.

Strontium

450–550 EC for plants, 600 EC for meat, 700 EC for milk and bone.

Technetium 725–750 EC for plants treated with ammonia.
Thorium

750 EC for bone.

Uranium

600 EC for coal, 750 EC for biological material; uranium reacts with porcelain glaze resulting in
sample losses.

Source: Bock (1979).
(Note that reducing conditions for platinum containers are given in Table 12.2)

12.3.1.3

Obtaining a Constant Weight

If required, constant weight is obtained by subjecting a sample to repetitive cycles of drying and
weighing until a series of weights meets specified requirements. Project-specific planning
documents or laboratory SOPs should define the acceptance criteria. For example, in Greenberg
et al. (1992), solids are repetitively heated for an hour, then weighed until successive weighings
agree within 4 percent of the mass or within 0.5 mg. In the ASTM guidelines for the preparation
of biological samples (ASTM D4638), an accurately weighed sample (1 to 2 g ± 0.1 mg, 5 to 10
g ± 1 mg, >10 g ± 10 mg) is heated for 2 hours, cooled in a desiccator, and weighed. Drying is
repeated at hourly intervals to attain a constant weight within the same accuracy. The consistent
drying of materials from a large sample set may require a qualitative evaluation of change in the
sample composition. If a qualitative change occurs the drying method may need to be checked for
completeness. One way to do this would be to perform routine dry-to-constant-weight
evaluations on separate samples.
Laboratory conditions and handling of the samples by the analyst during sample weight
determinations can increase the uncertainty of the final sample mass.

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12.3.1.4

Subsampling

Laboratories routinely receive larger samples than required for analysis. The challenge then
becomes to prepare a sample that is representative and large enough for analysis, but not so large
as to cause needless work in its final preparation. Generally, a raw sample first is crushed to a
reasonable particle size and a portion of the crushed material is taken for analysis. This step may
be repeated with intermittent sieving of the material until an appropriate sample size is obtained.
Then, this final portion is crushed to a size that minimizes sampling error and is fine enough for
the dissolution method (Dean 1995; Pitard, 1993).
French geologist Pierre Gy (1992) has developed a theory of particulate sampling that is
applicable to subsampling in the laboratory. Appendix F summarizes important aspects of the
theory and includes applications to radiochemistry. Some of the important points to remember
include the following:
• For most practical purposes, a subsample is guaranteed to be unbiased only if every particle
in the sample has the same probability of being selected for the subsample.
• The weight of the subsample should be many times greater than the weight of the largest
particle in the sample.
• The variance associated with subsampling may be reduced either by increasing the size of the
subsample or by reducing the particle sizes before subsampling.
• Grouping and segregation of particles tends to increase the subsampling variance.
• Grouping and segregation can be reduced by increment sampling, splitting, or mixing.
Increment sampling is a technique in which the subsample is formed from a number of smaller
portions selected from the sample. A subsample formed from many small increments will
generally be more representative than a subsample formed from only one increment. The more
increments the better. An example of increment sampling is the one-dimensional “Japanese slabcake” method (Appendix F, Laboratory Subsampling).
Splitting is a technique in which the sample is divided into a large number of equal-sized
portions and several portions are then recombined to form the subsample. Splitting may be
performed by a manual procedure, such as fractional shoveling, or by a mechanical device, such
as a riffle splitter. A riffle splitter consists of a series of chutes directed alternately to opposite
sides. The alternating chutes divide the sample into many portions, which are then recombined
into two. The riffle may be used repeatedly until the desired sample size is obtained. Riffle
splitters are normally used with free-flowing materials such as screened soils.
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Another traditional method for splitting is coning and quartering (Appendix F). Gy (1992) and
Pitard (1993) do not recommend coning and quartering because with similar tools and effort, one
can do fractional shoveling, which is a more reliable method.
If proper techniques and tools are used and adequate care is taken, samples of the sizes typically
encountered in the laboratory can be mixed effectively. However, the effects of mixing tend to be
short-lived because of the constant influence of gravity. Heterogeneous material may begin to
segregate immediately after mixing.
The method and duration needed to mix a sample adequately depends on the volume and type of
material to be mixed. Small volumes can be mixed by shaking for a relatively short time. Large
volumes may require hours. Pitard (1993) describes dynamic and discontinuous processes for
mixing samples including:
• Mechanical mixing of test tube samples is useful for small sample size and can be performed
on many samples at once. Some examples are a pipette shaker with a motor-activated,
rocking controlled motion; a nutator mixer with the test tubes fixed to an oscillating plate;
and a tube rotator where tubes are attached to a rotating plate mounted at an angle.
• Mechanical mixing of closed containers by rotating about a tumbling axis. A turbula
mechanical mixer is an example.
• Magnetic stirrers are commonly used to homogenize the contents of an open beaker.
• V-blenders are used to homogenize samples from several hundred grams to kilogram size.
• Stirrers coupled with propellers or paddles are used to mix large volumes of slurries or pulp.
• Sheet mixing or rolling technique, in which the sample is placed on a sheet of paper, cloth, or
other material, and the opposite corners are held while rolling the sample (see ASTM C702
for aggregates).
• Ball and rod mills homogenize as well as grind the sample (see ASTM C999 for soils).
When dealing with solid samples, it is often necessary to grind the sample to reduce the particle
size in order to ensure homogeneity and to facilitate attack by reagents. Obenauf et al. (2001) is
an excellent resource for information regarding grinding and blending.
For hand grinding, boron carbide mortars and pestles are recommended. For samples that can be
pulverized by impact at room temperature, a shatterbox, a mixer-mill, or a Wig-L-Bug™ is
appropriate, depending on the sample size. For brittle materials—such as wool, paper, dried
plants, wood, and soft rocks—which require shearing as well as impact, a hammer-cutter mill is
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warranted. For flexible or heat-sensitive samples such as polymers, cereal grains, and biological
materials, cryogenic grinding is necessary. Methods are described below:
• A shatterbox spins the sample, a puck, and a ring inside a dish-shaped grinding container in a
tight, high-speed horizontal circle. Within two to five minutes, approximately 100 grams of
brittle material can be reduced to less than 200 mesh. Shatterboxes are used typically to grind
soils, cement mix, rocks, slags, ceramics, and ores. They have also been used for hundreds of
other materials including dried marsh-grass, pharmaceuticals, fertilizers, and pesticides.
When used in a cryogenic atmosphere, this approach can be used to grind rubber, polymers,
bone, hair, and tissue.
• A mixer-mill grinds samples by placing them in a container along with one or more grinding
elements and imparting motion to the container. The containers are usually cylindrical, and
the grinding elements are ordinarily balls, but may be rods, cylinders or other shapes. As the
container is rolled, swung, vibrated or shaken, the inertia of the grinding elements causes
them to move independently into each other and against the container wall, thus, grinding the
sample. Mixer-mills are available for a wide-range of sample sizes. The length of time
necessary to grind a sample depends on the hardness of the material and the fineness desired
in the final product.
• The Wig-L-Bug™ is an example of a laboratory mill for pulverizing and blending very small
samples, typically in the range of 0.1 to 1 mL.
• A hammer-cutter mill uses high-speed revolving hammers and a serrated grinding chamber
lining to combine both shearing and impact. A slide at the bottom of the hopper feeds small
portions of the sample (up to 100 mL) into the grinding chamber. After the sample is
adequately pulverized, it passes through a perforated-steel screen at the bottom of the
grinding chamber and is then collected. With this approach, dried plants and roots, soils, coal
and peat, chemicals, and soft rocks all grind quickly with little sample loss.
• Many analytical samples—such as polymers, rubber, and tissues that are too flexible or
susceptible to degradation to be impact-ground at room temperature—can be embrittled by
chilling and then pulverized. Samples can be frozen and placed in a traditional grinder, or
alternatively, a freezer mill can be used. In a freezer mill, the grinding vial is immersed in
liquid nitrogen, and an alternating magnetic field shuttles a steel impactor against the ends of
the vial to pulverize the brittle material. Researchers at Los Alamos National Laboratory
developed a method of cryogenic grinding of samples to homogenize them and allow the
acquisition of a representative aliquant of the materials (LANL, 1996).
When samples agglomerate or “cake” during grinding, further particle size reduction is
suppressed. Caking can be caused from moisture, heat, static charge accumulation, the fusing of
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particles under pressure, etc. When it occurs, caking is a serious challenge. There are two main
approaches to this problem, slurry grinding and dry grinding.
• In slurry grinding, particles are suspended in solution during grinding. Water, alcohol, or
other liquids are added to the sample before grinding, and have to be removed afterwards.
Slurry grinding is a fairly reliable way of grinding a sample to micron-sized particles, but it is
sloppy and time-consuming.
• Dry grinding is often simpler and quicker, but requires careful matching of the technique to
the sample. If caking is due to moisture, as in many soils or cements, the sample should be
dried before grinding. Grinding aids such as lubricants, antistatic agents, abrasives, and
binding agents can also be used. Examples of grinding aids include dry soap or detergent (a
lubricant), graphite (an antistatic agent as well as a lubricant), polyvinyl alcohol, phenyl
acetate, propylene glycol, and aspirin. For example, propylene glycol (one drop for up to ten
grams of sample) is used for laboratory fine grinding of Portland cement and many minerals.
Grinding efficiency can be improved through intermittent screening of the material. The ground
sample is placed upon a wire or cloth sieve that passes particles of the desired size. The residual
particles are reground and this process is repeated until the entire sample passes through the
screen. Sieves with large openings can be used in the initial stages of sample preparation to
remove unwanted large rocks, sticks, etc.
The analysis of solid samples from the environment contaminated with radioactivity represents a
special challenge. In most cases, the radioactive materials will be from different sources than the
solid sample. Thus the contamination of solid samples with anthropogenic sources of radionuclides will result in a non-uniform particle mix as well as a non-uniform size distribution. This
further emphasizes the need for unbiased subsampling procedures.
12.3.2 Soil/Sediment Samples
For many studies, the majority of the solid samples will be soil/sediment samples or samples that
contain some soil. The definition of soil is given in Chapter 10 (Field and Sampling Issues that
Affect Laboratory Measurements). Size is used to distinguish between soils (consisting of sands,
silts, and clays) and gravels.
The procedures to be followed to process a raw soil sample to obtain a representative subsample
for analysis depend, to some extent, upon the size of the sample, the amount of processing
already undertaken in the field, and more importantly, the radionuclide of interest and the nature
of the contamination. Global fallout is relatively homogeneous in particle size and distribution in
the sample, and therefore, standard preparation procedures should be adequate for this
application. However, when sampling accidental or operational releases, the standard procedures
may be inadequate. Transuranic elements, especially plutonium, are notorious for being present
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as “hot-spots” ions (Eberhardt and Gilbert, 1980; Sill, 1975) and great care must be employed so
that the subsample taken for analysis accurately represents the total sample. This will depend on
the size and the degree of homogeneity. Multiple subsampling, larger aliquants, and multiple
analysis may be the only techniques available to adequately define the content of radionuclides in
heterogeneous samples. Therefore, it is imperative that the analyst choose a preparation approach
appropriate to the nature of the sample.
12.3.2.1

Soils

ASTM C999 provides guidance on the preparation of a homogenous soil sample from
composited core samples. The soil samples are dried at 110 EC until at constant weight, ground
and mixed in a ball mill, and processed through a U.S. Series No. 35 (500-µm or 32-mesh) sieve.
This method is intended to produce a homogeneous sample from which a relatively small
aliquant (10 g) may be drawn for radiochemical analyses.
A similar procedure for homogenizing soil samples is given in HASL-300 (DOE, 1997).
Unwanted material (e.g, vegetation, large rocks) is removed as warranted, and the sample is
dried. If the sample contains small rocks or pebbles, the entire soil sample is crushed to 6.35 mm,
or the entire sample is sieved through a 12.7-mm screen. The sample is blended, then reduced in
size by quartering. This subsample of soil is processed through a grinder, ball mill, sieve, or
pulverizer until the soil is reduced to <1.3 mm (15 mesh equivalent).
Sill et al. (1974) describe a procedure where they dried raw soil samples for two to three hours at
120 EC and then ground the cooled sample lightly in a mortar and pestle. All rocks larger than ¼
inch (6.25 mm) were removed. The sample was charred at 400 EC for two to three hours, cooled
and passed though a No. 35 U.S. standard sieve, and then blended prior to aliquanting (10.0 g are
taken for the analysis).
12.3.2.2

Sediments

ASTM D3976 is a standard practice for the preparation of sediment samples for chemical
analysis. It describes the preparation of test samples collected from streams, rivers, ponds, lakes,
and oceans. The procedures are applicable to the determination of volatile, semivolatile, and
nonvolatile constituents of sediments. Samples are first screened to remove foreign objects and
then mixed by stirring. The solids are allowed to settle and the supernatant liquid is decanted. To
minimize stratification effects due to differential rates of settling, the sample is mixed again
before aliquanting for drying and analysis.
12.3.3 Biota Samples
ASTM D4638 is a standard guide for the preparation of biological samples for inorganic
chemical analysis. It gives procedures for the preparation of test samples of plankton, mollusks,
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fish, and plants. The preparation techniques are applicable for the determination of volatile,
semivolatile, and nonvolatile inorganic compounds in biological materials. However, different
preparation steps are involved for the three classes of inorganic compounds. In the case of
nonvolatile compounds, the first step is to remove foreign objects and most of the occluded
water. For large samples such as fish, samples are homogenized using a tissue disrupter, blender,
or equivalent, and a moisture determination is performed on a one to two gram aliquant. The
samples then are dried by heating in an oven, by dessication, by air drying, by freeze drying, or
by low-temperature drying using an infrared lamp, hot plate, or a low setting on a muffle furnace.
Finally, the samples are dry ashed.
12.3.3.1

Food

The International Atomic Energy Agency offers a guidebook for the measurement of radionuclides in food and the environment, which includes guidance on sample preparation (IAEA, 1989).
Additionally, methods are presented in HASL-300 (DOE, 1997) for the preparation of milk,
vegetables, composite diets, etc. (Table 12.4). These methods involve dry ashing samples
containing non-volatile radionuclides. Initially the samples are completely dried at 125 EC, and
then the temperature is raised slowly over an eight-hour
TABLE 12.4 — Preliminary ashing
period to 500 EC. As the samples are heated, they will
temperature for food samples
reach ignition temperature. It is important to pass
(Method Sr-02-RC, HASL-300 [DOE, 1997])
through this ignition temperature range slowly without
Material
Temp ( EC)
sample ignition. With careful adjustment of the ashing
Eggs . . . . . . . . . . . . . . . 150-250
temperature in a stepwise fashion over this eight-hour
Meat . . . . . . . . . . . . . . . Burning
interval, sample ignition can be avoided. Table 12.4
Fish . . . . . . . . . . . . . . . . Burning
lists the ignition temperature ranges for various foods.
Fruit (fresh) . . . . . . . . . 175-325
Once through the ignition temperature range, the
Fruit (canned) . . . . . . . . 175-325
temperature can be raised more rapidly to 500 EC. The
Milk (dry) . . . . . . . . . . . —
samples can then be ashed at 500 EC for 16 hours.
Milk (wet) . . . . . . . . . . 175-325
Ignition sometimes cannot be avoided if the sample
Buttermilk (dry) . . . . . . —
type contains large amounts of fat. In addition, glowing
of carbonaceous material due to oxidation of carbon
Vegetables (fresh) . . . . 175-225
will be evident during the ashing process. If only a
Vegetables (canned) . . . 175-250
portion of ash is to be used for analysis, it is ground
Root vegetables . . . . . . 200-325
and sieved prior to aliquanting.
Grass . . . . . . . . . . . . . . 225-250
12.3.3.2

Vegetation

There are several DOE site references that contain
examples of sample preparation for vegetation. Los
Alamos National Laboratory (LANL, 1997) recently
grew pinto beans, sweet corn, and zucchini squash in a
field experiment at a site that contained observable
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Flour . . . . . . . . . . . . . . .

Burning

Dry beans . . . . . . . . . . .

175-250

Fruit juices . . . . . . . . . .

175-225

Grains . . . . . . . . . . . . . .

225-325

Macaroni . . . . . . . . . . .

225-325

Bread . . . . . . . . . . . . . .

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levels of surface gross gamma radioactivity within Los Alamos Canyon. Washed edible and
nonedible crop tissues (as well as the soil) were prepared for analysis for various radionuclides.
Brookhaven National Laboratory has also evaluated the effect of its operation on the local
environment. Their site environmental report (DOE, 1995) gives sample preparation steps for
radionuclide analysis of vegetation and fauna (along with ambient air, soil, sewage effluent,
surface water, and groundwater). HASL-300 (DOE, 1997) also describes sample preparation
techniques for vegetation samples for a variety of radionuclides.
12.3.3.3

Bone and Tissue

Bone and tissue samples can be dry ashed in a muffle furnace (DOE, 1997; Fisenne, 1994;
Fisenne et al.,1980), wet ashed with nitric acid and peroxide (Fisenne and Perry, 1978) or
alternately dry ashed and wet ashed with nitric acid until all visible signs of carbonaceous
material has disappeared (McInroy et al., 1985).
12.3.4 Other Samples
The category “other” includes such matrices as concrete, asphalt, coal, plastic, etc. The sample
preparation procedures applied to soils are generally applicable for the “other” category, except
for more aggressive grinding and blending in the initial step. For example, items such as plastic
or rubber that are too flexible to be impact-ground at room temperature must be ground
cryogenically. They are embrittled by chilling and then pulverized. ASTM C114 describes the
sample preparation steps for the chemical analysis of hydraulic cement, whereas ASTM C702
describes the sample preparation of aggregate samples, and is also applicable to lime and
limestone products as noted in ASTM C50. Additionally, ASTM D2013 describes the
preparation of coal samples for analysis.

12.4 Filters
Filters are used to collect analytes of interest from large volumes of liquids or gases. The exact
form of the filter depends on the media (e.g., air, aqueous liquid, nonaqueous liquid), the analyte
matrix (e.g., sediment, suspended particulates, radon gas), and the objectives of the project (e.g.,
volume of sample passing through the filter, flow rate through the filter, detection limits, etc. (see
Section 10.3.2, “Filtration”).
Filter samples from liquids usually consist of the filter with the associated solid material. For
samples with a large amount of sediment, the solid material may be removed from the filter and
analyzed as a solid. When there is a relatively small amount of solid material, the filter may be
considered as part of the sample for analytical purposes. When large volumes of liquid are
processed at high flow rates, filter cartridges often are used. Typically, the cartridge case is not
considered part of the sample, and laboratory sample preparation includes removing the filter
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material and sample from the cartridge case. Any special handling instructions should be
included as SOPs in the planning documents.
Air filters may be particulate filters, which are prepared in the same manner as liquid filters, or
they may be cartridges of absorbent material. Filters that absorb materials are typically designed
for a specific analysis. For example, activated charcoal cartridges are often used to collect
samples of iodine or radon. Silver zeolite cartridges generally are used for sampling iodine
isotopes. These cartridges are often designed to be analyzed intact, so no special sample
preparation is needed. If the cartridges need to be disassembled for analysis, a special SOP for
preparing these samples is usually required.
Homogenization is rarely an issue when preparing filter samples. Typically, the entire filter is
digested and analyzed. However, obtaining a representative sample of a filter does become an
issue when the entire filter is not analyzed. The planning document should give the details of
sample preparation for portions of a filter (e.g., sample size reduction through quartering). Steps
such as using tweezers for holding filters and using individual sample bags should be taken to
prevent the loss of material collected on the filter during handling and processing.

12.5 Wipe Samples
Wipe samples (also referred to as “swipes” or “smears”) are collected to indicate the presence
of removable surface contamination. The removable contamination is transferred from the
surface to the wipe material. The type of filter (paper, membrane, glass fiber, adhesive backing,
etc.) and counting method influence the preparation requirements (Section 10.6, “Wipe Sampling
for Assessing Surface Contamination”).
Wipes are usually counted directly without additional sample preparation. Wipe samples can be
counted directly with a gas flow proportional counter for alpha or beta radioactivity. For gammaemitting radionuclides, the wipe also can be counted directly. For very low-energy emissions,
wipe samples are commonly counted by liquid scintillation (see Chapter 15, Quantification of
Radionuclides).
When destructive analysis is required, the techniques in Chapter 13, Sample Dissolution, and
Chapter 14 Separation Techniques, should be followed. Some wipes have adhesive backing that
can complicate digestion and require more aggressive treatment with acid to dissolve. When
counting with liquid scintillation, the compatibility of the processed wipe with the cocktail is an
important consideration.

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12.6 Liquid Samples
Liquid samples are commonly classified as aqueous, nonaqueous, and mixtures. Aqueous liquids
are most often surface water, groundwater, drinking water, precipitation, effluent, or runoff.
Nonaqueous liquids may include solvents, oils, or other organic liquids. Mixtures may be
combinations of aqueous and nonaqueous liquids, but may include solid material mixed with
aqueous or nonaqueous liquids or both.
Preliminary sample measurements (e.g., conductivity, turbidity) may be performed to provide
information about the sample and to confirm field processing (see measurement of pH to confirm
field preservation in Chapter 11). These measurements are especially useful when there is no
prior historical information available from the sample collection site. In addition, this
information can also be helpful in the performance of certain radiochemical analyses. In many
cases, the results of preliminary measurements can be used to determine the quantity of sample to
be used for a specific analysis.
These preliminary measurements typically require little or no sample preparation. However, they
should be performed on a separate portion of the sample. This avoids any unexpected degradation of the sample parameters during transport and storage, and allows laboratory analysts to
focus on radiochemical analyses. Using a separate aliquant also helps to prevent crosscontamination of samples sent to the laboratory or loss of radionuclides through interaction with
field-measuring equipment.
12.6.1 Conductivity
In radiochemistry, conductivity measurements typically are used as a surrogate to estimate
dissolved solids content for gross-alpha and gross-beta measurements. Because the preservation
of samples with acid prevents the measurement of conductivity, the recommendation is to
perform the QC checks for conductivity in the field when the original measurements are
performed. If the sample is not preserved in the field, the measurement can be done in the
laboratory.
ASTM D1125 is the standard test method for determining the electrical conductivity of water.
The method is used for the measurement of ionic constituents, including dissolved electrolytes in
natural and treated water.
12.6.2 Turbidity
The presence of dissolved or suspended solids, liquids, or gases causes turbidity in water.
Measurement of turbidity provides a means to determine if removal of suspended matter is
necessary in order to meet the specifications for liquid samples as given in the plan document.
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ASTM D1889 is the standard test method for the determination of turbidity of water and
wastewater in the range from 0.05 to 40 nephelometric turbidity units (NTU). In the ASTM
method, a photoelectric nephelometer is used to measure the amount of light that a sample
scatters when the light is transmitted through the sample. Project planning documents should
specify the acceptable turbidity limit for of aqueous samples for direct sample processing without
removing solids.
12.6.3 Filtration
The filtration of samples is based on the appropriate plan document that should also give the
selection of the filter material to be used. If samples have not been filtered in the field, the
laboratory can perform the filtration. Guidance on filtration of liquid samples is provided in
Section 10.3.2. However, preservatives should not be added until sample filtration has been
performed (if stipulated in the project DQOs). This ensures that insoluble materials in the sample
that might be entrained during sample collection do not affect the analytical results.
12.6.4 Aqueous Liquids
Aqueous liquids are a common matrix analyzed by laboratories, and are often referred to as water
samples. Examples of possible aqueous liquids requiring radionuclide analysis include the
following:
•
•
•
•
•
•
•
•

Drinking water;
Surface water;
Ground water;
Soil pore water;
Storage tank water;
Oil production water or brine;
Trench or landfill leachate; and
Water from vegetation.

For certain samples that are not filtered, inversion is a form of homogenization. Typically, the
sample is homogenized by inverting the container several times to mix the sample thoroughly. If
there is some air in the container, the passage of air bubbles through the sample will create
sufficient turbulence to mix the sample thoroughly with three or four inversions of the sample
container. If the sample contains zero headspace (so there is no air in the sample container), the
sample should be inverted and allowed to stay inverted for several seconds before the next
inversion. Ten to twenty inversions of the sample container may be required to ensure that the
sample is mixed thoroughly under zero headspace conditions. Simply shaking the container will
not mix the contents as thoroughly as inverting the sample container. Mechanical shakers,
mixers, or rotators may be used to homogenize aqueous samples thoroughly.
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Filtration and acidification performed in the field is typically the only preparation required for
aqueous liquids (Chapter 10). A general discussion concerning preparation of water samples for
the measurement of radioactivity is presented in NCRP (1976). PNL/ACL (1992) gives a number
of sample preparation methods for various materials, including water samples.
ASTM gives standard test methods for the preparation of water samples for the determination of
alpha and beta radioactivity (ASTM D1943 and D1890, respectively). After collecting the water
sample in accordance with ASTM D3370, the sample is made radioactively homogeneous by
adding a reagent in which the radionuclides present in the sample are soluble in large
concentrations. Acids, complexing agents, or chemically similar stable carriers may be used to
obtain homogeneity. The chemical nature of the radionuclides and compounds present and the
subsequent steps in the method will indicate the action to be taken. Different radiochemical
preparation techniques for freshwater and seawater samples are illustrated in EPA (1979) and for
drinking water in EPA (1980).
12.6.5 Nonaqueous Liquids
Nonaqueous liquids can be substances other than water such as organic solvents, oil, or grease.
Many organic solvents are widely used to clean oil, grease, and residual material from electrical
and mechanical equipment. The resulting waste liquid may contain a significant amount of solid
material. It may be necessary to filter such liquids to determine (1) if the analyte is contained in
the filtrate and is soluble, or (2) if the analyte is contained in the solids and therefore is insoluble.
The appropriate plan document should be reviewed to determine if filtration is necessary. ASTM
C1234 describes the preparation of homogeneous samples from nuclear processing facilities.
Homogenization of nonaqueous samples is accomplished in a manner similar to that for aqueous
samples. Visual inspection is typically used as a qualitative measure of homogeneity in nonaqueous samples. If a quantitative measure of mixing is desired, turbidity measurements can be
performed after a predetermined amount of mixing (e.g., every 10 inversions, every 2 minutes,
etc.) until a steady level of turbidity is achieved (e.g., 1 to 10 percent variance, depending on the
project objectives—see ASTM D1889, Standard Test Method for Turbidity of Water).
DOE (ANL/ACL, 1995) evaluated sample preparation techniques used for the analysis of oils. In
evaluating the performance of a sample preparation technique, DOE considered the following
qualities to be important:
•
•
•
•
•
•

Thorough sample decomposition;
Retention of volatile analytes;
Acceptable analyte recovery;
Minimal contamination from the environment or the digestion vessel;
Low reagent blanks; and
Speed.

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One of the preparation methods involved combustion of oil under oxygen at 25 atm pressure
(ASTM E926) and another used nitric acid decomposition of the oil in a sealed vessel heated
with a microwave (EPA, 1990).
Many nonaqueous liquids present a health hazard (e.g., carcinogenicity) or require special safety
considerations (e.g., flammability). Any special handling requirements based on health and safety
considerations should be documented in the planning documents.
12.6.6 Mixtures
Some common examples of mixtures that may be encountered by the laboratory are water with
lots of total dissolved solids and undissolved solids or water and oil in separate layers. The
following sections discuss preparation procedures for these types of mixtures.
12.6.6.1

Liquid-Liquid Mixtures

When aqueous and nonaqueous liquids are combined, they usually form an immiscible mixture,
such as oil and water.1 In most cases, a separatory funnel helps in separating the liquids into two
samples. Each sample then is analyzed separately. If, in the rare case, both liquids must be
processed together, there is greater difficulty in preparing the combined liquids for analysis.
Obtaining a homogenous aliquant is a key consideration in this case. Often times, the entire
sample should be analyzed. This approach avoids processing problems and yields the desired
result.
12.6.6.2

Liquid-Solid Mixtures

Mixtures of liquids and solids are usually separated by filtering, centrifuging, or decanting, and
the two phases are analyzed separately. If the mixture is an aqueous liquid and a solid, and will
be analyzed as a single sample, the sample is often treated as a solid. Completely drying the
sample followed by dry ashing before any attempt at wet ashing is recommended to reduce the
chance of organic solids reacting with strong oxidizing acids (e.g., H 2SO4, HNO3, etc.). If the
mixture includes a nonaqueous liquid and a solid, it is suggested that the phases be separated by

1

It is often necessary to determine which liquid is aqueous and which liquid is nonaqueous. Never assume that the
top layer is always nonaqueous, or the bottom layer is always aqueous. The density of the bottom layer is always
greater than the density of the top layer. Halogenated solvents (e.g., carbon tetrachloride, CCl4) tend to have
densities greater than about 1 g/mL, so they typically represent the bottom layer. Other organic liquids (e.g., diethyl
ether, oil, etc.) tend to have densities less than 1g/mL, so they typically represent the top layer. Mixtures of organic
liquids may have almost any density. To test the liquids, add a drop of water to the top layer. If the drop dissolves in
the top layer, the top layer is aqueous. If the drop settles through the top layer and dissolves in the bottom layer, the
bottom layer is aqueous.
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filtration and the solid rinsed thoroughly with a volatile solvent such as ethanol or methanol
before continuing with the sample preparation process.
In rare cases where a sample contains a mixture of aqueous liquid, nonaqueous liquid, and solid
material, the sample can be separated into three different phases before analysis. The sample
should be allowed to settle overnight and the liquids decanted. The liquids can then be separated
in a separatory funnel without the solid material clogging the funnel. Each liquid should be
filtered to remove any remaining solid material. The solid should be filtered to remove any
remaining liquid and rinsed with a volatile solvent. This rinse removes any traces of organic
liquids to reduce problems during subsequent dissolution activities. The three phases are then
analyzed separately. If necessary, the results can be added together to obtain a single result for the
mixture after the separate analyses are completed.

12.7 Gases
Sample preparation steps are usually not required for gas samples. Lodge (1988) gives general
techniques, including any necessary sample preparation, for the sampling and storage of gases
and vapors. The determination of the tritium content of water vapor in the atmosphere is one of
the example procedures. ASTM D3442 is a standard test method for the measurement of total
tritium activity in the atmosphere. Sample preparation is covered in this test method.
EPA (1989) may be used to demonstrate compliance with the radionuclide National Emission
Standards for Hazardous Air Pollutants (NESHAP). This document includes references to air
sampling and sample preparation. Table 3-1 of EPA (1989) lists numerous references to
radionuclide air sampling and preparation, including Cehn (1979), Eichling (1983), Allied
Chemical (1982), and Browning et al. (1978).

12.8 Bioassay
Analyses of bioassay samples are necessary to monitor the health of employees involved in
radiological assessment work. Normally these types of samples include urine and fecal
specimens.
Urine samples are typically wet ashed with nitric acid (DOE, 1997) or with nitric acid and
peroxide (RESL, 1982). Alternatively, there are procedures that co-precipitate the target analytes
in urine by phosphate precipitation (Horwitz et al., 1990; Stradling and Popplewell, 1974; Elias,
1997). Fecal samples are normally dry ashed in a muffle furnace (DOE, 1997), or prepared by
lyophilization, “freeze drying” (Dugan and McKibbin, 1993).
It is important to note that although ANSI N13.30 indicates that aliquanting a homogeneous
sample to determine the activity present in the total sample is acceptable, this standard dictates
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that the entire sample should be prepared for analysis and the aliquant taken after the sample
preparation has been completed.

12.9 References
12.9.1 Cited Sources
Allied Chemical UF6 Conversion Plant. 1982. “Application for Renewal of Source Material
License: SUB-526, Docket 40-3392,” Metropolis, Illinois.
American National Standards Institute (ANSI) N13.30 Performance Criteria for Radiobioassay.
Health Physics Society. 1996.
Argonne National Laboratory/Analytical Chemistry Laboratory (ANL/ACL). 1992. Innovative
Methods for Inorganic Sample Preparation. April 1992.
Argonne National Laboratory/Analytical Chemistry Laboratory (ANL/ACL). 1995. Preparation
of Waste Oil for Analysis to Determine Hazardous Metals. July.
Association of Official Analytical Chemists International (AOAC) Official Method 985.14.
“Moisture in Meat and Poultry Products,” in Official Methods of Analysis of AOAC
International. P. Cuniff, Ed., Arlington, VA. 1995.
Association of Official Analytical Chemists International (AOAC) Official Method 985.26.
“Solids (Total) in Processed Tomato Products”, In Official Methods of Analysis of AOAC
International. P. Cuniff, Ed.; Association of Official Analytical Chemists International:
Arlington, VA. 1995.
American Society for Testing and Materials (ASTM) C50. Standard Practice for Sampling,
Inspection, Packing, and Marking of Lime and Limestone Products. West Conshohocken,
PA.
American Society for Testing and Materials (ASTM) C114. Standard Test Method for Chemical
Analysis of Hydraulic Cement. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) C702. Standard Practice for Reducing
Samples of Aggregate to Testing Size, Vol 04.02. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) C999. Standard Practice for Soil Sample
Preparation for the Determination of Radionuclides. West Conshohocken, PA.

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American Society for Testing and Materials (ASTM) C1234. Standard Test Method for
Preparation of Oils and Oily Waste Samples by High-Pressure, High-Temperature Digestion
for Trace Element Determinations. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) D1125. Standard Test Method for
Determining the Electrical Conductivity of Water. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) D1889. Standard Test Method for Turbidity
of Water. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) D1890. Standard Test Method for Beta
Particle Radioactivity of Water. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) D1943. Standard Test Method for Alpha
Particle Radioactivity of Water. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) D2013. Standard Method of Preparing
Coal Samples for Analysis. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) D3370. Standard Practices for Sampling
Water from Closed Conduits. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) D3442. Standard Test Method for Gaseous
Tritium Content of the Atmosphere. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) D3974. Standard Practice for Extraction of
Trace Elements from Sediments. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) D3975. Standard Practice for Development
and Use (Preparation) of Samples for Collaborative Testing of Methods for Analysis of
Sediments. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) D3976. Standard Practice for Preparation
of Sediment Samples for Chemical Analysis. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) D4638. Standard Guide for Preparation of
Biological Samples for Inorganic Chemical Analysis. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) D4643. Standard Test Method for
Determination of Water (Moisture) Content in Soil by the Microwave Oven Method. West
Conshohocken, PA.
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American Society for Testing and Materials (ASTM) E926. Standard Practices for Preparing
Refuse-Derived Fuel (RDF) Samples for Analyses of Metals. West Conshohocken, PA.
American Society for Testing and Materials (ASTM) E1358. Standard Test Method for
Determination of Moisture Content of Particulate Wood Fuels Using a Microwave Oven.
West Conshohocken, PA.
Beary, E.S. 1988. “Comparison of Microwave Drying and Conventional Drying Techniques For
Reference Materials.” Anal. Chem Vol. 60, pp. 742-746.
Bernabee, R. P., D. R. Percival, and D. B. Martin. 1980. “Fractionation of Radionuclides in
Liquid Samples from Nuclear Power Facilities.” Health Physics Vol. 39, pp. 57-67.
Bock, R. 1979. A Handbook of Decomposition Methods in Analytical Chemistry. International
Textbook Company, Limited. T. & A. Constable Ltd., Great Britain.
Browning, E.J., K. Banerjee, and W.E. Reisinger. 1978. “Airborne Concentrations of I-131 in a
Nuclear Medicine Laboratory,” Journal of Nuclear Medicine, Vol. 19, pp. 1078-1081.
Cameron, P., and Murgatroyd, K. 1996. Good Pharmaceutical Freeze-Drying Practice.
Interpharm Press.
Cehn, J.I. 1979. A Study of Airborne Radioactive Effluents from the Pharmaceutical Industry,
Fianl Report, Prepared by Teknekron, Inc., for the U.S. EPA Eastern Environmental Research
Facility, Montgomery, AL.
Dean, J.A. 1995. Analytical Chemistry Handbook, McGraw-Hill, Inc., New York.
DeVoe, J.R. 1961. Radioactive Contamination of Materials Used in Scientific Research.
Publication 895, NAS-NRC.
U.S. Department of Energy (DOE). 1995. Brookhaven National Laboratory Site Environmental
Report for Calendar Year 1995, Naidu, J. R., D. E. Paquette, G. L. Schroeder, BNL
December, 1996.
U.S. Department of Energy. 1997 (DOE).EML Procedures Manual (HASL-300-Ed.28), Edited
by N.A. Chieco, Environmental Measurements Laboratory.
Dugan, J.P. and T.T. McKibbin. 1993. “Preparation of Fecal Samples for Radiobioassay by
Lyophilization,” Radioactivity & Radiochemistry 4:3, pp. 12-15.

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Eberhardt, L.L., and R.O. Gilbert. 1980. “Statistics and Sampling in Transuranic Studies,” in
Transuranic Elements in the Environment , edited by W.C. Hanson, U.S. Department of
Energy. DOE/TIC-22800.
Eichling, J. 1983. “The Fraction of Material Released as Airborne Activity During Typical
Radioiodinations,” Proceedings of the 9th Biennial Conference of Campus Radiation Safety
Officers, University of Missouri-Columbia, June 6-8, 1983.
Elias, G. 1997. “A Rapid Method for the Analysis of Plutonium and Uranium in Urine Samples,”
Radioactivity & Radiochemistry 8:3, pp. 20-24.
U.S. Environmental Protection Agency (EPA). 1979. Radiochemical Analytical Procedures for
Analysis of Environmental Samples; F. B. Johns, P. B. Hahn, D. J. Thome, and E. W.
Bretthauer, EMSL, March 1979.
U.S. Environmental Protection Agency (EPA). 1980. Prescribed Procedures for Measurement of
Radioactivity in Drinking Water. H. L. Krieger and E. L. Whittaker, EPA 600-4-80-032,
August 1980.
U.S. Environmental Protection Agency (EPA). 1989. Background Information Document:
Procedures Approved for Demonstrating Compliance with 40 CFR Part 61, Subpart I. EPA
520-1-89-001, Office of Radiation Programs, October, 1989.
U.S. Environmental Protection Agency (EPA). 1990. Test Methods for Evaluating Solid Waste—
Physical/Chemical Methods. SW-846, Third Edition, Method 3051.
U.S. Environmental Protection Agency (EPA). 1992. Manual for the Certification of
Laboratories Analyzing Drinking Water: Criteria and Procedures. Fourth Edition, EPA 814B-92-002, Office of Ground Water and Drinking Water, Cincinnati, Ohio.
Fisenne, I.M. and P. Perry 1978. “The Determination of Plutonium in Tissue by Aliquat-336
Extraction,” Radiochem Radioanal. Letters Vol. 33, pp. 259-264.
Fisenne, I.M., P. Perry and G.A. Welford. 1980. “Determination of Uranium Isotopes in Human
Bone Ash,” Anal. Chem. Vol. 52, pp. 777-779.
Fisenne, I.M. 1994. “Lead-210 in Animal and Human Bone: A New Analytical Method,” Env.
Int. Vol. 20, pp. 627-632.
Greenberg, A.E., L.S. Clesceri, and A.D. Eaton (Eds). 1992. Standard Methods for the
Examination of Water and Wastewater. American Public Health Association.
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Laboratory Sample Preparation
Greenwood, N. N. and A. Earnshaw. 1984. Chemistry of the Elements. Pergamon Press, Inc.
Elmsford, New York.
Gy, Pierre M. 1992. Sampling of Heterogeneous and Dynamic Material Systems: Theories of
Heterogeneity, Sampling, and Homogenizing. Elsevier, Amsterdam, The Netherlands.
Horwitz, E.P., M.L. Dietz, D.M. Nelson, J.J. LaRosa, and W.D. Fairman 1990. “Concentration
and Separation of Actinides from Urine using a Supported Bifunctional Organophosphorous
Extractant,” Analytica Chimica Acta. Vol. 238, pp. 263-271.
IAEA. 1989. Measurement of Radionuclides in Food and the Environment—A Guidebook.
Technical Reports Series No. 295, International Atomic Energy Agency, Vienna.
Kingston, H.M., and Jassie, L.B. 1988. Introduction to Microwave Sample Preparation: Theory
and Practice, American Chemical Society, Washington, DC.
Koh, T.S. 1980. Anal Chem Vol. 52, pp. 1978-1979.
Kralian, M.A., M.J. Atkins, and S.A. Farber. 1990. “Guidelines for Effective Low-Level
Contamination Control in a Combination Environmental/Radioactive Waste Analysis
Facility,” Radioactivity & Radiochemistry. 1:3, pp. 8-18.
Lodge, J. 1988. Methods of Air Sampling and Analysis. Third Edition, CRC Press, Florida.
LANL. 1996. Application of Cryogenic Grinding to Achieve Homogenization of Transuranic
Waste, Atkins, W. H., LANL-13175.
Los Alamos National Laboratory (LANL). 1997. Radionuclide Concentration in pinto beans,
sweet corn, and zucchini squash grown in Los Alamos Canyon at Los Alamos National
Laboratory, Fresquez, P. R., M. A. Mullen, L. Naranjo, and D. R. Armstrong, May 1997.
Lucas, H.F., Jr. 1963. “A Fast and Accurate Survey Technique for Both Radon-222 and Radium226,” The Natural Radiation Environment, Proceedings of the International Symposium,
William Rice University, Houston, TX, 315-319.
Lucas, H.F. 1967. “A Radon Removal System for the NASA Lunar Sample Laboratory: Design
and Discussion,” Argonne National Laboratory Radiological Physics Division Annual
Report, ANL-7360.
McInroy, J.F., H.A. Boyd, B.C. Eutsler, and D. Romero. 1985. “Part IV: Preparation and
Analysis of the Tissue and Bones,” Health Physics, 49:4, pp. 585-621.
JULY 2004

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Laboratory Sample Preparation
NCRP Report No. 50. 1976. Environmental Radiation Measurements.
Obenhauf, R.H., R. Bostwick, W. Fithian, M. McCann, J.D. McCormack, and D. Selem. 2001.
SPEX CertiPrep Handbook of Sample Preparation and Handling, SPEX CertiPrep, Inc., 203
Norcross Avenue Metuchen, NJ 08840. Also at http://www.spexcsp.com/spmain/sprep/
handbook/tocprime.htm.
Pacific Northwest Laboratories/Analytical Chemistry Laboratory (PNL/ACL). 1992. Procedure
Compendium. Volume 2: Sample Preparation Methods. PNL-MA-559.
Pitard, F. F. 1993. Pierre Gy’s Sampling Theory and Practice. CRC Press, Inc., Boca Raton, FL.
Second Edition.
RESL Analytical Chemistry Branch Procedures Manual. 1982. U.S. Department of Energy, Idaho
Falls, Idaho, IDO-12096.
Schilt, A. 1979. Perchloric Acid and Perchlorates. The G. Frederick Smith Chemical Company,
Columbus, Ohio.
Schwochau, K. 2000. Technetium: Chemistry and Radiopharmaceutical Applications, WileyVCH (Federal Republic of Germany).
Scwedt, G. 1997. The Essential Guide to Analytical Chemistry (Translation of the revised and
updated German Second Edition. Translated by Brooks Haderlie), John Wiley & Sons,
England.
Sedlet, J. 1966. “Radon and Radium,” in Treatise on Analytical Chemistry, Part II, Vol. IV,
p219-366, edited by I.M. Kolthoff and P.J. Elving, John Wiley & Sons, Inc, New York.
Shugar, G.J. and J.T. Ballinger. 1996. Chemical Technicians’ Ready Reference Handbook.
McGraw-Hill, New York.
Sill, C.W., K.W. Puphal, and F.D. Hindman. 1974. “Simultaneous Determination of AlphaEmitting Nuclides of Radium through Californium in Soil,” Anal. Chem 46:12, pp. 17251737.
Sill, C.W. 1975. “Some Problems in Measuring Plutonium in the Environment,” Health Physics.
Vol. 29, pp. 619-626.
Sneed, M.C. and Brasted, R.C. 1958. Comprehensive Inorganic Chemistry, New York: D. Van
Nostrand.
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Laboratory Sample Preparation
Stradling, G.N. and D.S. Popplewell. 1974. “Rapid Determination of Plutonium in Urine by
Ultrafiltration,” Int. J. Appl. Radiation Isotopes. Vol. 25, p 217.
Walter, P., S. Chalk, and H. Kingston. 1997. “Overview of Microwave-Assisted Sample
Preparation.” Chapter 2, Microwave-Enhanced Chemistry, H. Kingston and S. Haswell,
editors, American Chemical Society, Washington, DC.
Wang, C.H., Willis, D.L., and Loveland W.D. 1975. Radiotracer Methodology in the Biological,
Environmental, and Physical Sciences. Prentice-Hall, Inc., New Jersey.
Windholz, M. 1976. The Merck Index (9th edition), Merck and Co. Inc., New Jersey.
Yamamoto, M., Komura, K., and Ueno, K. 1989. “Determination of Low-Level 226Ra in
Environmental Water Samples by Alpha-Ray Spectrometry,” Radiochimica Acta Vol. 46, pp.
137-142.
Zhang, H. and P. Dotson. 1998. The Use of Microwave Muffle Furnace for Dry Ashing Plant
Tissue Samples. CEM Corporation. Also, Commun. Soil Sci. Plant Anal. 25:9&10, pp. 13211327 (1994).
12.9.2 Other Sources
American Society for Testing and Materials (ASTM) D5245. Standard Practice for Cleaning
Laboratory Glassware, Plasticware, and Equipment Used in Microbiological Analyses. West
Conshohocken, PA.
American Society for Testing and Materials (ASTM) E1157. Standard Specification for
Sampling and Testing of Reusable Laboratory Glassware. West Conshohocken, PA.
U.S. Environmental Protection Agency. 1987. Eastern Environmental Radiation Facility
Radiochemistry Procedures Manual. Compiled and edited by R. Lieberman, EPA 520-5-84006, Office of Radiation Programs. August.
Kahn, B. 1973. “Determination of Radioactive Nuclides in Water,” in Water and Water Pollution
Handbook, Vol. 4, p. 1357 (L.L. Ciaccio, Ed.). M. Decker, New York.
Kahn, B., Shleien, B., and Weaver, C. 1972. “Environmental Experience with Radioactive
Effluents From Operating Nuclear Power Plants,” page 559 in Peaceful Uses of Atomic
Energy, Vol. 11 (United Nations, New York).

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Laboratory Sample Preparation
Krahenbuhl, M.P., Slaughter, D.M. 1998. “Improving Process Methodology for Measuring
Plutonium Burden in Human Urine Using Fission Track Analysis,” J. Radioanalytical and
Nuclear Chemistry, 220:1-2, pp. 153-160.
Krieger, H.L. and E.L. Whittaker. 1980. “Prescribed Procedures for Measurement of
Radioactivity in Drinking Water,” Environmental Monitoring and Support Laboratory,
Cincinnati, OH, EPA-600/4-80-032.
Laug, E.P. 1934. Ind. Eng. Chem., Anal Ed. Vol. 13, pp. 419.
McFarland, R.C. 1998a. “Determination of Alpha-Particle Counting Efficiency for Wipe-Test
Samples,” Radioactivity & Radiochemistry. 9:1, pp. 4-8.
McFarland, R.C. 1998b. “Determination of Counting Efficiency for Wipe-Test Samples
Containing Radionuclides that Emit High-Energy Beta Particles,” Radioactivity &
Radiochemistry 9:1, pp. 4-9.
Nichols, S.T. 2001. “New Fecal Method for Plutonium and Americium,” J. Radioanalytical and
Nuclear Chemistry, 250:1, pp. 117-121.
Shugar and Dean. 1990. The Chemist’s Ready Reference Handbook, McGraw-Hill.

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13 SAMPLE DISSOLUTION
13.1 Introduction
The overall success of any analytical procedure depends upon many factors, including proper
sample preparation, appropriate sample dissolution, and adequate separation and isolation of the
target analytes. This chapter describes sample dissolution techniques and strategies. Some of the
principles of dissolution are common to those of radiochemical separation that are described in
Chapter 14 (Separation Techniques), but their importance to dissolution is reviewed here.
Sample dissolution can be one of the biggest challenges facing the analytical chemist, because
most samples consist mainly of unknown compounds with unknown chemistries. There are many
factors for the analyst to consider: What are the measurement quality objectives of the program?
What is the nature of the sample; is it refractory or is there only surface contamination? How
effective is the dissolution technique? Will any analyte be lost? Will the vessel be attacked? Will
any of the reagents interfere in the subsequent analysis or can any excess reagent be removed?
What are the safety issues involved? What are the labor and material costs? How much and what
type of wastes are generated? The challenge for the analyst is to balance these factors and to
choose the method that is most applicable to the material to be analyzed.
The objective of sample dissolution is to mix a solid or nonaqueous liquid sample quantitatively
with water or mineral acids to produce a homogeneous aqueous solution, so that subsequent
separation and analyses may be performed. Because very few natural or organic materials are
water-soluble, these materials routinely require the use of acids or fusion salts to bring them into
solution. These reagents typically achieve dissolution through an oxidation-reduction process that
leaves the constituent elements in a more soluble form. Moreover, because radiochemists
routinely add carriers or use the technique of isotope dilution to determine certain radioisotopes,
dissolution helps to ensure exchange between the carrier or isotopic tracer and the element or
radioisotope to be determined, although additional chemical treatment might be required to
ensure exchange.
Contents
There are three main techniques for sample
decomposition discussed in this chapter: fusion; 13.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . 13-1
wet ashing, acid leaching, or acid dissolution;
13.2 The Chemistry of Dissolution . . . . . . . . . . . . 13-2
and microwave digestion.
13.3 Fusion Techniques . . . . . . . . . . . . . . . . . . . . 13-6

The choice of technique is determined by the
type of sample and knowledge of its physical
and chemical characteristics. Fusion and wet
ashing techniques may be used singly or in
combination to decompose most samples
analyzed in radioanalytical laboratories.
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13.4 Wet Ashing and Acid Dissolution
Techniques . . . . . . . . . . . . . . . . . . . . . . . . .
13.5 Microwave Digestion . . . . . . . . . . . . . . . . .
13.6 Verification of Total Dissolution . . . . . . . .
13.7 Special Matrix Considerations . . . . . . . . . .
13.8 Comparison of Total Dissolution and Acid
Leaching . . . . . . . . . . . . . . . . . . . . . . . . . . .
13.9 References . . . . . . . . . . . . . . . . . . . . . . . . . .

13-1

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13-21
13-23
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Sample Dissolution
Leaching techniques are used to determine the soluble fraction of the radionuclide of interest
under those specific leaching conditions. Different formulas for leaching agents will yield
different amounts of leachable analyte. It should be recognized that the information so obtained
leaves unknown the total amount of analyte present in the sample. Because recent advances in
microwave vessel design (e.g., better pressure control and programmable temperature control)
have allowed for the use of larger samples, microwave dissolution is becoming an important tool
in the radiochemistry laboratory. Leaching and the newer closed-vessel microwave methods
provide assurance that only minimal analyte loss will occur through volatilization.
Because of the potential for injury and explosions during sample treatment, it is essential that
proper laboratory safety procedures be in place, the appropriate safety equipment be available, a
safe work space be provided, and that the laboratory personnel undergo the necessary training to
ensure a safe working environment before any of these methods are used. Review the Material
Data Safety Sheets for all chemicals before their use.
Aspects of proper sample preparation, such as moisture removal, oxidation of organic matter, and
homogenization, were discussed in Chapter 12, Laboratory Sample Preparation. Fundamental
separation principles and techniques, such as complexation, solvent extraction, ion exchange, and
co-precipitation, are reviewed in Chapter 14, Separation Techniques.
There are many excellent references on sample dissolution (e.g., Bock, 1979; Bogen, 1978; Dean,
1995; Sulcek and Povondra, 1989).

13.2 The Chemistry of Dissolution
In order to dissolve a sample completely, each insoluble component must be converted into a
soluble form. Several different chemical methods may need to be employed to dissolve a sample
completely; usually, the tracer is added to the sample at the time of sample dissolution. Initially
the sample may be treated with acids yielding an insoluble residue. The residue may need to be
dissolved using fusion or hydrofluoric acid (HF) and then combined with the original mixture or
analyzed separately. In either case, the tracer/carrier should be added to the sample during the
first step of chemical change (e.g., acid dissolution as above) so that the yield for the entire
process may be determined accurately. An outline of the principles of these chemical methods is
provided in this section, but a complete description is available in Chapter 14, where the
principles are applied to a broader range of topics.
13.2.1 Solubility and the Solubility Product Constant, Ksp
The solubility data of many compounds, minerals, ores, and elements are available in reference
manuals. Solubilities typically are expressed in grams of substance per 100 mL of solvent,
although other units are sometimes used. The information is more complete for some substances
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Sample Dissolution
than others, and for many substances solubility is expressed only in general terms, such as
“soluble,” “slightly soluble,” or “insoluble.” Many environmental samples consist of complex
mixtures of elements, compounds, minerals, or ores, most of which are insoluble and must be
treated chemically to dissolve completely. In some cases, the sample constituents are known to
the analyst, but often they are not. Solubility data might not be available even for known
constituents, or the available data might be inadequate. Under these circumstances, sample
dissolution is not a simple case of following the solubilities of known substances. For known
constituents with solubility data, the solubilities indicate those that must be treated to complete
dissolution. This, in turn, provides a guide to the method of treatment of the sample. Given the
potential complexity of environmental samples, it is difficult to describe conditions for
dissolving all samples. Sometimes one method is used to dissolve one part of the sample while
another is used to dissolve the residue.
The solubility of many compounds in water is very low, on the order of small fractions of a
grams per 100 mL. The solubility may be expressed by a solubility product constant (K sp), an
equilibrium constant for dissolution of the compound in water (see Section 14.8.3.1, “Solubility
and Solubility Product Constant”). For example, the solubility product constant for strontium
carbonate, a highly insoluble salt (0.0006 g/100 mL), is the equilibrium constant for the process:
and is represented by:

SrCO3(s) 6 Sr+2(aq) + CO3!2(aq)
Ksp = [Sr+2][CO3!2] = 1.6×10!9

The brackets indicate the molar concentration (moles/liter) of the respective ions dissolved in
water. The very small value of the constant results from the low concentration of dissolved ions,
and the compound is referred to as “insoluble.” Chemical treatment is necessary sometimes to
dissolve the components of a compound in water. In this example, strontium carbonate requires
the addition of an acid to solubilize Sr+2. The next section describes chemical treatment to
dissolve compounds.
13.2.2 Chemical Exchange, Decomposition, and Simple Rearrangement Reactions
Chemical exchange, decomposition, and simple rearrangement reactions refer to one method for
solubilizing components of a sample. In this chemical process, the sample is treated to convert
insoluble components to a soluble chemical species using chemical exchange (double displacement), decomposition, or simple rearrangement reactions rather than oxidation-reduction
processes or complex formations. Some reagents solubilize sample components using chemical
exchange. Radium or strontium cations in radium or strontium carbonate (RaCO3 or SrCO3)
exchange the carbonate anion for the chloride ion on acid treatment with HCl to produce the
soluble chlorides; the carbonic acid product decomposes to carbon dioxide and water:
RaCO3 + 2 HCl 6 RaCl2 + H2CO3
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Sample Dissolution
and the net reaction is as follows:

H2CO3 6 CO2 + H2O

RaCO3 + 2 HCl 6 RaCl2 + CO2 + H2O
Sodium pyrosulfate fusion, for example, converts zirconia (ZrO2) into zirconium sulfate
[Zr(SO4)2], which is soluble in acid solution by a simple (nonoxidative) rearrangement of oxygen
atoms (Hahn, 1961; Steinberg, 1960):
ZrO2 + 2 Na2S2O7 6 2 Na2SO4 + Zr(SO4)2
Many environmental samples contain insoluble silicates, such as aluminum silicate [Al2(SiO3)3 or
Al2O3 · 3SiO2], which can be converted into soluble silicates by fusion with sodium carbonate:
Al2(SiO3)3 + 4 Na2CO3 6 3 Na2SiO3 + 2 NaAlO2 + 4 CO2
Dissolution of radium from some ores depends on the exchange of anions associated with the
radium cation (sulfate for example) to generate a soluble compound. Extraction with nitric acid is
partly based on this process, generating soluble radium nitrate.
13.2.3 Oxidation-Reduction Processes
Oxidation-reduction (redox) processes are an extremely important aspect of sample dissolution.
The analyte may be present in a sample in several different chemical forms or oxidation states.
As an example, consider a ground-water sample that contains 129I as the analyte. The iodine may
be present in any of the following inorganic forms: I!, I2, IO!, or IO3!. If the ground water has a
high reduction potential or certain bacteria are present, the iodine also may be present as CH3I. It
is of paramount importance to ensure that all of these different forms of iodine are brought to the
same oxidation state (e.g., to iodate) at the time of first change in redox environment or change in
sample composition. Furthermore, accurate assessment of chemical yield only can be determined
if the tracer or carrier is added prior to a change in chemical form or oxidation state of the analyte
at an initial point in the digestion process. This process is referred to as “equilibration of the
tracer/carrier and analyte.” From this point on during the sample analysis, any loss that occurs to
the analyte will occur to an equal extent for the tracer/carrier, thus allowing the calculation of a
chemical yield for the process.
A redox reaction redistributes electrons among the atoms, molecules, or ions in the reaction. In
some redox reactions, electrons actually are transferred from one reacting species to another. In
other redox reactions, electrons are not transferred completely from one reacting species to
another; the electron density about one atom decreases, while it increases about another atom. A
complete discussion of oxidation and reduction is found in Section 14.2, “Oxidation-Reduction
Processes.”
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Sample Dissolution
Many oxidizing agents used in sample dissolution convert metals to a stable oxidation state
displacing hydrogen from hydrochloric, nitric, sulfuric, and perchloric acids. (This redox process
often is referred to as nonoxidative hydrogen replacement by an active metal, but it is a redox
process where the metal is oxidized to a cation, usually in its highest oxidation state, and the
hydrogen ion is reduced to its elemental form.) Dissolution of uranium for analysis is an example
of hydrogen-ion displacement to produce a soluble substance (Grindler, 1962):
U + 8 HNO3 6 UO2(NO3)2 + 6 NO2 + 4 H2O
Prediction of the reactivity of a metal with acids is dependent on its position in the electromotive
force series (activity series). A discussion of the series appears in Section 13.4.1, “Acids and
Oxidants.” In general, metals with a negative standard reduction potential will replace hydrogen
and be dissolved. Perchloric acid offers a particular advantage because very soluble metal
perchlorate salts are formed.
Other important oxidizing processes depend on either oxidizing a lower, less soluble oxidation
state of a metal to a higher, more soluble state or oxidizing the counter anion to generate a more
soluble compound. Oxidation to a higher state is common when dissolving uranium samples in
acids or during treatment with fusion fluxes. The uranyl ion (UO2+2) forms soluble salts—such as
chloride, nitrate, and perchlorate—with anions of the common acids (Grindler, 1962). (Complexion formation also plays a role in these dissolutions; see the next section). Dissolution of oxides,
sulfides, or halides of technetium by alkaline hydrogen peroxide converts all oxidation states to
the soluble pertechnetate salts (Cobble, 1964):
2 TcO2 + 2 NaOH + 3 H2O2 6 2 NaTcO4 + 4 H2O
13.2.4 Complexation
The formation of complex ions (see also Section 14.3, “Complexation”) is important in some
dissolution processes, usually occurs in conjunction with treatment by an acid, and also can occur
during fusion. Complexation increases solubility in the dissolution mixture and helps to minimize hydrolysis of the cations. The solubility of radium sulfate in concentrated sulfuric acid is
the result of forming a complex-ion, Ra(SO4)2!2. The ability of both hydrochloric and hydrofluoric acids to act as a solubilizing agent is dependent on their abilities to form stable complex
ions with cations. Refractory plutonium samples are solubilized in a nitric acid-hydrofluoric acid
solution forming cationic fluorocomplexes such as PuF+3 (Booman and Rein, 1962). Numerous
stable complexes of anions from solubilizing acids (HCl, HF, HNO 3, H2SO4, HClO4) contribute
to the dissolution of other elements, such as americium, cobalt, technetium, thorium, uranium,
and zirconium (see Section 14.10, “Analysis of Specific Radionuclides”). The process of fusion
with sodium carbonate to solubilize uranium samples is also based on the formation of
UO2(CO3)2!4 after the metal is oxidized to U+6 (Grindler, 1962).
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Sample Dissolution
13.2.5 Equilibrium: Carriers and Tracers
Carriers and tracers that are sometimes required for radiochemical separation procedures usually
are added to samples before dissolution in order to subject them to the same chemical treatment
as the analyte. Addition as soon as practical promotes equilibrium with the analyte. The dissolution process tends to bring the carriers and tracers to the same oxidation state as the analyte and
ensures complete mixing of all the components in solution. Acid mixtures also create a large
hydrogen-ion concentration that minimizes the tendency of cations to hydrolyze and subsequently
form insoluble complexes. Detailed discussions of carriers and tracers as well as radioactive
equilibrium are found in Section 14.9, “Carriers and Tracers,” Section 14.10, “Analysis of
Specific Radionuclides,” and Attachment 14A, “Radioactive Decay and Equilibrium.” The
immediate and final forms of these tracers, carriers, and analytes are crucial information during
the analytical process. During each of the steps in a given separation method, the analyst should
be aware of the expected oxidation states of the analyte and its tendency to hydrolyze, polymerize, and form complexes and radiocolloids, and other possible interactions. Knowledge of these
processes will ensure that the analyst will be able to recognize and address problems if they arise.

13.3 Fusion Techniques
Sample decomposition through fusion is employed most often for samples that are difficult to
dissolve in acids such as soils, sludges, silicates, and some metal oxides. Fusion is accomplished
by heating a salt (the flux) mixed with an appropriate amount of sample. The mixture is heated to
a temperature above the melting point of the salt, and the sample is allowed to react in the molten
mixture. When the reaction is completed, the mixture is allowed to cool to room temperature.
The fused sample is then dissolved, and the analysis is continued. Any residue remaining may be
treated by repeating the fusion with the same salt, performing a fusion with a different salt, acid
treatment, or any combination of the three.
Decomposition of the sample matrix depends on the high temperatures required to melt a flux
salt and the ratio of the flux salt to the sample. For a fusion to be successful, the sample must
contain chemically bound oxygen as in oxides, carbonates, and silicates. Samples that contain no
chemically bound oxygen, such as sulfides, metals, and organics, must be oxidized before the
fusion process.
Samples to be fused should be oven-dried to remove moisture. Samples with significant amounts
of organic material are typically dry ashed or wet ashed before fusion. Solid samples are ground
to increase the surface area, allowing the fusion process to proceed more readily. The sample
must be mixed thoroughly with the flux in an appropriate ratio. Generally, the crucible should
never be more than half-filled at the outset of the fusion process. Fusions may be performed
using sand or oil baths on a hot plate, in a muffle furnace, or over a burner. Crucibles are made of
platinum, zirconium, nickel, or porcelain (Table 13.1). The choice of heat source and crucible
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Sample Dissolution
material generally depends on the salt used for the fusion.
During fusion, samples are heated slowly and evenly to prevent ignition of the sample before the
reaction with the molten salt can begin. It is especially important to raise the temperature slowly
when using a gas flame because the evolution of water and gases is a common occurrence at the
beginning of the fusion, and hence a source of spattering. The crucible can be covered with a lid
as an added precaution. Sand and oil baths provide the most even source of heat, but they are
difficult to maintain at very high temperatures. Muffle furnaces provide an even source of heat,
but when using them it is difficult to monitor the progress of the reaction and impossible to work
with the sample during the fusion. Burners are used often as a convenient heat source although
they make it difficult to heat the sample evenly.
TABLE 13.1 — Common fusion fluxes
Flux
(mp, EC)

Fusion
Temperature, EC

Type of
Crucible

Na2S2O7 (403E) or
K2S2O7 (419E)

Up to red heat

Pt, quartz,
porcelain

NaOH (321E)
or
KOH (404E)

450-600E

Na2CO3 (853) or
K2CO3 (903)

900-1,000E

Na2O2

600E

H3BO3

250E

Pt

Na2B4O7 (878E)

1,000-1,200E

Pt

Li2B4O7 (920E)
or
LiBO2 (845E)

1,000-1,100E

Pt, graphite

NH4HF2 (125E) NaF
(992E)
KF (857E)
900E
or
KHF2 (239E)
Source: Dean (1995) and Bock (1979).

Types of Sample Decomposed
For insoluble oxides and oxide-containing samples,
particularly those of Al, Be, Ta, Ti, Zr, Pu, and the
rare earths.

Ni, Ag, glassy For silicates, oxides, phosphates, and fluorides.
carbon
Ni
Pt for short
periods (use lid)
Ni; Ag, Au, Zr;
Pt (<500 EC)

For silicates and silica-containing samples (clays,
minerals, rocks, glasses), refractory oxides, quartz,
and insoluble phosphates and sulfates.
For sulfides; acid-insoluble alloys of Fe, Ni, Cr, Mo,
W, and Li; Pt alloys; Cr, Sn, and Zn minerals.
For analysis of sand, aluminum silicates, titanite,
natural aluminum oxide (corundum), and enamels.
For Al2O3; ZrO2 and zirconium ores, minerals of the
rare earths, Ti, Nb, and Ta, aluminum-containing
materials; iron ores and slags.
For almost anything except metals and sulfides. The
tetraborate salt is especially good for basic oxides and
some resistant silicates. The metaborate is better
suited for dissolving acidic oxides such as silica and
TiO2 and nearly all minerals.
For the removal of silicon, the destruction of silicates
and rare earth minerals, and the analysis of oxides of
Nb, Ta, Ti, and Zr.

Pt

The maximum temperature employed varies considerably and depends on the sample and the
flux. In order to minimize attack of the crucible and decomposition of the flux, excessive
temperatures should be avoided. Once the salt has melted, the melt is swirled gently to monitor
the reaction. The fusion continues until visible signs of reaction are completed (e.g., formation of
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gases, foaming, fumes). It is frequently difficult to decide when heating should be discontinued.
In ideal cases, a clear melt serves to indicate the completeness of sample decomposition. In other
cases, it is not as obvious, and the analyst must base the heating time on past experience with the
sample type.
The melt sometimes is swirled during cooling to spread it over the inside of the crucible. Thin
layers of salt on the sides of the crucible often will crack and flake into small pieces during
cooling. These small fragments are easier to remove and dissolve.
After the sample has returned to room temperature, the fused material is dissolved. The solvent is
usually warm water or a dilute acid solution, depending on the salt. For example, dilute acid
typically would not be used to dissolve a carbonate fusion because of losses to spray caused by
release of CO2. The aqueous solution from the dissolution of the fusion melt should be examined
carefully for particles of undissolved sample. If undissolved particles are present, they should be
separated from solution by centrifugation or filtration, and a second fusion should be performed.
Several types of materials are used for crucibles, but platinum, other metals (Ni, Zr, Ag), and
graphite are most common. Graphite crucibles are a cost-effective alternative to metal crucibles;
they are disposable, which eliminates the need for cleaning and the possibility of cross-sample
contamination. Graphite crucibles are chemically inert and heat-resistant, although they do
oxidize slowly at temperatures above 430 EC. Graphite is not recommended for extremely
lengthy fusions or for reactions where the sample may be reduced. Platinum is probably the most
commonly used crucible material. It is virtually unaffected by most of the usual acids, including
hydrofluoric, and it is attacked only by concentrated phosphoric acid at very high temperatures,
and by sodium carbonate. However, it dissolves readily in mixtures of hydrochloric and nitric
acids (aqua regia), nitric acid containing added chlorides, or chlorine water or bromine water.
Platinum offers adequate resistance toward molten alkali metal, borates, fluorides, nitrates, and
bisulfates. When using a platinum crucible, one should avoid using aqua regia, sodium peroxide,
free elements (C, P, S, Ag, Bi, Cu, Pb, Zn, Se, and Te), ammonium, chlorine and volatile
chlorides, sulfur dioxide, and gases with carbon content. Platinum crucibles can be cleaned in
boiling HNO3, by hand cleaning with sea sand or by performing a blank fusion with sodium
hydrogen sulfate.
Many kinds of salts are used in fusions. The lowest melting flux capable of reacting completely
with the sample is usually the optimum choice. Basic fluxes, such as the carbonates, the
hydroxides, and the borates, are used to attack acidic materials. Sodium or potassium nitrate may
be added to furnish an oxidizing agent when one is needed, as with the sulfides, certain oxides,
ferroalloys, and some silicate materials. The most effective alkaline oxidizing flux is sodium
peroxide; it is both a strong base and a powerful oxidizing agent. Because it is such a strong
alkali, sodium peroxide is often used even when no oxidant is required. Alternatively, acid fluxes
are the pyrosulfates, the acid fluorides, and boric acids. Table 13.1 lists several types of fusions,
examples of salts used for each type of fusion, and the melting points of the salts.
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SULFATE FUSION is useful for the conversion of ignited oxides to sulfates, but is generally an
ineffective approach for silicates. Sulfate fusion is particularly useful for BeO, Fe2O3, Cr2O3,
MoO3, TeO2, TiO2, ZrO2, Nb2O5, Ta2O5, PuO2, and rare earth oxides (Bock, 1979). Pyrosulfate
fusions are prepared routinely in the laboratory by heating a mixture of sodium or potassium
sulfate with a stoichiometric excess of sulfuric acid:
Na2SO4 + H2SO4 6 [2NaHSO4] 6 Na2S2O7 + H2O
Na2S2O7 6 Na2SO4 + SO38
Na2SO4 etc.
The rate of heating is increased with time until the sulfuric acid has volatilized and a clear
pyrosulfate fusion is obtained. A pyrosulfate melt can be reprocessed if necessary to achieve
complete sample dissolution. The analyst must distinguish between insoluble material that has
not yet or will not dissolve, and material that has precipitated during the final stages of a
prolonged pyrosulfate fusion. In the latter situation the fusion must be cooled, additional sulfuric
acid added, and the sample refused until the precipitated material redissolves and a clear melt is
obtained. Otherwise, the precipitated material will be extremely difficult, if not impossible, to
dissolve in subsequent steps. Platinum or quartz crucibles are recommended for this type of
fusion, with quartz being preferred for analysis of the platinum group metals. After the melt is
cooled and solidified, it should be dissolved in dilute sulfuric or hydrochloric acid rather than in
water to avoid hydrolysis and precipitation of Ti, Zr, etc. Niobium and tantalum may precipitate
even in the presence of more concentrated acid. In order to avoid precipitation of Nb or Ta,
concentrated sulfuric acid, tartaric acid, ammonium oxalate, hydrogen peroxide, or hydrofluoric
acid must be used. Mercury and the anions of volatile acids are largely volatilized during these
fusion procedures.
13.3.1 Alkali-Metal Hydroxide Fusions
Alkali metal hydroxide fusions are used for silicate analysis of ash and slag; for decomposition of
oxides, phosphates, and fluorides (Bock, 1979, pp. 102-108); and for dissolution of soils for
actinide analyses (Smith et al., 1995). Sodium hydroxide (NaOH) generally is used because of its
lower melting point, but potassium hydroxide (KOH) is just as effective. These fusions generally
are rapid, the melts are easy to dissolve in water, and the losses due to volatility are reduced
because of the low temperature of the melt. Nickel, silver, or glassy carbon crucibles are
recommended for this type of fusion. The maximum suggested temperature for nickel crucibles is
600 EC, but silver crucibles can be used up to 700 EC. Generally, crucibles made of platinum,
palladium, and their alloys should not be used with hydroxide fusions because the crucibles are
easily attacked in the presence of atmospheric oxygen. The weight ratio of fusion salt to sample
is normally 5-10:1. Typically, these fusions are carried out below red heat at 450 to 500 EC for
15 to 20 minutes, or sometimes at higher temperatures between 600 to 700 EC for 5 to 10
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minutes. The solidified melt dissolves readily in water; and therefore, this step may be carried out
directly in the crucible, or alternatively in a nickel dish. Under no circumstances should the
dissolution be carried out in a glass vessel because the resulting concentrated hydroxide solution
attacks glass quite readily.
FUSION WITH SODIUM CARBONATE (Na2CO3) is a common procedure for decomposing silicates
(clays, rocks, mineral, slags, glasses, etc.), refractory oxides (magnesia, alumina, beryllia,
zirconia, quartz, etc.), and insoluble phosphates and sulfates (Bogen, 1978). The fusion may
result in the formation of a specific compound such as sodium aluminate, or it may simply
convert a refractory oxide into a condition where it is soluble in hydrochloric acid—this is the
method of choice when silica in a silicate is to be determined, because the fusion converts an
insoluble silicate into a mixture that is easily decomposed by hydrochloric acid (“M” represents a
metal in the equations below):
MSiO3 + Na2CO3 6 Na2SiO3 + MCO3 (or MO + CO2),
followed by acidification to form a more soluble chloride salt,
Na2SiO3 + MCO3 + 4 HCl + x H2O 6 H2SiO3 · x H2O + MCl2 + CO2 + H2O + NaCl.
Carbonate fusions provide an oxidizing melt for the analysis of chromium, manganese, sulfur,
boron, and the platinum group metals. Organic material is destroyed, sometimes violently.
Na2CO3 generally is used because of its lower melting point. However, despite its higher melting
point and hygroscopic nature, K2CO3 is preferred for niobium and tantalum analyses because the
resulting potassium salts are soluble, whereas the analogous sodium salts are insoluble.
The required temperature and duration of the fusion depend on the nature of the sample as well
as particle size. In the typical carbonate fusion, 1 g of the powdered sample is mixed with 4 to 6 g
of sodium carbonate and heated at 900 to 1,000 EC for 10 to 30 minutes. Very refractory
materials may require heating at 1,200 EC for as long as 1 to 2 hours. Silica will begin to react at
500 EC, while barium sulfate and alumina react at temperatures above 700 EC. Volatility could
be a problem at these temperatures. Mercury and thallium are lost completely, while selenium,
arsenic, and iodine suffer considerable losses. Nonsilicate samples should be dissolved in water,
while silicate samples should be treated with acid (Bock, 1979).
Platinum crucibles are recommended for fusion of solid samples even though there is a 1 to 2 mg
loss of platinum per fusion. Attack on the crucible can be reduced significantly by covering the
melt with a lid during the fusion process, or virtually eliminated by working in an inert atmosphere. Moreover, nitrate is often added to prevent the reduction of metals and the subsequent
alloying with the platinum crucibles. The platinum crucibles may be seriously attacked by
samples containing high concentrations of Fe2+, Fe3+, Sn4+, Pb2+, and compounds of Sb and As,
because these ions are reduced easily to the metallic state and then form intermetallic alloys with
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platinum that are not easily dissolved in mineral acids. This problem is especially prevalent when
fusion is carried out in a gas flame. Porcelain crucibles are corroded rapidly and should be
discarded after a single use.
13.3.2 Boron Fusions
Fusions with boron compounds are recommended for analysis of sand, slag, aluminum silicates,
alumina (Al2O3), iron and rare earth ores, zirconium dioxide, titanium, niobium, and tantalum.
Relatively large amounts of flux are required for these types of fusions. The melts are quite
viscous and require swirling or stirring, so they should not be performed in a furnace. Platinum
crucibles should be used for these fusions because other materials are rapidly attacked by the
melt, even though some platinum is lost in each fusion.
BORIC ACID (H3BO3) can be used to fuse a number of otherwise inert substances such as sand,
aluminum silicates, titanite, natural aluminum oxide (corundum), and enamels. Boric acid
fusions generally require 4 to 8 times as much reagent as sample. Initially, the mixture should be
heated cautiously while water is being driven off, then more strongly until gas evolution is
completed, and then more vigorously if the sample has yet to be fully decomposed. Normally, the
procedure is complete within 20 to 30 minutes. The cooled and solidified melt usually is
dissolved in dilute acid. Additionally, boric acid has one great advantage over all other fluxes in
that it can be completely removed by addition of methanol and subsequent volatilization of the
methyl ester.
Because MOLTEN SODIUM TETRABORATE (Na2B4O7) dissolves so many inorganic compounds, it is
an important analytical tool for dissolving very resistant substances. Fusions with sodium tetraborate alone are useful for Al2O3, ZrO2 and zirconium ores, minerals of the rare earths, titanium,
niobium, and tantalum, aluminum-containing materials, and iron ores and slags (Bock, 1979).
Relatively large amounts of borax are mixed with the sample, and the fusion is carried out at a
relatively high temperature (1,000 to 1,200 EC) until the melt becomes clear. Thallium, mercury,
selenium, arsenic, and the halogens are volatilized under these conditions. Boric acid can be
removed from the melt as previously described. By dissolving the melt in dilute hydrofluoric
acid, calcium, thorium, and the rare earths can be separated from titanium, niobium, and tantalum
as insoluble fluorides.
LITHIUM METABORATE (Li2B4O7) is well-suited for dissolving basic oxides, such as alumina
(Al2O3), quicklime (CaO), and silicates. Platinum dishes are normally used for this type of fusion,
but occasionally graphite crucibles are advantageous because they can be heated rapidly by
induction, and because they are not wetted by Li2B4O7 melts. The fusion melt typically is
dissolved in dilute acid, usually nitric but sometimes sulfuric. When easily hydrolyzed metal ions
are present, dissolution should be carried out in the presence of ethylenediamine tetracetic acid
(EDTA) or its di-sodium salt in 0.01 M HCl (Bock, 1979).
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LITHIUM METABORATE (LiBO2), or a mixture of the meta- and tetraborates, is a more basic
flux and is better for dissolving highly acidic oxides or very insoluble ones, such as silica (SiO2)
or rutile (TiO2). The metaborate is, however, suitable for dissolving all metal oxides. After the
melt of sample and metaborate are dissolved, hydrogen peroxide should be used to maintain the
titanium in solution.
13.3.3 Fluoride Fusions
Fluoride fusions are used for the removal of silicon, the destruction of silicates and rare earth
minerals, and the analysis of oxides of niobium, tantalum, titanium, and zirconium. Sill et al.
(1974) and Sill and Sill (1995) describe a method using potassium fluoride/potassium pyrosulfate
fusion for determining alpha-emitting nuclides in soil (see Section 13.8, “Comparison of Total
Dissolution and Acid Leaching”). Sulcek and Povondra (1989) describe the isolation of the rare
earth elements and thorium from silicate materials and their minerals, especially monazite,
through potassium hydrofluoride fusion. The silicate matrix is first degraded by evaporation with
HF, then the residue is fused with tenfold excess flux, and finally the melt is digested with dilute
acid. The resulting fluorides (rare earths + Th + Ca + U) are filtered out, dissolved, and further
separated.
Platinum crucibles are recommended for fluoride fusions. Silicon, boron, lead, and polonium are
volatilized during these fusion procedures, and if the temperature is high enough, some
molybdenum, tantalum, and niobium also are lost. Residual fluoride can be a problem for
subsequent analysis of many elements such as aluminum, tin, beryllium, and zirconium. This
excess fluoride usually is removed by evaporation with sulfuric acid.
13.3.4 Sodium Hydroxide Fusion
Burnett et al. (1997) presented a technique that employs sodium hydroxide as the fusion agent in
a 5:1 ratio to the soil. The fusion is performed in an alumina crucible, and deioinized water is
added to the resultant cake. Sufficient iron exists in most samples to from an Fe(OH)3 scavenging
precipitate for the actinides. The addition of sodium formaldehyde sulfoxylate (“Rongalite”)
ensures all actinides are in the +4 or +3 valence state.

13.4 Wet Ashing and Acid Dissolution Techniques
“Wet ashing” and “acid dissolution” are terms used to describe sample decomposition using hot,
concentrated acid solutions. Because many inorganic matrices such as oxides, silicates, nitrides,
carbides, and borides can be difficult to dissolve completely, geological or ceramic samples can
be particularly challenging. Therefore, different acids are used alone or in combination to decompose specific compounds that may be present in the sample. Few techniques will decompose all
types of samples completely. Many decomposition procedures use wet ashing to dissolve the
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major portion of the sample but leave a minor fraction as residue. Whether or not this residue
requires additional treatment (by wet ashing or fusion) depends on the amount of residue and
whether it is expected to contain the radionuclides of interest. The residue should not be
discarded until all of the results have been reviewed and determined to be acceptable.
13.4.1 Acids and Oxidants
Numerous acids are commonly used in wet ashing procedures. Table 13.2 lists several acids and
the types of compounds they generally react with during acid dissolution. The electromotive
force series (Table 13.3) is a summary of oxidation-reduction half-reactions arranged in
decreasing oxidation strength and is also useful in selecting reagent systems (Dean, 1995).
TABLE 13.2 — Examples of acids used for wet ashing
Acid

Typical Uses

Hydrofluoric Acid, HF

Removal of silicon and destruction of silicates; dissolves oxides of Nb, Ta,
Ti, and Zr, and Nb, and Ta ores.

Hydrochloric Acid, HCl

Dissolves many carbonates, oxides, hydroxides, phosphates, borates, and
sulfides; dissolves cement.

Hydrobromic Acid, HBr

Distillation of bromides (e.g., As, Sb, Sn, Se).

Hydroiodic Acid, HI

Effective reducing agent; dissolves Sn+4 oxide and Hg+2 sulfide.

Sulfuric Acid, H2SO4

Dissolves oxides, hydroxides, carbonates, and various sulfide ores; hot
concentrated acid will oxidize most organic compounds.

Phosphoric Acid, H3PO4

Dissolves Al2O3, chrome ores, iron oxide ores, and slag.

Nitric Acid, HNO3

Oxidizes many metals and alloys to soluble nitrates; organic material
oxidized slowly.

Perchloric Acid, HClO4

Extremely strong oxidizer; reacts violently or explosively to oxidize organic
compounds; attacks nearly all metals.

The table allows one to predict which metals will dissolve in nonoxidizing acids, such as hydrochloric, hydrobromic, hydrofluoric, phosphoric, dilute sulfuric, and dilute perchloric acid The
dissolution process is simply a replacement of hydrogen by the metal (Dean, 1995). In practice,
however, what actually occurs is influenced by a number of factors, and the behavior of the
metals cannot be predicted from the potentials alone. Generally, metals below hydrogen in Table
13.3 displace hydrogen and dissolve in nonoxidizing acids with the evolution of hydrogen.
Notable exceptions include the very slow dissolution by hydrochloric acid of lead, cobalt, nickel,
cadmium, and chromium. Also, lead is insoluble in sulfuric acid because of the formation of a
surface film of insoluble lead sulfate.

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TABLE 13.3 — Standard reduction potentials of
selected half-reactions at 25 EC
Half-Reaction
E0 (volts)
1.980
Ag2+ + e! 6 Ag+ . . . . . . . . . . . . . . . . . . . . . . . . .
1.96
S2O82- + 2e! 6 2SO42- . . . . . . . . . . . . . . . . . . . . .
Ce4+ + e! 6 Ce3+ . . . . . . . . . . . . . . . . . . . . . . . . .
1.72
MnO4! + 4H+ + 3e! 6 MnO2 (s) + 2H2O . . . . . .
1.70
+
!
2HClO + 2H + 2e 6 Cl2 + 2H2O . . . . . . . . . . .
1.630
2HBrO + 2H+ + 2e! 6 Br2 + 2H2O . . . . . . . . . . .
1.604
1.593
NiO2 + 4H+ + 2e! 6 Ni2+ + 2H2O . . . . . . . . . . . .
Bi2O4 (bismuthate) + 4H+ + 2e! 6 2BiO+ + 2H2O 1.59
MnO4! + 8H+ + 5e! 6 Mn2+ + 4H2O . . . . . . . . . .
1.51
2BrO3! + 12H+ + 10e! 6 Br2 + 6H2O . . . . . . . . .
1.478
PbO2 + 4H+ + 2e! 6 Pb2+ + 2H2O . . . . . . . . . . . .
1.468
+
!
3+
21.36
Cr2O7 + 14H + 6e 6 2Cr + 7H2O . . . . . . . .
Cl2 + 2e! 6 2Cl! . . . . . . . . . . . . . . . . . . . . . . . . .
1.3583
2HNO2 + 4H+ + 4e! 6 N2O + 3H2O . . . . . . . . . .
1.297
1.23
MnO2 + 4H++ 2e! 6 Mn2+ + 2H2O . . . . . . . . . . .
O2 + 4H+ + 4e! 6 2H2O . . . . . . . . . . . . . . . . . . .
1.229
1.201
ClO4! + 2H+ + 2e! 6 ClO!3 + H2O . . . . . . . . . . .
2IO3! + 12H+ + 10e! 6 I2 + 3H2O . . . . . . . . . . . .
1.19
+
!
N2O4 + 2H + 2e 6 2HNO2 . . . . . . . . . . . . . . . .
1.07
1.07
2ICl!2 + 2e! 6 4Cl! + I2 . . . . . . . . . . . . . . . . . . .
Br2 (aq) + 2e! 6 2Br- . . . . . . . . . . . . . . . . . . . . .
1.065
1.039
N2O4 + 4H+ + 4e! 6 2NO + 2H2O . . . . . . . . . . .
HNO2 + H+ + e! 6 NO + H2O . . . . . . . . . . . . . . .
0.996
NO3! + 4H+ + 3e! 6 NO + 2H2O . . . . . . . . . . . .
0.957
+
!
!
0.94
NO3 + 3H + 2e 6 HNO2 + H2O . . . . . . . . . . .
2Hg2+ + 2e! 6 Hg22+ . . . . . . . . . . . . . . . . . . . . . .
0.911
0.861
Cu2+ + I! + e! 6 CuI (s) . . . . . . . . . . . . . . . . . . .
OsO4 (s) + 8H+ + 8e! 6 Os + 4H2O . . . . . . . . . .
0.84
Ag+ + e! 6 Ag . . . . . . . . . . . . . . . . . . . . . . . . . . .
0.7991
Hg22+ + 2e! 6 2Hg . . . . . . . . . . . . . . . . . . . . . . .
0.7960
Fe3+ + e! 6 Fe2+ . . . . . . . . . . . . . . . . . . . . . . . . . .
0.771
+
!
0.739
H2SeO3 + 4H + 4e 6 Se + 3H2O . . . . . . . . . . .
HN3 + 11H+ + 8e! 6 3NH4+ . . . . . . . . . . . . . . . .
0.695
O2 + 2H+ + 2e- 6 H2O2 . . . . . . . . . . . . . . . . . . . .
0.695
Ag2SO4 + 2e! 6 2Ag + SO42- . . . . . . . . . . . . . . .
0.654
0.654
Cu2+ + Br! + e! 6 CuBr (s) . . . . . . . . . . . . . . . . .
2HgCl2 + 2e! 6 Hg2Cl2 (s) + 2Cl! . . . . . . . . . . .
0.63
Sb2O5 + 6H+ + 4e! 6 2SbO+ + 3H2O . . . . . . . . .
0.605
+
!
0.560
H3AsO4 + 2H + 2e 6 HAsO2 + 2 H2O . . . . . . .
TeOOH+ + 3H+ + 4e! 6 Te + 2H2O . . . . . . . . . .
0.559
Cu2+ + Cl! + e! 6 CuCl (s) . . . . . . . . . . . . . . . . .
0.559

Half-Reaction
I!3 + 3e! 6 3I! . . . . . . . . . . . . . . . . . .
I2 + 2e! 6 2I! . . . . . . . . . . . . . . . . . .
Cu+ + e! 6 Cu . . . . . . . . . . . . . . . . . .
4H2SO3 + 4H+ + 6e! 6 S4O62- + 6H2O
Ag2CrO4 + 2e! 6 2Ag + CrO42- . . . .
2H2SO3 + 2H+ + 4e! 6 S2O32- + 3H2O
UO2+ + 4H+ + e! 6 U4+ + 2H2O . . . .
Cu2+ + 2e! 6 Cu . . . . . . . . . . . . . . . .
VO2+ + 2H+ + e! 6 V3+ + H2O . . . . .
BiO+ + 2H+ + 3e! 6 Bi + H2O . . . . .
UO22+ + 4H+ + 2e! 6 U4+ + 2H2O . . .
Hg2Cl2 (s) + 2e! 6 2Hg + 2Cl! . . . . .
AgCl (s) + e! 6 Ag + Cl! . . . . . . . . .
SbO+ + 2H+ + 3e! 6 Sb + H2O . . . . .
CuCl32- + e! 6 Cu + 3Cl! . . . . . . . . .
SO42- + 4H+ + 2e! 6 H2SO3 + H2O . .
Sn4+ + 2e! 6 Sn2+ . . . . . . . . . . . . . . .
CuCl + e! 6 Cu + Cl! . . . . . . . . . . . .
TiO2+ + 2H+ + e- 6 Ti3+ + H2O . . . . .
S4O62- + 2e! 6 2S2O32- . . . . . . . . . . . .
2H+ + 2e! 6 H2 . . . . . . . . . . . . . . . . .
Hg2I2 (s) + 2e! 6 2Hg + 2I! . . . . . . .
Pb2+ + 2e! 6 Pb . . . . . . . . . . . . . . . . .
Sn2+ + 2e! 6 Sn . . . . . . . . . . . . . . . . .
AgI (s) + e! 6 Ag + I! . . . . . . . . . . .
V3+ + e! 6 V2+ . . . . . . . . . . . . . . . . . .
Ni2+ + 2e! 6 Ni . . . . . . . . . . . . . . . . .
Co2+ + 2e! 6 Co . . . . . . . . . . . . . . . .
PbSO4 + 2e! 6 Pb + SO42- . . . . . . . .
Cd2+ + 2e! 6 Cd . . . . . . . . . . . . . . . .
Cr3+ + e! 6 Cr2+ . . . . . . . . . . . . . . . . .
Fe2+ + 2e! 6 Fe . . . . . . . . . . . . . . . . .
H3PO3 + 2H+ + 2e! 6 HPH2O2 + H2O
U4+ + e! 6 U3+ . . . . . . . . . . . . . . . . . .
Zn2+ + 2e! 6 Zn . . . . . . . . . . . . . . . .
Mn2+ + 2e! 6 Mn . . . . . . . . . . . . . . .
Al3+ + 3e! 6 Al . . . . . . . . . . . . . . . . .
Mg2+ + 2e! 6 Mg . . . . . . . . . . . . . . .
Na+ + e! 6 Na . . . . . . . . . . . . . . . . . .
K+ + e! 6 K . . . . . . . . . . . . . . . . . . . .
Li+ + e! 6 Li . . . . . . . . . . . . . . . . . . .
3N2 + 2H+ + 2e! 6 2HN3 . . . . . . . . .

E0 (volts)
0.536
0.536
0.53
0.507
0.449
0.400
0.38
0.340
0.337
0.32
0.27
0.2676
0.2223
0.212
0.178
0.158
0.15
0.121
0.100
0.08
0.0000
-0.0405
-0.125
-0.136
-0.1522
-0.255
-0.257
-0.277
-0.3505
-0.4025
-0.424
-0.44
-0.499
-0.52
-0.7626
-1.18
-1.67
-2.356
-2.714
-2.925
-3.045
-3.1

Source: Dean, 1995.
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Oxidizing acids, such as nitric acid, hot concentrated sulfuric acid, or hot concentrated perchloric
acid, are used to dissolve metals whose E0 values are greater than hydrogen. For nitric acid, the
potential of the nitrate ion-nitric oxide couple can be employed as a rough estimate of the solvent
power. For aqua regia, the presence of free chlorine ions allows one to make predictions based
upon the potential of the chlorine-chloride couple, although NOCl also plays a significant role.
Some oxidizing acids exhibit a passivating effect with transition elements such as chromium and
pure tungsten, resulting in a very slow attack because of the formation of an insoluble surface
film of the oxide in the acid (Bogen, 1978). Moreover, oxides are often resistant to dissolution in
oxidizing acids and, in fact, dissolve much more readily in nonoxidizing acids. A common
example is ferric oxide, which is readily soluble in hydrochloric acid but is relatively inert in
nitric acid.
However, insoluble oxides of the lower oxidation states of an element sometime dissolve in
oxidizing acids with concurrent oxidation of the element. For example, UO 2 and U3O8 dissolve
readily in nitric acid to produce a solution of uranyl ion (UO2+2).
HYDROFLUORIC ACID. The most important property of HF is its ability to dissolve silica and
other silicates. For example:
SiO2 + 6HF 6 H2SiF6 + 2H2O
whereby the fluorosilicic acid formed dissociates into gaseous silicon tetrafluoride and hydrogen
fluoride upon heating:
H2SiF6 6 SiF48 + 2HF
HF also exhibits pronounced complexing properties that are widely used in analytical chemistry.
Hydrofluoric acid prevents the formation of sparingly soluble hydrolytic products in solution,
especially of compounds of elements from the IV to VI groups of the periodic table (Sulcek and
Povondra, 1989). In the presence of fluoride, soluble hydrolytic products that are often polymeric
depolymerize to form reactive monomeric species suitable for further analytical operations.
Formation of colloidal solutions is avoided and the stability of solutions is increased even with
compounds of elements that are hydrolyzed easily in aqueous solution (e.g., Si, Sn, Ti, Zr, Hf,
Nb, Ta, and Pa).
HF should never be used or stored in glass, or porcelain containers. Digestion in platinum
containers is preferred, and Teflon™ is acceptable as long as the temperature does not exceed
250 EC. This would occur only with HF if the mix were taken to dryness, because the constant
boiling azeotrope is 112 EC. HF works most effectively when used alone, as all other acids or
oxidizing agents used are less volatile than HF and would cause the HF concentration to be
decreased at elevated temperatures. HF is most effective when used on a solid residue. Samples
should be ground to a fine powder to increase the surface area and moistened with a minimal
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Sample Dissolution
amount of water to prevent losses as dust and spray when the acid is added to the sample. After
the addition of HF, the sample may be allowed to react overnight to dissolve the silicates.
However, heating the solution to 80 EC will allow reaction to occur within 1-2 hours. Because it
is such a strong complexing agent, excess fluoride ion can cause problems with many separation
methods. Residual fluoride is usually removed by evaporation to fumes in a low-volatility acid
(e.g., H2SO4, HNO3, HClO4) or, in extreme cases, excess fluoride ion can be removed by fusing
the residue with boric acid or sodium tetraborate. The fluorides are converted to BF3 that is then
removed by evaporation.
HYDROCHLORIC ACID (HCl) is one of the most widely used acids for sample dissolution because
of the wide range of compounds it reacts with and the low boiling point of the azeotrope
(110 EC); after a period of heating in an open container, a constant boiling 6M solution remains.
HCl forms strong complexes with Au+3, Ti+3, and Hg+2. The concentrated acid will also complex
Fe+3, Ga+3, In+3, and Sn+4. Most chloride compounds are readily soluble in water except for silver
chloride, mercury chloride, titanium chloride, and lead chloride. HCl can be oxidized to form
chlorine gas by manganese dioxide, permanganate, and persulfate. While HCl dissolves many
carbonates, oxides, hydroxides, phosphates, borates, sulfides, and cement, it does not dissolve the
following:
•
•
•
•
•
•
•

Most silicates or ignited oxides of Al, Be, Cr, Fe, Ti, Zr, or Th;
Oxides of Sn, Sb, Nb, or Ta;
Zr phosphate;
Sulfates of Sr, Ba, Ra, or Pb;
Alkaline earth fluorides;
Sulfides of Hg; or
Ores of Nb, Ta, U, or Th.

The dissolution behavior of specific actinides by hydrochloric acid is discussed by Sulcek and
Povondra (1989):
“The rate of decomposition of oxidic uranium ores depends on the U(VI)/U(+4) ratio.
The so-called uranium blacks with minimal contents of U(+4) are even dissolved in dilute
hydrochloric acid. Uraninite (UO2) requires an oxidizing mixture of hydrochloric acid
with hydrogen peroxide, chlorate, or nitric acid for dissolution. Uranium and thorium
compounds cannot be completely leached from granites by hydrochloric acid. Natural and
synthetic thorium dioxides are highly resistant toward hydrochloric acid and must be
decomposed in a pressure vessel. Binary phosphates of uranyl and divalent cations, e.g.,
autunite and tobernite, are dissolved without difficulties. On the other hand, phosphates
of thorium, tetravalent uranium, and the rare earths (monazite and xenotime) are only
negligibly attacked, even with the concentrated acid.”
As+3, Sb+3, Ge+3, and Se+4 are volatilized easily in HCl solutions, while Hg+2, Sn+4, and Rh(VII)
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Sample Dissolution
are volatilized in the latter stages of evaporation. Glass is the preferred container for HCl
solutions.
HYDROBROMIC ACID (HBr) has no important advantages over HCl for sample dissolution. HBr
forms an azeotrope with water containing 47.6 percent by weight of HBr, boiling at 124.3 EC.
HBr is used to distill off volatile bromides of arsenic, antimony, tin, and selenium. HBr can also
be used as a complexing agent for liquid-liquid extractions of gold, titanium, and indium.
HYDROIODIC ACID (HI) is readily oxidized. Solutions often appear yellowish-brown because of
the formation of the triiodide complex (I!3). HI is most often used as a reducing agent during
dissolutions. HI also dissolves Sn+4 oxide, and complexes and dissolves Hg+2 sulfide. HI forms an
azeotrope with water containing 56.9 percent by weight of HI, boiling at 127 EC.
SULFURIC ACID (H2SO4) is another widely used acid for sample decomposition. Part of its
effectiveness is due to its high boiling point (about 340 EC). Oxides, hydroxides, carbonates, and
sulfide ores can be dissolved in H2SO4. The boiling point can be raised by the addition of sodium
or potassium sulfate to improve the attack on ignited oxides, although silicates will still not
dissolve. H2SO4 is not appropriate when calcium is a major constituent because of the low
solubility of CaSO4. Other inorganic sulfates are typically soluble in water, with the notable
exceptions of strontium, barium, radium, and lead.
Non-fuming H2SO4 does not exhibit oxidizing properties, but the concentrated acid will dissolve
many elements and react with almost all organic compounds. Concentrated sulfuric acid is a
powerful dehydrating agent. Its action on organic materials is a result of removing OH and H
groups (to form water) from adjacent carbon atoms. This forms a black char (residue) that is not
easily dissolved using wet-ashing techniques. Moreover, because of the high boiling point of
H2SO4, there is an increased risk of losses because of volatilization. Iodine can be distilled
quantitatively, and boron, mercury, selenium, osmium, ruthenium, and rhenium may be lost to
some extent. The method of choice is to oxidize the organic substances with HNO3, volatilize the
nitric acid, add H2SO4 until charred, followed by HNO3 again, repeating the process until the
sample will not char with either HNO3 or H2SO4. Dissolution is then continued with HClO4.
Glass, quartz, platinum, and porcelain are resistant to H2SO4 up to the boiling point. Teflon™
should not be used above 250 EC, and, therefore, it is not recommended for applications
involving concentrated H2SO4 that require elevated temperature.
Glass, quartz, platinum, and porcelain are resistant to H2SO4 up to the boiling point. Teflon
decomposes at 300 EC, below the boiling point, and, therefore, is not recommended for
applications involving H2SO4 that require elevated temperature.
PHOSPHORIC ACID (H3PO4) seldom is used for wet ashing because the residual phosphates
interfere with many separation procedures. H3PO4 attacks glass, although glass containers are
usually acceptable at temperatures below 300 EC. Alumina, chromium ores, iron oxide ores, and
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Sample Dissolution
slags can be dissolved in H3PO4. The acid also has been used to dissolve silicates selectively
without attacking quartz.
NITRIC ACID (HNO3) is one of the most widely used oxidizing acids for sample decomposition.
Most metals and alloys are oxidized to nitrates, which are usually very soluble in water, although
many metals exhibit a pronounced tendency to hydrolyze in nitric acid solution. Nitric acid does
not attack gold, hafnium, tantalum, zirconium, and the metals of the platinum group (except
palladium). Aluminum, boron, chromium, gallium, indium, niobium, thorium, titanium, calcium,
magnesium, and iron form an adherent layer of insoluble oxide when treated with HNO3, thereby
passivating the metal surface. However, calcium, magnesium, and iron will dissolve in more
dilute acid.
Complexing agents (e.g., Cl!, F!, citrate, tartrate) can assist HNO3 in dissolving most metals. For
example, Sulcek and Povondra (1989) describe the decomposition of thorium and uranium
dioxides in nitric acid, which is catalytically accelerated by the addition of 0.05 to 0.1 M HF.
They also report that a solid solution of the mixed oxides (Pu, U)O2 or PuO2 ignited at
temperatures below 800 EC behaves analogously.
Although nitric acid is a good oxidizing agent, it usually boils away before sample oxidation is
complete. Oxidation of organic materials proceeds slowly and is usually accomplished by
repeatedly heating the solution to HNO3 fumes. Refluxing in the concentrated acid can help
facilitate the treatment, but HNO3 is seldom used alone to decompose organic materials.
PERCHLORIC ACID (HClO4). Hot concentrated solutions of HClO4 act as a powerful oxidizer, but
dilute aqueous solutions are not oxidizing. Hot concentrated HClO4 will attack nearly all metals
(except gold and platinum group metals) and oxidize them to the highest oxidation state, except
for lead and manganese, which are oxidized only to the +2 oxidation state. Perchloric acid is an
excellent solvent for stainless steel, oxidizing the chromium and vanadium to the hexavalent and
pentavalent acids, respectively. Many nonmetals also will react with HClO4. Because of the
violence of the oxidation reactions, HClO4 is rarely used alone for the destruction of organic
materials. H2SO4 or HNO3 are used to dilute the solution and break down easily oxidized material
before HClO4 becomes an oxidizer above 160 EC.
The concentrated acid is a dangerous oxidant that can explode violently. The following are
examples of some reactions with HClO4 that should never be attempted:
•
•
•
•
•
•

Heating bismuth metal and alloys with concentrated acid.
Dissolving metals (e.g., steel) in concentrated acid when gaseous hydrogen is heated.
Heating uranium turnings or powder in concentrated acid.
Heating finely divided aluminum and silicon in concentrated acid.
Heating antimony or Sb+3 compounds in HClO4.
Mixing HClO4 with hydrazine or hydroxylamine.

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Sample Dissolution
•
•
•
•
•
•
•

Mixing HClO4 with hypophosphates.
Mixing HClO4 with fats, oils, greases, or waxes.
Evaporating solutions of metal salts to dryness in HClO4.
Evaporating alcoholic filtrates after collection of KClO4 precipitates.
Heating HClO4 with cellulose, sugar, and polyhydroxy alcohols.
Heating HClO4 with N-heterocyclic compounds.
Mixing HClO4 with any dehydrating agent.

Perchloric acid vapor should never be allowed to contact organic materials such as rubber
stoppers. The acid should be stored only in glass bottles. Splashed or spilled acid should be
diluted with water immediately and mopped up with a woolen cloth, never cotton. HClO4 should
only be used only in specially designed fume hoods incorporating a washdown system.
Acid dissolutions involving HClO4 should only be performed by analysts experienced in working
with this acid. When any procedure is designed, the experimental details should be recorded
exactly. These records are used to develop a detailed standard operating procedure that must be
followed exactly to ensure the safety of the analyst (Schilt, 1979).
AQUA REGIA. One part concentrated HNO3 and three parts concentrated HCl (by volume) are
combined to form aqua regia:
3HCl + HNO3 6 NOCl + Cl2 + 2H2O
However, the interaction of these two acids is much more complex than indicated by this simple
equation. Both the elemental chlorine and the trivalent nitrogen of the nitrosyl chloride exhibit
oxidizing effects, as do other unstable products formed during the reaction of these two acids.
Coupled with the catalytic effect of Cl2 and NOCl, this mixture combines the acidity and
complexing power of the chloride ions. The solution is more effective if allowed to stand for 10
to 20 minutes after it is prepared.
Aqua regia dissolves sulfides, phosphates, and many metals and alloys including gold, platinum,
and palladium. Ammonium salts are decomposed in this acid mixture. Aqua regia volatilizes
osmium as the tetroxide; has little effect on rhodium, iridium, and ruthenium; and has no effect
on titanium. Oxidic uranium ores with uraninite and synthetic mixed oxides (U3O8) are dissolved
in aqua regia, with oxidation of the U+4 to UO2+2 ions (Sulcek and Povondra, 1989). However,
this dissolution procedure is insufficient for poor ores; the resistant, insoluble fraction must be
further attacked (e.g., by sodium peroxide or borate fusion) or by mixed-acid digestion with HF,
HNO3, and HClO4.
Oxysalts, such as KMnO4 (potassium permanganate) and K2Cr2O7 (potassium dichromate), are
commonly not used to solubilize or wet ash environmental samples for radiochemical analysis
because of their limited ability to oxidize metals and the residue that they leave in the sample
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Sample Dissolution
mixture. These oxysalts are more commonly used to oxidize organic compounds.
POTASSIUM PERMANGANATE (KMnO4) is a strong oxidizer whose use is limited primarily to the
decomposition of organic substances and mixtures, although it oxidizes metals such as mercury
to the ionic form. Oxidation can be performed in an acid, neutral, or basic medium; near-neutral
or basic solutions produce an insoluble residue of manganese dioxide (MnO2) that can be
removed by filtration. Oxidation in acid media leaves the Mn+2 ion in solution, which might
interfere with additional chemical procedures or analyses. Extreme caution must be taken when
using this reagent because KMnO4 reacts violently with some organic substances such as acetic
acid and glycerol, with some metals such as antimony and arsenic, and with common laboratory
reagents such as hydrochloric acid and hydrogen peroxide.
POTASSIUM DICHROMATE (K2Cr2O7) is a strong oxidizing agent for organic compounds but is not
as strong as KMnO4. K2Cr2O7 has been used to determine carbon and halogen in organic
materials, but the procedure is not used extensively. K2Cr2O7 is commonly mixed with sulfuric
acid and heated as a strong oxidizing agent to dissolve carbonaceous compounds. The Cr+3 ion
remains after sample oxidation and this might interfere with other chemical procedures or
analyses. K2Cr2O7 can react violently with certain organic substances such as ethanol and might
ignite in the presence of boron. Caution also must be observed in handling this oxidizing agent
because of human safety concerns, particularly with the hexavalent form of chromium.
SODIUM BROMATE (NaBrO3) is an oxidizing agent for organic compounds but is not used for
metals. Unlike KMnO4 and K2Cr2O7, the bromate ion can be removed from solution after sample
oxidation by boiling with excess HCl to produce water and Br 2. Caution must be observed when
using this oxidizing agent because it can react violently with some organic and inorganic
substances.
13.4.2 Acid Digestion Bombs
Some materials that would not be totally dissolved by acid digestion in an open vessel on a
hotplate, can be completely dissolved in an acid digestion bomb. These pressure vessels hold
strong mineral acids or alkalies at temperatures well above normal boiling points, thereby
allowing one to obtain complete digestion or dissolution of samples that would react slowly or
incompletely at atmospheric pressure. Sample dissolution is obtained without losing volatile
elements and without adding contaminants from the digestion vessel. Ores, rock samples, glass
and other inorganic samples can be dissolved quickly using strong mineral acids such as HF,
HCl, H2SO4, HNO3, or aqua regia.
These sealed pressure vessels are lined with Teflon™, which offers resistance to cross-contamination between samples and to attack by HF. In all reactions, the bomb must never be completely
filled; there must be adequate vapor space above the contents. When working with inorganic
materials, the total volume of sample plus reagents must never exceed two-thirds of the capacity
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Sample Dissolution
of the bomb. Moreover, many organic materials can be treated satisfactorily in these bombs, but
critical attention must be given to the nature of the sample as well to possible explosive reactions
with the digestion media.

13.5 Microwave Digestion
Microwave energy as a heat source for sample digestion was first described more than 20 years
ago (Abu-Samra et al., 1975). Its popularity is derived from the fact that it is faster, cleaner, more
reproducible, and more accurate than traditional hot-plate digestion. However, until recently, this
technology has had limited application in the radiochemical laboratory because of constraints on
sample size resulting from vessel pressure limitations. Because of this drawback, microwave
dissolution was not practical for many radiochemical procedures where larger sample sizes are
dictated to achieve required detection limits. However, recent advances in vessel design and
improved detection methods, such as ICP-MS (inductively coupled plasma-mass spectrometry)
and ion chromatography have eliminated this disadvantage, and microwave dissolution is an
important radiochemical tool (Smith and Yaeger, 1996; Alvarado et al., 1996). A series of
articles in Spectroscopy describes recent advances in microwave dissolution technology
(Kammin and Brandt, 1989; Grillo, 1989 and 1990; Gilman and Engelhardt, 1989; Lautenschlager, 1989; Noltner et al., 1990), and Dean (1995) presents a synopsis of current microwave
theory and technology. Kingston and Jassie (1988) and Kingston and Haswell (1997) are other
excellent resources for this topic.
The American Society for Testing and Materials (ASTM) has issued several protocols for various
media. ASTM D5258 describes the decomposition of soil and sediment samples for subsequent
analyte extraction; ASTM D4309 addresses the decomposition of surface, saline, domestic, and
industrial waste water samples; and ASTM D5513 covers the multistage decomposition of
samples of cement raw feed materials, waste-derived fuels, and other industrial feedstreams for
subsequent trace metal analysis. A method for acid digestion of siliceous and organically based
matrices is given in EPA (1996).
There are various microwave instruments that may be satisfactory depending on sample
preparation considerations. The three main approaches to microwave dissolution are: focused
open-vessel, low-pressure closed-vessel, and high-pressure closed-vessel. Each has certain
advantages and disadvantages and the choice of system depends upon the application.
13.5.1 Focused Open-Vessel Systems
A focused open-vessel system has no oven but consists of a magnetron to generate microwaves, a
waveguide to direct and focus the microwaves and a cavity to contain the sample (Grillo, 1989).
Because of the open-vessel design, there is no pressure buildup during processing, and reagents
may be added during the digestion program. These systems are quite universal in that any reagent
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Sample Dissolution
and any type of vessel (glass, Perfluoroalcoholoxil™ [PFA], or quartz) can be used.
The waveguide ensures that energy is directed only at the portion of the vessel in the path of the
focused microwaves thereby allowing the neck of the vessel and refluxer to remain cool and
ensuring refluxing action. Because of this refluxing action, the system maintains all elements,
even selenium and mercury. The focused microwaves cause solutions to reach higher
temperatures faster than with conventional hotplates or block-type digesters and do so with
superior reproducibility. An aspirator removes excess acid vapors and decomposition gases.
Depending on the system, up to 20 g of solids or 50 to 100 mL of liquids can be digested within
10 to 30 minutes on average.
13.5.2 Low-Pressure, Closed-Vessel Systems
These systems consist of a microwave oven equipped with a turntable, a rotor to hold the sample
vessels, and a pressure-control module (Grillo, 1990). The PFA vessels used with these systems
are limited to approximately 225 EC, and, therefore, low-boiling reagents or mixtures of reagents
should be used. Waste is minimized in these systems because smaller quantities of acid are
required. Moreover, because little or no acid is lost during the digestion, additional portions of
acid may not be required and blank values are minimized. Additionally, these sealed vessels are
limited to 100 to 300 psi (689 to 2,068 kPa), depending on the model thereby limiting the size of
organic samples utilized. However, inorganic materials such as metals, water and waste waters,
minerals, and most soils and sediments are easily digested without generating large amounts of
gaseous by-products. Typical sample sizes are on the order of 0.5 g for solids and 45 mL for
aqueous samples.
The pressure control module regulates the digestion cycle by monitoring, controlling, and
dwelling at several preferred pressure levels for specified time periods in order to obtain
complete dissolution and precise recoveries in the minimum amount of time. As the samples are
irradiated, temperatures in the vessels rise thereby increasing the pressure. The pressure
transducer will cycle the magnetron to maintain sufficient heat to hold the samples at the
programmed pressure level for a preset dwell time. The vessels are designed to vent safely in
case of excessive internal pressure.
13.5.3 High-Pressure, Closed-Vessel Systems
Recent advances in vessel design have produced microwave vessels capable of withstanding
pressures on the order of 1,500 psi (10 mPa; Lautenschlager, 1989), allowing for larger sample
sizes on the order of 1 to 2 g for soil (Smith and Yaeger, 1996) or 0.5 to 3 g for vegetation
(Alvarado et al., 1996) and, consequently, better detection limits. These high-pressure vessels are
used to digest organic and inorganic substances, such as coals, heavy oils, refractories, and
ceramic oxides, which cannot easily be digested with other techniques. Additionally, vessel
composition continues to improve. Noltner et al. (1990) have demonstrated that
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Sample Dissolution
Tetrafluorometoxil™ (TFM) vessels exhibit significantly lower blank background values from
residual contamination and reuse than vessels produced with the more traditional PFA. This
lower “memory” results in lower detection limits, a clear advantage for environmental
laboratories.

13.6 Verification of Total Dissolution
Following aggressive acid digestion or fusion, the analyst often must determine if the sample has
indeed been dissolved. This determination is made first through visual inspection for particulate
matter in the acid leachate, post-digestion solution, or dissolved fusion melt. (The analyst should
allow the solution to cool prior to making an assessment of total dissolution.) A hot digestate
may appear to be free from particulate matter. However, upon cooling, finely divided particulate
or colloidal matter may agglomerate, forming a residue. If a residue is observed, this residue
must be physically separated, or the sample digestate must be retreated to ensure a single final
aqueous phase. Sometimes these residues are inconsequential and contain no analyte of interest.
Project-specific requirements will dictate how these residues are handled.
If no particles are readily observed, small undissolved particles that are invisible to the unaided
eye may be present. A method to assess this may be to filter a duplicate cooled solution (see
Section 10.3.2, “Liquid Sample Preparation: Filtration”) and count it using a gamma spectrometer,
alpha spectrometer, or proportional counter. The analyst should focus on the analytes of interest
to assess whether any activity is lost in this residue. Finally, for those cases where the laboratory
has decided to perform an acid leaching, rather than a total dissolution or fusion, it is advisable to
perform total dissolution on a subset of the samples and compare the results to those obtained
from the acid digestion. This check will help to substantiate that the acid leaching approach is
adequate for the particular sample matrix.

13.7 Special Matrix Considerations
13.7.1 Liquid Samples
Aqueous samples usually are considered to be in solution. This may not always be true, and,
based on the objectives of the project, additional decomposition of aqueous samples may be
requested.
Most radiochemical analyses are performed in aqueous solutions. Because nonaqueous liquids
are incompatible with this requirement, these samples must be converted into an aqueous form.
In most cases, the nonaqueous liquid is simply a solvent that does not contain the radionuclide of
interest, and the nonaqueous solvent simply can be removed and the residue dissolved as
described in Sections 13.3 (“Fusion Techniques”) and 13.4 (“Wet Ashing and Acid Dissolution
Techniques”).
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Occasionally, the nonaqueous phase must be analyzed. A procedure for the decomposition of
petroleum products is described by Coomber (1975). There are restrictions on how many
nonaqueous liquids can be disposed of, even as laboratory samples. Evaporation of volatile
solvents may initially be an attractive alternative, but the legal restrictions on evaporating
solvents into the air should be investigated before this method is implemented. Burning flammable liquids such as oil may also initially appear attractive, but legal restrictions on incineration
of organic liquids need to be considered. A liquid-liquid extraction or separation using ion
exchange resin may be the only alternative for transferring the radionuclide of interest into an
aqueous solution. Unfortunately, these methods require extensive knowledge of the sample
matrix and chemical form of the contaminant, which is seldom available. Often, gross
radioactivity measurements using liquid scintillation counting techniques or broad spectrum
direct measurements such as gamma spectroscopy are the only measurements that can be
practically performed on nonaqueous liquids.
13.7.2 Solid Samples
Decomposition of solid samples is accomplished by applying fusion, wet ashing, leaching, or
combustion techniques singly or in some combination. A discussion of each of these techniques
is included in this chapter.
13.7.3 Filters
Air filter samples generally have a small amount of fine particulate material on a relatively small
amount of filter media. In many cases, filters of liquid samples also have limited amounts of
sample associated with the filter material. This situation may initially appear to make the sample
decomposition process much easier, the small amount of sample appears to dissolve readily in a
simple acid dissolution. The ease with which many filters dissolve in concentrated acid does not
always mean that the sample has dissolved, and the fine particles are often impossible to observe
in an acid solution. If the radionuclides of concern are known to be in the oxide form, or if the
chemical form of the contaminants is unknown, a simple acid dissolution will not completely
dissolve the sample. In these cases, the sample may be dry ashed to destroy the filter and the
residue subjected to fusion or other decomposition of oxides in the sample.
13.7.4 Wipe Samples
If oxides and silicates are not present in wipe samples, acid dissolutions are generally acceptable
for sample decomposition. In many cases, it is not the sample but the material from which the
wipe is constructed that causes problems with acid dissolution. Paper wipes are decomposed
easily in sulfuric-nitric solutions or in perchloric nitric solutions or by combustion, and it may be
necessary to dry ash the sample before dissolution. If volatile isotopes are expected, precautions
must be taken to prevent loss when heating (see Section 14.5, “Volatilization and Distillation).
“Sticky” smears can be more difficult to dissolve—the glue can be especially troublesome and
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should be watched closely if perchloric acid is used. Other materials used for wipe samples
should be evaluated on an individual basis to determine the best method for sample decomposition. In some cases, the sample will be a problem to decompose as well. Oil and grease are often
collected on wipe samples from machinery, and these samples are usually dry ashed before acid
dissolution to remove the organic material. If large amounts of solid material (i.e., soil, dust, etc.)
are collected with the wipe, it is recommended that the sample be treated as a solid (the analytical
protocol specification or the project manager should be consulted before removing the wipe and
simply analyzing the solid sample).

13.8 Comparison of Total Dissolution and Acid Leaching
Sample dissolution can be one of the biggest challenges facing the analyst because the adequacy
of the dissolution has direct and profound effects on the resultant data. The analyst must balance
numerous factors such as the nature of the sample and the analyte (e.g., is it refractory or
volatile?), the effects of excess reagents during subsequent analyses, the accuracy and precision
requirements for the data, and the costs associated with effort, materials, and waste generation.
Consequently, the question of total dissolution through fusion or digestion, or through acid
leaching, is under constant debate, and it is important for the analyst to be aware of the
limitations of both methods.
The MARLAP process enables one to make a decision concerning the dissolution required
through its process of establishing data quality objectives, analytical protocol specification, and
measurement quality objectives. During this process, all pertinent information is available to the
radioanalytical specialist who then evaluates the alternatives and assists with the decision. The
following discussion on acid leaching focuses on its use for the complete dissolution of the
analyte of interest and not for such procedures as the Environmental Protection Agency’s
“Toxicity Characteristic Leaching Procedure” (TCLP; 40 CFR 261, Appendix II, Method 1311),
which are intended to determine the leachability of a nonradioactive analyte.
“Acid leaching” has no accepted definition, but will be defined here as the use of nitric or
hydrochloric acid to put the radionuclide into solution. The acid concentration may vary up to
and include concentrated acid. Normally, the use of hydrofluoric acid and aqua regia are not
included in this definition. Sample size is usually relatively much larger than that used for fusion.
Although mineral acids might not totally break down all matrices, they have been shown to be
effective leaching solvents for metals, oxides, and salts in some samples. In some cases, leaching
requires fewer chemicals and less time to accomplish than complete sample dissolution. For
matrices amenable to leaching, multiple samples are easily processed simultaneously using a
hotplate or microwave system, and excess reagents can be removed through evaporation.
Complete dissolution of a sample is not necessary if it can be demonstrated confidently that the
radionuclide of interest is completely leached from the sample medium. However, as indicated
by Sill and Sill (1995), this may not always be possible:
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“In many cases, the mono-, di-, and small tervalent elements can be leached fairly
completely from simple solids by boiling with concentrated hydrochloric or nitric acids.
However, even these elements cannot necessarily be guaranteed to be dissolved completely by selective leaching. If they are included in a refractory matrix, they will not be
removed completely without dissolution of the matrix. If the samples have been exposed
to water over long periods of time, such as with sediments in a radioactive waste pond,
small ions such as divalent cobalt will have diffused deeply into the rock lattice from
which they cannot be removed without complete dissolution of the host matrix. In
contrast, because of its large size, ionic cesium has a marked tendency to undergo isomorphous replacement in the lattice of complex silicates from which it too cannot be
removed completely.”
Thus, the results of acid leaching processes should be used with caution.
There are those within the radiochemistry community who contend that total sample dissolution
provides the most analytically accurate and reproducible analyte concentration in the sample. Sill
and Sill (1995), longtime proponents of total dissolution, state:
“Any procedure that fails to obtain complete sample dissolution …will inevitably give
low and erratic results. The large ter-, quadri-, and pentavalent elements are extremely
hydrolytic and form hydroxides, phosphates, silicates, carbides, etc., that are very
insoluble and difficult to dissolve in common acids, particularly if they have been heated
strongly and converted to refractory forms.”
However, there are also disadvantages and challenges associated with the fusion approach.
Fusions are frequently more labor intensive than the leaching approach. More often than not,
single-sample processing requires a dedicated analyst. Large quantities of the flux are generally
required to decompose most substances, often 5 to10 times the sample weight. Therefore,
contamination of the sample by impurities in the reagent is quite possible. Furthermore, the
aqueous solutions resulting from the fusions will have a very high salt content, which may lead to
difficulties in subsequent steps of the analysis, i.e., difficulties of entrainment, partial replacements, etc. The high temperatures associated with some fusion processes increase the danger of
loss of certain analytes by volatilization. Finally, the crucible itself may be attacked by the flux,
once again leading to possible contamination of the sample. The typical sample size for fusions
ranges from typically one to ten grams. The analyst must consider whether this sample is
representative.

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13.9 References
13.9.1 Cited References
Abu-Samra, A., Morris, J.S., and Koirtyohann, S.R. 1975. “Wet Ashing of Some Biological
Samples in a Microwave Oven,” Analytical Chemistry, 47:8, pp 1475-1477.
Alvarado, J.S., Neal, T.J., Smith, L.L., and Erickson, M.D. 1996. “Microwave Dissolution of
Plant Tissue and the Subsequent Determination of Trace Lanthanide and Actinide Elements
by Inductively Coupled Plasma-Mass Spectrometry,” Analytica Chimica Acta, Vol. 322, pp.
11-20.
American Society for Testing Materials (ASTM) D4309. “Standard Practice for Sample
Digestion Using Closed Vessel Microwave Heating Technique for the Determination of Total
Metals in Water,” in 1994 Annual Book of ASTM Standards, Vol. 11.01, 1996.
American Society for Testing Materials (ASTM) D5258. “Standard Practice for Acid-Extraction
from Sediments Using Closed Vessel Microwave Heating,” in 1992 Annual Book of ASTM
Standards, Vol. 11.02, 1992.
American Society for Testing Materials (ASTM) D5513. “Standard Practice for Microwave
Digestion of Industrial Furnace Feedstreams for Trace Element Analysis,” in 1994 Annual
Book of ASTM Standards, Vol. 11.04, 1994.
Bock, R. 1979. A Handbook of Decomposition Methods in Analytical Chemistry, Halsted Press,
John Wiley and Sons, New York.
Bogen, DC. 1978. “Decomposition and Dissolution of Samples: Inorganic,” in Kolthoff, I.M. and
Elving, P.J., Eds., Treatise on Analytical Chemistry, Part I, Vol. 5, Wiley-Interscience, New
York, pp. 1-22.
Booman, G.L. and Rein, J.E. 1962. “Uranium,” in Kolthoff, I.M. and Elving, P.J., Eds., Treatise
on Analytical Chemistry, Part, Volume 9, John Wiley and Sons, New York, pp. 1-188.
Burnett, W.C., Corbett, D.R., Schultz, M., and Fern, M. 1997. “Analysis of Actinide Elements in
Soils and Sediments,” presented at the 44th Bioassay Analytical and Environmental
Radioactivity (BAER) Conference, Charleston.
Cobble, J.W. 1964. “Technetium,” in Kolthoff, I.M. and Elving, P.J., Eds., Treatise on
Analytical Chemistry, Part II, Volume 6, John Wiley and Sons, New York, pp. 404-434.

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Coomber, D.I. 1975. “Separation Methods for Inorganic Species,” in Radiochemical Methods in
Analysis, Coomber, D.I., Ed., Plenum Press, New York, pp. 175-218.
Dean, J. 1995. Analytical Chemistry Handbook, McGraw-Hill, New York.
U.S. Environmental Protection Agency (EPA). 1996. “Microwave Assisted Digestion of
Siliceous and Organically Based Materials,” in Test Methods for Evaluating Solid Waste,
Physical/Chemical Methods, SW-846, Method 3052. December.
Gibbs, J., Everett, L., and Moore, D. 1978. Sample Preparation for Liquid Scintillation
Counting, Packard Instrument Co., Downers Grove, IL., pp 65-78.
Gilman, L.B., and Engelhardt, W.G. 1989. “Recent Advances in Microwave Sample
Preparation,” Spectroscopy, 4:8, pp. 4-21.
Grillo, A.C. 1989. “Microwave Digestion by Means of a Focused Open-Vessel System,”
Spectroscopy, 4:7, pp. 16-21.
Grillo, A.C. 1990. “Microwave Digestion Using a Closed Vessel System,” Spectroscopy, 5:1, pp.
14, 16, 55.
Grindler, J.E. 1962. The Radiochemistry of Uranium, National Academy of Sciences-National
Research Council (NAS-NS), NAS-NS 3050, Washington, DC.
Hahn, R.B. 1961. “Zirconium and Hafnium,” in Kolthoff, I.M. and Elving, P.J., Eds., Treatise on
Analytical Chemistry, Part II, Volume 5, John Wiley and Sons, New York, pp. 61-138.
Kammin, W.R., and Brandt, M.J. 1989. “The Simulation of EPA Method 3050 Using a HighTemperature and High-Pressure Microwave Bomb,” Spectroscopy, 4:6, pp. 22, 24.
Kingston, H.M., and Jassie, L.B. 1988. Introduction to Microwave Sample Preparation: Theory
and Practice, American Chemical Society, Washington, DC.
Kingston, H.M., and S.J. Haswell. 1997. Microwave-Enhanced Chemistry: Fundamentals,
Sample Preparation, and Applications, American Chemical Society, Washington, DC.
Lautenschlager, W. 1989. “Microwave Digestion in a Closed-Vessel, High-Pressure System,”
Spectroscopy, 4:9, pp. 16-21.
Noltner, T., Maisenbacher, P., and Puchelt, H. 1990. “Microwave Acid Digestion of Geological
and Biological Standard Reference Materials for Trace Element Analysis by ICP-MS,”
Spectroscopy, 5:4, pp. 49-53.
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Sample Dissolution
Peng, T. 1977. Sample Preparation in Liquid Scintillation Counting, Amersham Corporation,
Arlington Heights, IL., pp. 48-54.
Schilt, A. 1979. Perchloric Acids and Perchlorates, The G. Frederick Smith Company,
Columbus, Ohio.
Sill, C.W., Puphal, K.W., and Hindman, F.D. 1974. “Simultaneous Determination of AlphaEmitting Nuclides from Radium through Californium in Soil,” Analytical Chemistry, 46:12,
pp. 1725-1737.
Sill, C.W. 1975. “Some Problems in Measuring Plutonium in the Environment,” Health Physics,
Vol. 29, pp. 619-626.
Sill, C.W. 1981. “A Critique of Current Practices in the Determination of Actinides,” in
Actinides in Man and Animal, Wren, M.E., Ed., RD Press, Salt Lake City, Utah, pp. 1-28.
Sill, C.W. and Sill, D.S. 1995. “Sample Dissolution,” Radioactivity and Radiochemistry, 6:2, pp.
8-14.
Smith, LL., Crain, J.S., Yaeger, J.S., Horwitz, E.P., Diamond, H., and Chiarizia, R. 1995.
“Improved Separation Method for Determining Actinides in Soil Samples,” Journal of
Radioanaytical Nuclear Chemistry, Articles, 194:1, pp. 151-156.
Smith, L.L. and Yaeger, J.S. 1996. “High-Pressure Microwave Digestion: A Waste-Minimization
Tool for the Radiochemistry Laboratory,” Radioactivity and Radiochemistry, 7:2, pp. 35-38.
Steinberg, E.O. 1960. The Radiochemistry of Zirconium and Hafnium, National Academy of
Sciences-National Research Council (NAS-NRC), NAS-NRC 3011, Washington, DC.
Sulcek, Z., and Povondra, P. 1989. Methods for Decomposition in Inorganic Analysis, CRC
Press, Inc., Boca Raton, Florida.
13.9.2 Other Sources
Bishop, C.T., Sheehan, W.E., Gillette, R.K., and Robinson, B. 1971. “Comparison of a Leaching
Method and a Fusion Method for the Determination of Plutonium-238 in Soil,” Proceedings
of Environmental Symposium, Los Alamos Scientific Laboratory, Los Alamos, NM, U.S.
Atomic Energy Commission, Document LA-4756, December, pp. 63-71.
Burnett, W.C., Corbett, D.R. Schultz, M., Horwitz, E.P., Chiarizia, R., Dietz, M., Thakkar, A.,
and Fern, M. 1997. “Preconcentration of Actinide Elements from Soils and Large Volume
Water Samples Using Extraction Chromatography,” J. Radioanalytical and Nuclear
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Sample Dissolution
Chemistry, 226, pp.121-127.
U.S. Department of Energy (DOE). 1990. EML Procedures Manual, Chieco, N.A., Bogen, DC.,
and Knutson, E.O., Eds., HASL-300, 27th Edition, DOE Environmental Measurements
Laboratory, New York.
U.S. Environmental Protection Agency (EPA). 1992. Guidance for Preforming Site Inspections
Under CERCLA, EPA/540-R-92-021, Office of Solid Waste and Emergency Response,
Washington, DC.
MARSSIM. 2000. Multi-Agency Radiation Survey and Site Investigation Manual, Revision 1.
NUREG-1575 Rev 1, EPA 402-R-97-016 Rev1, DOE/EH-0624 Rev1. August. Available
from www.epa.gov/radiation/marssim/.
Grimaldi, F.S. 1961. “Thorium,” in Treatise on Analytical Chemistry, Kolthoff, I.M. and Elving,
P.J., Eds., Part II, Volume 5, John Wiley and Sons, New York, pp. 142-216.
Kim, G., Burnett, W.C., and Horwitz, E.P. 2000. “Efficient Preconcentration and Separation of
Actinide Elements from Large Soil and Sediment Samples,” Analytical Chemistry, 72, pp.
4882-4887.
Krey, P.W. and Bogen, DC. 1987. “Determination of Acid Leachable and Total Plutonium in
Large Soil Samples,” Journal of Radioanalytical and Nuclear Chemistry, 115:2, pp. 335-355.
Maxwell, S. and Nichols, S.T. 2000. “Actinide Recovery Method for Large Soil Samples,”
Radioactivity and Radiochemistry, 11:4, pp. 46-54.
Noyes, A.A. and Bray, W.C. 1927, reprinted 1943. A System of Qualitative Analysis for the
Rarer Elements, MacMillan, New York.
Sill, C.W. 1975. “Some Problems in Measuring Plutonium in the Environment,” Health Physics,
29, pp. 619-626.
Sill, D.S. and Bohrer, S.E. 2000. “Sequential Determination of U, Pu, Am, Th and Np in Fecal
and Urine Samples with Total Sample Dissolution,” Radioactivity and Radiochemistry, 11:3,
pp. 7-18.
Smith, L.L., Markun, F., and TenKate, T. 1992. “Comparison of Acid Leachate and Fusion
Methods to Determine Plutonium and Americium in Environmental Samples,” Argonne
National Laboratory, ANL/ACL-92/2.

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14 SEPARATION TECHNIQUES
14.1 Introduction
The methods for separating, collecting, and detecting radionuclides are similar to ordinary
analytical procedures and employ many of the chemical and physical principles that apply to their
nonradioactive isotopes. However, some important aspects of the behavior of radionuclides are
significantly different, resulting in challenges to the radiochemist to find a means for isolation of
a pure sample for analysis (Friedlander et al., 1981).
While separation techniques and principles may be found in standard textbooks, Chapter 14
addresses the basic chemical principles that apply to the analysis of radionuclides, with an
emphasis on their unique behavior. It is not a comprehensive review of all techniques. This
chapter provides: (1) a review of the important chemical principles underlying radiochemical
separations, (2) a survey of the important separation methods used in radiochemistry with a
discussion of their advantages and disadvantages, and (3) an examination of the particular
features of radioanalytical chemistry that distinguish it from ordinary analytical chemistry.
Extensive examples have been provided throughout the chapter to illustrate various principles,
practices, and procedures in radiochemistry. Many were selected purposely as familiar
illustrations from agency procedural manuals. Others were taken from the classical and recent
radiochemical literature to provide a broad, general overview of the subject.
This chapter integrates the concepts of classical chemistry with those topics unique to radionuclide analysis. The first eight sections of the chapter describe the bases for chemical
separations involving oxidation-reduction, complex-ion formation, distillation/volatilization,
solvent extraction, precipitation and coprecipitation, electrochemistry, and chromatography.
Carriers and tracers, which are unique to radiochemistry, are described in Section 14.9 together
with specific separation examples for each of the elements covered in this manual. Section 14.10
also provides an overview of the solution chemistry
Contents
and appropriate separation techniques for 17
elements. An attachment at the end of the chapter
14.1 Introduction . . . . . . . . . . . . . . . . . . . . 14-1
14.2 Oxidation-Reduction Processes . . . . . 14-2
describes the phenomenon of radioactive
14.3 Complexation . . . . . . . . . . . . . . . . . . 14-18
equilibrium, also unique to radioactive materials.
Because the radiochemist detects atoms by their
radiation, the success or failure of a radiochemical
procedure often depends on the ability to separate
extremely small quantities of radionuclides (e.g.,
10!6 to 10!12 g) that might interfere with detection
of the analyte. For example, isolation of trace
quantities of a radionuclide that will not precipitate
on its own with a counter-ion requires judicious
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14.4 Solvent Extraction . . . . . . . . . . . . . . 14-25
14.5 Volatilization and Distillation . . . . . 14-36
14.6 Electrodeposition . . . . . . . . . . . . . . . 14-41
14.7 Chromatography . . . . . . . . . . . . . . . 14-44
14.8 Precipitation and Coprecipitation . . 14-56
14.9 Carriers and Tracers . . . . . . . . . . . . 14-82
14.10 Analysis of Specific Radionuclides . 14-97
14.11 References . . . . . . . . . . . . . . . . . . . 14-201
14.12 Selected Bibliography . . . . . . . . . . 14-218
Attachment 14A Radioactive Decay and
Equilibrium . . . . . . . . . . . . . . . . . . 14-223
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selection of a carrier and careful technique to produce a coprecipitate containing the pure
radionuclide, free of interfering ions.
In detection procedures, the differences in the behavior of radionuclides provide unique opportunities not available in the traditional analytical chemistry of nonradioactive elements. Radionuclides often can be detected by their unique radiation regardless of the chemical form of the
element. There is also a time factor involved because of the short half-lives of some radionuclides. Traditional procedures involving long digestion or slow filtration cannot be used for shortlived radionuclides, thereby requiring that rapid separations be developed. Another distinction is
the hazards associated with radioactive materials. At very high activity levels, chemical effects of
the radiation, such as decomposition of solvents (through radiolysis) and heat effects (caused by
interaction of decay particles with the solution), can affect the procedures. Equally important,
even at lower activity levels, is the radiation dose that the radiochemist can receive unless
protected by shielding, ventilation, time, or distance. Even at levels where the health concerns are
minimal, special care needs to be taken to guard against laboratory and equipment contamination.
Moreover, the radiochemist should be concerned about the type and quantity of the waste
generated by the chemical procedures employed, because the costs and difficulties associated
with the disposal of low-level and mixed radioactive waste continue to rise (see Chapter 17,
Waste Management in a Radioanalytical Laboratory).
The past 10 years have seen significant improvements to some of the classical techniques as well
as the development of new methods of radiochemical analysis. Knowledge of these analytical
developments, as well as maintenance of a working familiarity with developing techniques in the
radiochemistry field will further enhance the waste reduction effort.

14.2 Oxidation-Reduction Processes
14.2.1 Introduction
Oxidation and reduction (redox) processes play an important role in radioanalytical chemistry,
particularly from the standpoint of the dissolution, separation, and detection of analytes, tracers,
and carriers. Ion exchange, solvent extraction, and solid-phase extraction separation techniques,
for example, are highly dependent upon the oxidation state of the analytes. Moreover, most
radiochemical procedures involve the addition of a carrier or isotope tracer. There must be
complete equilibration (isotopic exchange) between the added isotope(s) and all the analyte
species present in order to achieve quantitative yields. The oxidation number of a radionuclide
can affect its chemical stability in the presence of water, oxygen, and other natural substances in
solution; reactivity with reagents used in the radioanalytical procedure; solubility in the presence
of other ions and molecules; and behavior in the presence of carriers and tracers. The oxidation
numbers of radionuclides in solution and their susceptibility to change, because of natural or
induced redox processes, are critical, therefore, to the physical and chemical behavior of
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radionuclides during these analytical procedures. The differences in mass number of all
radionuclides of an element are so small that they will exhibit the same chemical behavior during
radiochemical analysis (i.e., no mass isotope effects).
14.2.2 Oxidation-Reduction Reactions
An oxidation-reduction reaction (redox reaction) is a reaction in which electrons are redistributed
among the atoms, molecules, or ions in solution. In some redox reactions, electrons are actually
transferred from one reacting species to another. Oxidation under these conditions is defined as
the loss of electron(s) by an atom or other chemical species, whereas reduction is the gain of
electron(s). Two examples will illustrate this type of redox reaction:
U + 3 F2 6 U+6 + 6 F!1
Pu+4 + Fe+2 6 Pu+3 + Fe+3
In the first reaction, uranium loses electrons, becoming a cation (oxidized), and fluorine gains an
electron (reduced), becoming an anion. In the second reaction, the reactants are already ions, but
the plutonium cation (Pu+4) gains an electron, becoming Pu+3 (reduced), and the ferrous ion (Fe+2)
loses an electron, becoming Fe+3 (oxidized).
In other redox reactions, electrons are not completely transferred from one reacting species to
another; the electron density of one atom decreases while it increases at another atom. The
change in electron density occurs as covalent bonds (in which electrons are shared between two
atoms) are broken or made during a chemical reaction. In covalent bonds between two atoms of
different elements, one atom is more electronegative than the other atom. Electronegativity is the
ability of an atom to attract electrons in a covalent bond. One atom, therefore, attracts the shared
pair of electrons more effectively, causing a difference in electron density about the atoms in the
bond. An atom that ends up bonded to a more electronegative atom at the end of a chemical
reaction loses net electron density. Conversely, an atom that ends up bonded to a less electronegative atom gains net electron density. Electrons are not transferred completely to other atoms,
and ions are not formed because the electrons are still shared between the atoms in the covalent
bond. Oxidation, in this case, is defined as the loss of electron density, and reduction is defined
as the gain of electron density. When carbon is oxidized to carbon dioxide by oxygen:
C + O2 6 CO2
the electron density associated with the carbon atom decreases, and that of the oxygen atoms
increases, because the electronegativity of oxygen is greater than the electronegativity of carbon.
In this example, carbon is oxidized and oxygen is reduced. Another example from the chemistry
of the preparation of gaseous uranium hexafluoride (UF6) illustrates this type of redox reaction:
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3 UF4 + 2 ClF3 6 3 UF6 + Cl2
Because the order of electronegativity of the atoms increases in the order U < Cl < F, the uranium
atom in uranium tetrafluoride (UF4) is oxidized further as more electronegative fluorine atoms
are added to the metal and shift the electron density away from uranium. Chlorine atoms break
their bonds with fluorine and gain electron density (are reduced) when they bond with each other
instead of the more electronegative fluorine atoms.
In a redox reaction, at least one species is oxidized and at least one species is reduced simultaneously; one process cannot occur without the other. The oxidizing agent is defined as the
substance that causes oxidation of another species by accepting electron(s) from it or increasing
in electron density; it is thereby reduced itself. Reducing agents lose electron(s) or electron
density and are therefore oxidized. In the reduction of Pu+4 to Pu+3 by Fe+2, the reducing agent
donates an electron to Pu+4 and is itself oxidized, while Pu+4, the oxidizing agent, accepts an
electron from Fe+2 and is reduced. Generally, the nonmetallic elements are strong oxidizing
reagents, and the metals are strong reducing agents.
To keep track of electrons in oxidation-reduction reactions, it is useful to assign oxidation
numbers to atoms undergoing the changes. Oxidation numbers (oxidation states) are a relative
indication of the electron density associated with an atom of an element. The numbers change
during redox reactions, whether they occur by actual transfer of electrons or by unequal sharing
of electrons in a covalent bond. The number increases as the electron density decreases, and it
decreases as the electron density increases. From the standpoint of oxidation numbers and in
more general terms, oxidation is defined as an increase in oxidation number, and reduction is
defined as the decrease in oxidation number. Different sets of rules have been developed to
assign oxidation numbers to monatomic ions and to each individual atom in polyatomic
molecules. One set of rules is simple and especially easy to use. It can be used to determine the
oxidation number of atoms in many, but not all, chemical species. In this set, the rules for
assigning oxidation numbers are listed in order by priority of application; the rule written first in
the list has priority over the rule below it. The rules are applied in the order in which they come
in the list, starting at the top and proceeding down the list of rules until each atom of each
element, not the element only, in a species has been assigned an oxidation number. Generally, all
atoms of each element in a chemical species will have the same oxidation number in that species.
For example, all oxygen in sulfate are !2. (A specific exception is nitrogen in the cation and
anion in ammonium nitrate, NH4NO3.) It is important to remember that in many cases, oxidation
numbers are not actual electrical charges but only a helpful bookkeeping method for following
redox reactions or examining various oxidation states. The oxidation number of atoms in isolated
elements and monatomic ions are actually the charge on the chemical species. The priority rules
are:
1. The sum of oxidation numbers of all atoms in a chemical species adds up to equal the
charge on the species. This is zero for elements and compounds because they are
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electrically neutral species and are the total charge for a monatomic or polyatomic ion.
2. The alkali metals (the Group IA elements, Li, Na, K, Rb, Cs, and Fr) have an oxidation
number of +1; the alkaline earth metals (the Group IIA elements, Be, Mg, Ca, Sr, Ba, and
Ra) have an oxidation number of +2.
3. Fluorine has an oxidation number of !1; hydrogen has an oxidation number of +1.
4. Oxygen has an oxidation number of !2.
5. The halogens (the Group VIIA elements, F, Cl, Br, I, and At) have an oxidation number
of !1.
6. In binary compounds (compounds containing elements), the oxidation number of the
oxygen family of elements (the Group VIA elements, O, S, Se, Te, and Po) is !2; for the
nitrogen family of elements (the VA elements except N, P, As, and Sb), it is -3.
Applying these rules illustrates their use:
1. The oxidation number of metallic uranium and molecular oxygen is 0. Applying rule one,
the charge on elements is 0.
2. The oxidation number of Pu+4 is +4. Applying rule one again, the charge is +4.
3. The oxidation numbers of carbon and oxygen in CO2 are +4 and !2, respectively.
Applying rule one, the oxidation numbers of each atom must add up to the charge of 0
because the net charge on the molecule is zero. The next rule that applies is rule four.
Therefore, the oxidation number of each oxygen atom is !2. The oxidation number of
carbon is determined by C + 2(!2) = 0, or +4. Notice that there is no charge on carbon
and oxygen in carbon dioxide because the compound is molecular and does not consist of
ions.
4. The oxidation numbers of calcium and hydrogen in calcium hydride (CaH2) are +2 and
!1, respectively. The compound is neutral, and the application of rule one requires that
the oxidation numbers of all atoms add up to 0. By rule two, the oxidation number of
calcium is +2. Applying rule one, the oxidation number of hydrogen is: 2H + 2=0, or !1.
Notice that in this example, the oxidation number as predicted by the rules does not agree
with rule three, but the number is determined by rules one and two, which take
precedence over rule three.
5. The oxidation numbers of uranium and oxygen in the uranyl ion, UO2+2, are +6 and !2,
respectively. Applying rule one, the oxidation numbers of each atom must add up to the
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charge of +2. Rule four indicates that the oxygen atoms are !2 each. Applying rule one,
the oxidation number of uranium is U + 2(!2) = +2, and uranium is +6. In this example,
the charges on uranium and oxygen are not actually +6 and !2, respectively, because the
polyatomic ion is held together through covalent bonds. The charge on the ion is the
result of a deficiency of two electrons.
Oxidation numbers (states) are commonly represented by zero and positive and negative
numbers, such as +4, !2, etc. They are sometimes represented by Roman numerals for metals,
especially the oxidation numbers of atoms participating in covalent bonds or those of polyatomic
ions, such as chromium(VI) in CrO4!2. In general, elements in solution whose oxidation number
is greater than +4 or less than -4 can exist only as complexed ions in solution. Many of the
transuranic elements can occur in multiple oxidation states, and the transformation from one to
another is a critical step of the separation process. In this chapter, all species whose oxidation
number is greater than +4 will be represented either by their complexed form in solution or by its
symbol with a Roman numeral signifying the oxidation state [UO2+ or U(V)]. This conforms to
the intent of IUPAC (1990) nomenclature.
14.2.3 Common Oxidation States
The oxidation state for any element in its free state (when not combined with any other element,
as in Cl2 or Ag metal) is zero. The oxidation state of a monatomic ion is equal to the electrical
charge of that ion. The Group IA elements form ions with a single positive charge (Li +1, Na+1,
K+1, Rb+1, and Cs+1), whereas the Group IIA elements form +2 ions (Be+2, Mg+2, Sr+2, Ba+2, and
Ra+2). The halogens generally form !1 ions (F!1, Br!1, Cl!1, and I!1); however, except for fluorine,
the other halogens form oxygen compounds in which several other oxidation states are present
[Cl(I) in HClO and I(V) in HIO3]. For example, iodine can exist as I !1, I2, IO!1, IO3!1, and IO4!1.
Oxygen exhibits a !2 oxidation state except when it is bonded to fluorine (where it can be +1 or
+2); in peroxides, where the oxidation state is !1; or in superoxides, where it is -½.
Some radionuclides, such as those of cesium and thorium, exist in solution in single oxidation
states, as indicated by their position in the periodic table. Others, such as technetium and
uranium, can exist in multiple oxidation states. Multiple oxidation states of plutonium are
commonly found in the same solution.
Each of the transition metals has at least two stable oxidation states, except for Sc, Y, and La
(Group IIIB), which exhibit only the +3 oxidation state. Generally, negative oxidation states are
not observed for these metallic elements. The large number of oxidation states exhibited by the
transition elements leads to an extensive, often complicated, oxidation-reduction chemistry. For
example, oxidation states from !1 through +7 have been observed for technetium, although the
+7 and +4 are most common (Anders, 1960). In an oxidizing environment, Tc exists predominantly in the heptavalent state as the pertechnetate ion, TcO4!1, which is water soluble, but which
can yield insoluble salts with large cations. Technetium forms volatile heptoxides and acidMARLAP

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insoluble heptasulfides. Subsequently, pertechnetate is easily lost upon evaporation of acid
solutions unless a reducing agent is present or the evaporation is conducted at low temperatures.
Technetium(VII) can be reduced to lower oxidation states by reducing agents such as bisulfite
(HSO3!1). This process proceeds through several intermediate steps, some of which are slow;
therefore, unless precautions are taken to maintain technetium in the appropriate oxidation state,
erratic results can occur. The (VII) and +4 ions behave very differently in solution. For instance,
pertechnetate does not coprecipitate with ferric hydroxide, while Tc+4 does.
The oxidation states of the actinide elements have been comprehensively discussed by Ahrland
(1986) and Cotton and Wilkinson (1988). The actinides exhibit an unusually broad range of
oxidation states, of from +2 to +7 in solution. Similar to the lanthanides, the most common
oxidation state is +3 for actinium, americium, and curium. The +4 state is common for thorium
and plutonium, whereas (V) is most common for protactinium and neptunium. The most stable
state for uranium is the (VI) oxidation state.
In compounds of the +3 and +4 oxidation states, the elements are present as simple M+3 or M+4
cations (where “M” is the metal ion); but for higher oxidation states, the most common forms in
compounds and in solution are the oxygenated actinyl ions, MO2+1 and MO2+2:
• M+3. The +3 oxidation state is the most stable condition for actinium, americium, and curium,
and it is easy to produce Pu+3. This stability is of critical importance to the radiochemistry of
plutonium. Many separation schemes take advantage of the fact that Pu can be selectively
maintained in either the +3 or +4 oxidation state. Unlike Pu and Np, U+3 is such a strong
reducing agent that it is difficult to keep in solution.
• M+4. The only oxidation state of thorium that is experienced in radiochemical separations is
+4. Pa+4, U+4, and Np+4 are stable, but they are easily oxidized by O2. In acid solutions with
low plutonium concentrations, Pu+4 is stable. Americium and curium can be oxidized to the
+4 state with strong oxidizing agents such as persulfate.
• M(V). The actinides, from protactinium through americium, form MO2+1 ions in solution.
PuO2+1 can be the dominant species in solution at low concentration in natural waters that are
relatively free of organic material.
• M(VI). This is the most stable oxidation state of uranium, which exist as the UO2+2 species.
Neptunium, plutonium, and americium also form MO2+2 ions in solution. The bond strength,
as well as the chemical stability toward reduction for these MO2+2 ions, decrease in the order
U > Np > Pu > Am.
Reactions that do not involve making or breaking bonds, M+3 6 M+4 or MO2+1 6 MO2+2, are fast
and reversible, while reactions that involve chemical bond formation, M+3 6 MO2+1 or
M+4 6 MO2+2, are slow and irreversible.
JULY 2004

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Plutonium exhibits redox behavior unmatched in the periodic table. It is possible to prepare
solutions of plutonium ions with appreciable concentrations of four oxidation states, +3, +4, (V),
and (VI), as Pu+3, Pu+4, PuO2+1, and PuO2+2, respectively. Detailed discussions can be found in
Cleveland (1970), Seaborg and Loveland (1990), and in Coleman (1965). According to
Cleveland (1970), this polyvalent behavior occurs because of the tendency of Pu+4 and Pu(V) to
disproportionate:
3 Pu+4 + 2 H2O 6 2 Pu+3 + PuO2+2 + 4H+1
3 PuO2+1 + 4 H+1 6 Pu+3 + 2 PuO2+2 + 2 H2O
and because of the slow rates of reaction involving formation or rupture of Pu-O bonds (such as
PuO2+ and PuO22+) compared to the much faster reactions involving only electron transfer. The
distribution depends on the type and concentration of acid used for dissolution, the method of
solution preparation, and the initial concentration of the different oxidation states. In HCl, HNO3,
and HClO4, appreciable concentrations of all four states exist in equilibrium. Seaborg and
Loveland (1990) report that in 0.5 M HCl at 25 EC, the equilibrium percentages of plutonium in
the various oxidation states are found to be as follows:
Pu+3
Pu+4
Pu(V)
Pu(VI)

27.2%
58.4%
~0.7%
13.6%

Apart from the disproportionation reactions, the oxidation state of plutonium ions in solution is
affected by its own decay radiation or external gamma and X-rays. At high levels, radiolysis
products of the solution can oxidize or reduce the plutonium, depending on the nature of the
solution and the oxidation state of plutonium. Therefore, the stated oxidation states of old
plutonium solutions, particularly old HClO4 and H2SO4 solutions, should be viewed with
suspicion. Plutonium also tends to hydrolyze and polymerize in solution, further complicating the
situation (see Section 14.10, “Analysis of Specific Radionuclides”).
Tables 14.1 and 14.2 summarize the common oxidation number(s) of some important elements
encountered in the radioanalytical chemistry of environmental samples and the common
chemical form of the oxidation state.
TABLE 14.1 — Oxidation states of elements
Element
Am

MARLAP

Oxidation
State(1)

Chemical Form

Notes(2)

+3
+4
(V)

Am+3
Am+4
AmO2+1

(VI)

AmO2+2

Pink; stable; difficult to oxidize
Pink-red; unstable in acid
Pink-yellow; disproportionates in strong acid; reduced by products of
its own radiation
Rum color; stable

14-8

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Separation Techniques
Oxidation
State(1)

Chemical Form

+1
+2
+3
+2
+3

Cs(H2O)x+1
Co(H2O)6+2
Co(H2O)6+3
Fe(H2O)6+2
Fe(H2O)6+3

H

+1

3

I

!1
-1/3
+1
(V)
(VII)
+2
+3
+5
+4
+3
+4
(V)

Element
Cs
Co
Fe

3

Ni
Nb
Po
Pu

(VI)
(VII)

Ra
Sr
Tc

Th
U

Zr

+2
+2
+4
(V)
(VII)
+4
+3
+4
(V)
(VI)
+4

HOH and
3
HOH2+1
I!1
I3!1
OI-1
IO3!1
IO4!1
Ni(H2O)6+2
Unknown
HNb6O19!7
Pu(H2O)x+3
Pu(H2O)x+4
Pu(H2O)x+5
or
PuO2+1
PuO2+2
PuO5!3
or PuO4(OH)2!3
Ra(H2O)x+2
Sr(H2O)x+2
TcO3!2
TcO3!1
TcO4!1
Th(H2O)8+4
U(H2O)x+3
U(H2O)8 or 9+4
UO2+1
UO2(H2O)5+2
Zr(H2O)6+4
Zr4(OH)8(H2O)16+2

Notes(2)
Colorless; x probably is 6
Pink to red; oxidation is very unfavorable in solution
Rapidly reduced to +2 by water unless acidic
Green
Pale yellow; hydrolyses in solution to form yellow or brown
complexes
Isotopic exchange of tritium is extremely rapid in samples that have
water introduced.
Colorless
Brown; commonly in solutions of I!1 exposed to air
Colorless
Colorless; formed in vigorously oxidized solutions
Colorless
Green
In sulfuric acid solutions of Nb2O5

Violet to blue; stable to air and water; easily oxidized to +4
Tan to brown; first state formed in freshly prepared solutions; stable
in 6 M acid; disproportionates in low acidity to +3 and +6
Never observed alone; always disproportionates; most stable in low
acidity
Purple
Yellow-pink; stable but fairly easy to reduce
Green
PuO4(OH)2!3 more likely form
Colorless; behaves chemically like Sr and Ba
Colorless

Colorless; at pH>3 forms complex hydrolysis products
Red-brown; slowly oxidized by water and rapidly by air to +4
Green; stable but slowly oxidized by air to (VI)
Unstable but more stable at pH 2-4; disproportionates to +4 and (VI)
Yellow; only form stable in solution containing air; difficult to reduce
Only at very low ion concentrations and high acidity
At typical concentrations in absence of complexing agents

(1) Most common form is in bold.
(2) Color shades may vary depending on the concentration of the isotope.
Sources: Booman and Rein, 1962; Cotton and Wilkinson, 1988; Emsley, 1989; Greenwood and Earnshaw,
1984; Grinder, 1962; Hampel, 1968; Katzin, 1986; Latimer, 1952; and 1970.

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TABLE 14.2 — Oxidation states of selected elements
Element
Titanium
Vanadium
Chromium
Manganese
Iron
Cobalt
Nickel
Strontium
Yttrium
Molybdenum
Technetium
Silver
Cesium
Barium
Lanthanides
Lead
Polonium
Radium
Actinium
Thorium
Protactinium
Uranium
Neptunium
Plutonium
Americium
Curium

+1

+2
"
"
!
!
!
!
!
!
"
"

!
!

+3
"
"
!
"
!
!
"

+4
!
!
"
!
"

!
"
"
"

V
!
"
"

VI

VII

!
"
"

!

VIII

"
!
!
"

!
"

!
"

!

!
!
!
"
!

"
!

"

!
"
"
"
!
!

!
"
"
"
!
"
"

!
"
!
"
"

!
"
"
"

"

The stable nonzero oxidation states are indicated. The more common oxidation states
are indicated by solid black circles.
Sources: Seaborg and Loveland (1990) and the NAS–NRC monographs listed in the
references.

14.2.4 Oxidation State in Solution
For the short-lived isotopes that decay by alpha emission or spontaneous fission, high levels of
radioactivity cause heating and chemical effects that can alter the nature and behavior of the ions
in solution and produce chemical reactions not observed with longer-lived isotopes. Decomposition of water by radiation (radiolysis) leads to H and OH free radicals and formation of H2 and
H2O2, among other reactive species, and higher oxidation states of plutonium and americium are
produced.
The solutions of some ions are also complicated by disproportionation, the autooxidationreduction of a chemical species in a single oxidation state to higher and lower oxidation states.
The processes are particularly dependent on the pH of the solution. Oxidation of iodine, uranium,
americium, and plutonium are all susceptible to this change in solution. The disproportionation
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Separation Techniques
of UO2+1, for example, is represented by the chemical equation:
2 UO2+1 + 4 H+1 º U+4 + UO2+2 + 2 H2O (K = 1.7×106)
The magnitude of the equilibrium constant reflects the instability of the (V) oxidation state of
uranium in UO2+1 described in Table 14.1, and the presence of hydrogen ions reveals the
influence of acidity on the redox process. An increase in acidity promotes the reaction.
14.2.5 Common Oxidizing and Reducing Agents
HYDROGEN PEROXIDE. Hydrogen peroxide (H2O2) has many practical applications in the
laboratory. It is a very strong oxidizing agent that will spontaneously oxidize many organic
substances, and water samples are frequently boiled with peroxide to destroy organic compounds
before separation procedures. When hydrogen peroxide serves as an oxidizing reagent, each
oxygen atom changes its oxidation state from !1 to !2. For example, the reaction for the
oxidation of ferrous ion is as follows:
H2O2 + 2H+1 + 2Fe+2 º 2H2O + 2Fe+3
Hydrogen peroxide is frequently employed to oxidize Tc+4 to the pertechnetate:
4 H2O2 + Tc+4 º TcO4!1 + 4H2O
Hydrogen peroxide can also serve as a reducing agent, with an increase in oxidation state from
-1 to 0, and the liberation of molecular oxygen. For example, hydrogen peroxide will reduce
permanganate ion (MnO4!1) in basic solution, forming a precipitate of manganese dioxide:
2 MnO4!1 + 3 H2O2 6 2 MnO29 + 3 O28 + 2 H2O + 2 OH!1
Furthermore, hydrogen peroxide can decompose by the reaction:
2 H2O2 6 2 H2O + O2
This reaction is another example of a disproportionation (auto-oxidation-reduction) in which a
chemical species acts simultaneously as an oxidizing and reducing agent; half of the oxygen
atoms are reduced to O!2, and the other half are oxidized to elemental oxygen (O0) in the
diatomic state, O2.
OXYANIONS. Oxyanions (NO3!1, Cr2O7!2, ClO3!1, and MnO4!1) differ greatly in their oxidizing
strength, but they do share certain characteristics. They are stronger oxidizing agents in acidic
rather than basic or neutral conditions, and they can be reduced to a variety of species depending
on the experimental conditions. For example, on reduction in acidic solutions, the permanganate
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Separation Techniques
ion accepts five electrons, forming the manganous ion Mn+2:
MnO4!1 + 5 e!1 + 8 H+1 6 Mn+2 + 4 H2O
In neutral or basic solution, permanganate accepts 3 electrons, and forms manganese dioxide
(MnO2), which precipitates:
MnO4!1 + 3 e!1 + 4 H+1 6 MnO2 9 + 2 H2O
These oxidizing agents are discussed further in Section 13.4, “Wet Ashing and Acid Dissolution
Techniques.”
NITRITE. Nitrite ion (NO2!1), plays an important role in the manipulation of Pu oxidation states in
solution. It is capable of oxidizing Pu+3 to Pu+4 and of reducing Pu(VI) to Pu+4. Because most
aqueous processes center around Pu+4, sodium nitrite (NaNO2) is frequently used as a valence
adjuster to convert all Pu to the +4 state. And because the Pu(VI) 6 Pu+4 reaction by nitrite is
slow, another reducing agent, such as the ferrous ion, often is added to increase the rate of
reaction.
PERCHLORIC ACID. The use of perchloric acid (HClO4) as an oxidizing agent is covered in depth
in Section 13.4, “Wet Ashing and Acid Dissolution Techniques.”
METALS IONS. Generally, metals ions (Ti+3, Cr+2, Fe+2, etc.) are strong reducing agents. For
example, both Ti+3 and Cr+2 have been shown to reduce Pu+4 to Pu+3 rapidly in acidic media.
Fe+2 rapidly reduces Np(V) to Np+4 and Pu+4 to Pu+3 in acidic media.
Ti+3 is used extensively as a reducing agent in both inorganic and organic analyses. Ti +3 is
obtained by reducing Ti+4, either electrolytically or with zinc. Ti+4 is the most stable and common
oxidation state of titanium. Compounds in the lower oxidation states (!1, 0, +2, and +3) are quite
readily oxidized to Ti+4 by air, water, or other reagents.
ASCORBIC ACID. Commonly known as vitamin C, ascorbic acid is an important reducing agent
for the radiochemist. Because the ferric ion interferes with the uptake of Am+3 in several popular
extraction schemes, ascorbic acid is used frequently to reduce Fe+3 to Fe+2 to remove this
interference. Ascorbic acid is also used to reduce Pu+4 to Pu+3.
SULFAMIC ACID. Aqueous solutions of this solid material are strongly acidic and act selectively
as oxidizing agents. It is of particular value in its ability to oxidize nitrites to nitrates while not
affecting Pu+3 or Np+4 ions.

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14.2.6 Oxidation State and Radiochemical Analysis
Most radiochemical analyses require the radionuclide be in aqueous solution. Thus, the first step
of an analysis is the complete dissolution of the sample, so that all components remaining at the
end of the process are in a true solution, and chemical equilibration with tracers or carriers can be
established. Dissolution of many samples requires vigorous conditions to release the radionuclides from its natural matrix. Strong mineral acids or strong bases, which also serve as powerful
oxidizing agents, are used in boiling mixtures or under fusion conditions to decompose the
matrix—evaporating portions of the acid or base from the mixture and oxidizing the radionuclide
to a common oxidation state. The final state depends, generally, on the radionuclide, oxidizers
used, and pH of the solution (see notes to Table 14.1, page 14-9). Even water samples might
contain radionuclides at various states of oxidation because of their exposure to a variety of
natural oxidizing conditions in the environment and the pH of the sample.
Once the analyte is in solution, the radionuclide and the tracers and carriers used in the procedure
must be in the same oxidation state to ensure the same chemical behavior (Section 14.10.2,
“Oxidation State”). For radionuclides that can exist in multiple oxidation states, one state must
be achieved; for those such as plutonium, which disproportionates, a reproducible equilibrium
mixture of all oxidation states can be established. Oxidizing or reducing agents are added to the
reaction mixture to establish the required conditions. Table 14.3 contains a summary of several
chemical methods for the oxidation and reduction of select radionuclides.
In some radioanalytical procedures, establishing different states at different steps in the procedure
is necessary to ensure the requisite chemical behavior of the analyte.
TABLE 14.3 — Redox reagents for radionuclides(1)
Redox Reaction

Reagent

Am+3 6 AmO2+2
Am+4 6 AmO2+2
AmO2+1 6 AmO2+2

Ag+2, Ag+/S2O8!2
O3
Ce+4
O3
Br!1, Cl!1
Na2CO3
I!1, H2O2, NO2!1, SO2
alpha radiation effects
O3
O2, H2O2

AmO2+2 6 AmO2+1
+2

+3

AmO2 6 Am
Am+4 6 Am+3
Co+2 6 Co+3
Co+3 6 Co+2
Fe+2 6 Fe+3

H2O
O2
Ce+4, MnO4!1, NO3!1,
H2O2, S2O8!2
Cr2O7!2

JULY 2004

Conditions
13 M NH4F
HClO4
Heated HNO3 or HClO4
Heat to precipitate NaAmO2CO3; dissolve in H+1
Spontaneous
Cold HClO4
Complexed cobalt
Rapid with evolution of H2
Faster in base; slower in neutral and acid solution; decreases
with H+1

HCl or H2SO4

14-13

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Separation Techniques
Redox Reaction

Reagent

Fe+3 6 Fe+2

H2S, H2SO3
Zn, Cd, Al, Ag amalgams
Sn+2, I!1, Cu+1, Ti+3
NH2OH
HNO2 (NaNO2 in acid)
MnO2 in acid
6M HNO3
NaHSO3 or NaHSO3 in H+1
Na2SO3; Na2S2O3
KMnO4
50% CrO3 in 9M H2SO4
NaClO in base
NH2OH@HCl
H2C2O4
NaHSO3 in acid
SO2; NaHSO3

I!1 6 I2

I!1 6 IO3!1
I!1 6 IO4!1
IO4!1 6 I2
IO4!1 6 I!1
I2 6 I!1
Np+3 6 Np+4
Np+4 6 NpO2+1
Np+4 6 NpO2+2
NpO2+1 6 NpO2+2
NpO2+2 6 NpO5!3
NpO2+16 Np+4
Pu+3 6 Pu+4

Pu+4 6 PuO2+2

PuO2+1 6 PuO2+2
PuO2+2 6 PuO2+1

MARLAP

NO2!1
MnO4!1
Fe+2
Ti+3
BrO3!1
Ce+4
!2
Cr2O7 , IO3!1, MnO4!1
NO2!1
NO3!1
HNO2
NaBiO3
BrO3!1
Ce+4
HOCl (KClO)
MnO4!1
O3
Ag+2
Cr2O7!2
Cl2
NO3!1
Ag2O
IO3!
HNO3
+3
V or Ti+3
I!1
SO2

Conditions
Excess removed by boiling

Boiling solution
Does not affect other halides
Well suited for lab work

9 M H2SO4

Dilute acid
HNO3
Dilute alkaline
Acid
Acid
Dilute H2SO4
1–2 M HCl
Dilute H+1
HCl or H2SO4 solution
Dilute H+1
HNO3
HNO3 or dilute HCl (100EC)
HNO3
Dilute HNO3 at 85EC
Dilute HNO3 or HClO4
pH 4.5 at 80EC or 45% K2CO3 at 40EC
Dilute HNO3
Ce+4 or Ag+1 catalyst or dilute H2SO4/60EC
Ag+1/S2O8!1 in dilute HNO3
Dilute H2SO4
Dilute H2SO4 at 80EC or dil.HClO4/Cl!1
Dilute HNO3 at 95EC
43% K2CO3 at 75EC
Dilute; slow
HClO4; slow
pH 2
H+1

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JULY 2004

Separation Techniques
Redox Reaction

PuO2+2 6 Pu+4

PuO2+1 6 Pu+4
Pu+4 6 Pu+3

Tc+4 6 TcO4!1
TcO2(hydrated) 6
TcO4!1
TcCl6!2 6 TcO4!1

TcO4!1 6 Tc+4 or
TcO2(hyd)

U+3 6 U+4

JULY 2004

Reagent
Fe+2
V+3 or U+4
HNO2
Ag
C2O4!2
I!1
Fe+2
Sn+2
H2O2
Ti+3
Cu2O
HNO2
Zn
HNO2
NH2OH@HCl
hydroquinone
H2/Pt
I!1
HSO3!1
NH2OH@HCl
Zn
SO2
Ti+3
ascorbic acid
U+4
H2S
HNO3
H2O2
O2 (air)
Ce+4
H2O2
H2O2
Cl2
Ce+4
MnO4_1
N2H4
NH2OH
Ascorbic acid
Sn+2
Zn
Concentrated HCl
ClO4!1
Co+3 complexes

Conditions
HClO4 or HCl
HClO4
Dilute HNO3NaNO3
Dilute HCl
75EC; RT with dilute HCl
HNO3
HCl, HNO3, or H2SO4
HCl/HClO4
HNO3; continues to Pu+3 in absence of Fe+3
HClO4
45% K2CO3 75EC
HNO3/75EC
Dilute HCl
Slow
Dilute HCl, slow
Dilute HNO3
HCl
Dilute HCl
Dilute HNO3
Dilute HCl
Dilute HNO3
HCl, dilute H2SO4, or dilute HNO3/H2SO4
HNO3
Dilute HClO4
Dilute acid

Dilute H2SO4
Dilute H2SO4
Dilute H2SO4
Dilute H2SO4
Dilute HCl
6 TcCl6!2
Dilute HClO4
Dilute HClO4 or LiClO4

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Separation Techniques
Redox Reaction

Reagent

Conditions

Cr+3 and Cr+3 complexes
Dilute HClO4 or LiClO4
H2O
Dilute or concentrated HCl or H2SO4
O2 (air)
U+4 6 UO2+2
Br2
Catalyzed by Fe+3 or Mn+2
!1
BrO3
HClO4
Ce+4
Dilute HClO4
ClO3!1
Catalyzed by Fe+2 or V+5
+3
Fe
HClO2
Phenol
HCrO4!1
HNO2
Catalyzed by Fe+2
HNO3
H2O2
O2
MnO2
UO2+1 6 UO2+2
Fe+3
+2
+4
UO2 6 U
Cr+2
Eu+2
Np+3
Ti+3
+2
V and V+3
Rongalite (an aqueous
Dilute basic solution
solution of sodium
hydroxymethanesulfonate)
+2
+3
Zn(Hg)
UO2 6 U
Cr+2
UO2+1 6 U+4
H2
Zn(Hg)
(1) Compiled from: Anders, 1960; Bailar et al., 1984; Bate and Leddicotte, 1961; Cobble, 1964; Coleman, 1965;
Cotton and Wilkinson, 1988; Greenwood and Earnshaw, 1984; Hassinsky and Adloff, 1965; Kleinberg and
Cowan, 1960; Kolthoff et al., 1969; Latimer, 1952; Metz and Waterbury, 1962; Schulz and Penneman, 1986;
Weigel, 1986; and Weigel et al., 1986.

One method for the analysis of radioiodine in aqueous solutions illustrates the use of oxidation
and reduction chemistry to bring the radionuclide to a specific oxidation state so that it can be
isolated from other radionuclides and other elements (DOE, 1997, Method RP230). Iodine
species in the water sample are first oxidized to iodate (IO4!1) by sodium hypochlorite (NaClO),
and then reduced to iodide (I!1) by sodium bisulfite. The iodine is finally oxidized to molecular
iodine (I2) and extracted from most other radionuclides and elements in solution by a nonpolar
organic solvent such as carbon tetrachloride (CCl4) or chloroform (CHCl3) (see Section 14.4,
“Solvent Extraction”).
Plutonium and its tracers can be equilibrated in a reproducible mixture of oxidation states by the
rapid reduction of all forms of the ion to the +3 state, momentarily, with iodide ion (I!1) in acid
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Separation Techniques
solution. Disproportionation begins immediately, but all radionuclide forms of the analyte and
tracer begin at the same time from the same oxidation state, and a true equilibrium mixture of the
radionuclide and its tracer is achieved. All plutonium radionuclides in the same oxidation state
can be expected to behave the same chemically in subsequent separation and detection
procedures.
In addition to dissolution and separation strategies, oxidation-reduction processes are used in
several quantitation steps of radiochemical analyses. These processes include titration of the
analyte and electrochemical deposition on a target for counting.
The classical titrimetric method is not commonly employed in the quantitation of environmental
level samples because the concentrations of radionuclides in these samples are typically too low
for detection of the endpoint of the titration, even by electrometric or spectroscopic means.
However, the method is used for the determination of radionuclides in other samples containing
larger quantities of long-lived radionuclides. Millimole quantities of uranium and plutonium in
nuclear fuels have been determined by titration using methods of endpoint detection as well as
chemical indicators (IAEA, 1972). In one method, uranium in the (VI) oxidation state is first
reduced to +3 and +4 with Ti+3, then uranium in the +3 state is oxidized to +4 with air bubbles
(Baetsel and Demildt, 1972). The solution is then treated with a slight excess of Ce+4 solution of
known concentration, which oxidizes U+4 to U(VI) (as UO2+2) while being reduced, as follows:
U+4 + 2 Ce+4 6 U+6 + 2 Ce+3
(U+4 + 2 Ce+4 +2 H2O 6 UO2+2 + 2 Ce+3 + 4 H+1)
The excess Ce+4 is back-titrated with Fe+2 solution, using ferroin as indicator for the endpoint of
the titration:
Fe+2 + Ce+4 6 Fe+3 + Ce+3
Electrochemical methods are typically used in radiochemistry to reduce ions in solution, plating
them onto a target metal for counting. Americium ions (Am+3) from soil samples ultimately are
reduced from solution onto a platinum electrode by application of an electrical current in an
electrolytic cell (DOE, 1990 and 1997, Method Am-01). The amount of americium on the
electrode is determined by alpha spectrometry.
In some cases, the deposition process occurs spontaneously without the necessity of an applied
current. Polonium and lead spontaneously deposit from a solution of hydrochloric acid onto a
nickel disk at 85 EC (Blanchard, 1966). Alpha and beta counting are used to determine 210Po and
210
Pb. Wahl and Bonner (1951) contains a table of electrochemical methods used for the
oxidation and reduction of carrier-free tracers.
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Oxidation-reduction chemistry often is used to separate mixtures of transuranics. This is because
mixtures of several transuranics (e.g., U, Pu, Cm) or transition metals will generate different
oxidation states of each element as a result of inter-element redox reactions. An example would
be :
2 H2O + U+4 + 2 Pu+4 6 UO22+ + 2 Pu+3 + 4 H+
Thus, when attempting to determine plutonium (as the Pu+4 ion) in a solution containing U+4, it
would be necessary to isolate most of the plutonium from the uranium before Pu+4 can be
analyzed successfully. The isolation would take place using extraction, precipitation, or
chromatographic methods.

14.3 Complexation
14.3.1 Introduction
A complex ion is formed when a metal atom or ion bonds with one or more molecules or anions
through an atom capable of donating one or more electron pairs. A ligand is any molecule or ion
that has at least one electron pair that can be donated to the metal. The bond is called a
coordination bond, and a compound containing a complex ion is a coordination compound. The
following are several examples of the formation of complex ions:
Th+4 + 2 NO3!1 6 Th(NO3)2+2
Ra+2 + EDTA-4 * 6 Ra(EDTA)!2
U+4 + 5 CO3!2 6 U(CO3)5!6
* EDTA!4 = Ethylene diamine tetraacetate, !1(OOC)2-NH-CH2-CH2-NH-(COO)2!1
In a fundamental sense, every ion in solution can be considered complexed; there are no free or
“naked” ions. Dissolved ions are surrounded by solvent molecules. In aqueous solutions, the
complexed water molecules, referred to as the inner hydration sphere, form aquo ions that can be
either weakly or strongly bound:
Fe+2 + 6 H2O 6 Fe(H2O)6+2
From an elementary standpoint, the process of complexation is simply the dynamic process of
replacing one set of ligands, the solvent molecules, with another. The complexation of a metal
ion in aqueous solution with a ligand, L, can be expressed as:
M(H2O)n+x + L-y 6 M(H2O)n!1Lx-y + H2O

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Successive aquo groups can be replaced by other ligand groups until the complex MLnx-ny is
formed as follows:
M(H2O)n!1 Lx!y + L!y 6 M(H2O)n!2 Lx-22 y + H2O, etc.
In the absence of other complexing agents, in dilute aqueous solution solvated metal ions are
simply written as M+n for simplicity.
Ligands are classified by the number of electrons they donate to the metal to form coordination
bonds to the metal. If only one atom in the ligand is bonded to the metal, it is called a “unidentate
ligand” (from the Latin word for teeth). It is a categorization of ligands that describe the number
of atoms with electron pairs a ligand has available for donation in complex-ion formation; if two
atoms, bidentate, and so on for tridentate, tetradentate, pentadentate, and hexadentate. The term
“coordination number” is also used to indicate the number of atoms donating electrons to the
metal atom. The coordination number is 10 in U(CO3)5!6, as illustrated above. EDTA, also
illustrated above, is a hexadentate ligand, because it bonds to the metal through the four oxygen
atoms and two nitrogen atoms. Table 14.4 lists some common ligands arranged by type.
A ligand can be characterized by the nature and basicity of its ligand atom. Oxygen donors and
the fluoride ion are general complexing agents. They combine with any metal ion (cation) with a
charge of more than one. Acetates, citrates, tartrate, and β-diketones generally complex all
metals. Conversely, cyanide (CN!1), the heavy halides, sulfur donors, and—to a lesser extent—
nitrogen donors, are more selective complexing agents than the oxygen donors. These ligands do
not complex the A-metals of the periodic table; only the cations of the B-metals and the
transition metals coordinate to carbon, sulfur, nitrogen, chlorine, bromine, and iodine.
TABLE 14.4 — Common ligands
Ligand Type
Unidentate

Bidentate
Tridentate
Polydentate

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(1)

Examples
Water (H2O), halides (X ), hydroxide (OH!1), ammonia (NH3),
cyanide (CN!1), nitrite (NO2!1), thiocyanate (SCN!1), carbon
monoxide (CO)
Oxalate, ethylene diamine, citrate
Diethylene triamine, 1,3,5 triaminocyclohexane
8-hydroxyquinoline, β-diketones (thenoyltrifluoroacetone
[TTA]), ethylene diamine tetraacetic acid (EDTA), diethylene
triamine pentaacetic acid (DTPA)
Organophosphates: (octyl(phenyl)-N,N-diiso-butylcarbamoylmethylphosphine oxide [CMPO]); tributylphosphate (TBP),
trioctylphosphinic oxide (TOPO), quaternary amines (tricaprylylmethylammonium chloride [Aliquat-336®]), triisooctylamine
(TIOA), tri-n-octylamine (TnOA), macrocyclic polyethers (crown
ethers such as [18]-crown-6), cryptates
!1

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(1) Ligands are categorized by the number of electron pairs available for donation. Unidentate
ligands donate one pair of electrons; bidentate donate two pairs, etc.

14.3.2 Chelates
When a multidentate ligand is bound to the metal atom or ion by two or more electron pairs,
forming a ring structure, it is referred to as a “chelate” and the multidentate ligand is called a
“chelating agent” or reagent. Chelates are organic compounds containing two, four, or six
carboxylic acid (RCOOH) or amine (RNH2) functional groups. A chelate is effective at a pH
where the acid groups are in the anionic form as carboxylates, RCOO !1, but the nitrogen is not
protonated so that its lone pair of electrons is free for bonding. The chelate bonds to the metal
through the lone pair of electrons of these groups as bi-, tetra-, or hexadentate ligands, forming a
coordination complex with the metal. Binding through multiple sites wraps up the metal in a
claw-like fashion, thus the name chelate, which means claw. Practically all chelates form five- or
six-membered rings on coordinating with the metal. Chelates are much more stable than complex
compounds formed by unidentate reagents. Moreover, if multiple ring systems are formed with a
single metal atom or ion, stability improves. For example, EDTA, a hexadentate ligand, forms
especially stable complexes with most metals. As illustrated in Figure 14.1, EDTA has two donor
pairs from the nitrogen atoms, and four donor pairs from the oxygen atoms.

H2
C-C

..

..

N
C-C
H2

H2

O

..
..

..
..

O

..
..
..
..

H-O

H2 H2
C- C

H-O

..
..

H2
.. C - C
H-O
..
N
..
H-O
.. C - C

..

..

O

O

FIGURE 14.1 — Ethylene diamine tetraacetic acid (EDTA)
EDTA forms very stable complexes with most metal atoms because it has
two pairs of electrons available from the nitrogen atoms, and four pairs of
electrons from the oxygen atoms. It is often used as a complexing agent in a
basic solution. Under these conditions, the four carboxylic-acid groups
ionize with the loss of a hydrogen ion (H+1), forming EDTA!4, a stronger
complexing agent. EDTA is often used as a food additive to increase shelf
life, because it combines with transition metal ions that catalyze the
decomposition of food. It is also used as a water softener to remove Ca+2
and Mg+2 ions from hard water.

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Various chelating agents bind more readily to certain cations, providing the specificity for
separating ions by selective bonding. Occasionally, the complex is insoluble under the solvent
conditions used, allowing the collection of the complex by precipitation. Selectivity of a chelate
can be partially controlled by adjusting the pH of the medium to vary the net charge on its
functional groups. Different chelates provide specificity through the number of functional groups
available for bonding and the size of claw formed by the molecular structure, providing a select
fit for the diameter of a specific cation. The electron-donating atoms of the chelate form a ring
system with the metal atom when they participate in the coordination bond. In most cases,
chelates form much more stable complexes than unidentate ligands. For example, the complex
ion formed between Ni+2 and the bidentate ligand ethylenediamine (H2N-CH2-CH2-NH2, or en),
Ni(en)3+2, is almost 108 times more stable than the complex ion formed between the metal ion
and ammonia, Ni(NH3)+2.
Another class of ligands that is becoming increasingly important to the radiochemist doing
laboratory analyses is the macrocyclic polyethers, commonly called crown ethers (Horwitz et al.,
1991 and 1992a; Smith et al., 1996 and 1997). These compounds are cyclic ethers containing a
number of regularly spaced oxygen atoms. Some examples are given in Figure 14.2.

FIGURE 14.2 — Crown ethers

First identified in 1967, crown ethers have been shown to form particularly stable coordination
complexes. The term, “crown ether,” was suggested by the three-dimensional shape of the
molecule. In the common names of the crown ethers, the ring size is given in brackets, and the
number of oxygen atoms follows the word “crown.”
Crown ethers have been shown to react rapidly and with high selectivity (Gokel, 1991; Hiraoka,
1992). This property is particularly significant when a separation requires high selectivity and
efficiency in removing low-level species from complex and concentrated matrices, a situation
frequently encountered in environmental or mixed-waste analyses. Because crown ethers are
multidentate chelating ligands, they have very high formation constants. Moreover, because the
metal ion must fit within the cavity, crown ethers demonstrate some selectivity for metal ions
according to their size. Crown ethers can be designed to be very selective by changing the ring
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size, the ring substituents, the ring number, the donor atom type, etc. For example, dibenzo-18crown-6 forms a strong complex with potassium; weaker complexes with sodium, cesium, and
rubidium; and no complex with lithium or ammonium, while 12-crown-4, with its smaller cavity,
specifically complexes with lithium.
Other crown ethers are selective for radionuclide ions such as radium and UO2+2. Addition of 18crown-6 to solutions containing NpO2+2 causes the reduction of neptunium to Np(V) as NpO2+1,
which is encircled by the ether ligand (Clark et al., 1998).
14.3.3 The Formation (Stability) Constant
The stability of the complex is represented by the magnitude of an equilibrium constant
representing its formation. The complex ion, [Th(NO3)2+2], forms in two equilibrium steps:
Th+4 + NO3!1 6 Th(NO3)+3
Th(NO3)+3 + NO3!1 6 Th(NO3)2+2
The final equation is:
Th+4 + 2NO3! 6 Th(NO3)2+2
The stepwise formation (stability) constants are:
K1 =

[Th(NO 3 ) +3 ]
[Th +4 ][NO 3−1 ]

and
K2 =

[Th(NO3 ) 2 +2 ]
[Th(NO 3 ) +3 ][NO 3−1 ]

The overall formation (stability) constant is:
K f = K1 ⋅ K 2

+2

[Th(NO 3 ) 2 ]
−1
[Th +4 ][NO 3 ]2

In the Ni+2 examples cited in the preceding section, the relative stabilities of the complex ions are
represented by the values of K; for Ni(en)3+2 it is 1018.28, and for Ni(NH3)+2 it is 108.61 (Cotton and
Wilkinson, 1988).
Many radionuclides form stable complex ions and coordination compounds that are important to
the separation and determination steps in radioanalytical chemistry. Formation of a complex
changes the properties of the ion in several ways. For example:

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• Complexation of UO2+2 with carbonate to form UO2(CO3)3-4 increases the solubility of the
uranium species in groundwater (Lindsay, 1988).
• Thorium (+2) forms Th(NO3)6!2 in nitric acid solution (optimally at 7 M) that is the basis for
separation of thorium from other actinides and thorium progeny, because they do not form
anionic complexes under these conditions (Hyde, 1960).
• Radium (+2) forms a very insoluble compound with sulfate (RaSO4) but is soluble in hot
concentrated sulfuric acid because of the formation of Ra(SO4)2!2 (Kirby and Salutsky, 1964).
In addition, the complex ion in solution is in equilibrium with the free (hydrated) ion, and the
equilibrium mixture might, therefore, contain sufficient concentration of the free ion for it to be
available for other reactions, depending on the stability of the complex ion.
14.3.4 Complexation and Radiochemical Analysis
Property changes also accompany the formation of complex ions and coordination compounds
from simple radionuclide ions. These changes provide a valuable approach in radiochemistry for
isolating, separating, and measuring radionuclide concentrations, and are important in several
areas of radiochemistry.
14.3.4.1

Extraction of Laboratory Samples and Ores

Uranium ores are leached with alkaline carbonates to dissolve uranium as the UO2(CO3)3!4
complex ion after oxygen is used to convert U+4 to U(VI) (Grindler, 1962). Samples containing
refractory plutonium oxides are dissolved with the aid of a nitric acid-hydrofluoric acid solution
to produce the complex cation PuF+3 and similar cationic fluorocomplexes (Booman and Rein,
1962). Refractory silicates containing niobium (Nb) also yield to fluoride treatment. Potassium
bifluoride (KF2!1) is used as a low-temperature flux to produce a fluoride complex NbF6!1
(Willard and Rulfs, 1961; Greenwood and Earnshaw, 1984).
14.3.4.2

Separation by Solvent Extraction and Ion-Exchange Chromatography

Many ion-exchange separations of radionuclides are based on the formation of complex ions
from the metal ions in solution or the displacement of ions bound to an exchanger by complex
formation. Uranium in urine samples, for example, is partly purified by forming a chlorocomplex
of U+4 and UO2+2 ions, UCl6!2 and UO2Cl3!1, that bind preferentially to the anion-exchange
ligands in 7 M HCl. Other cations pass through the column under these conditions. Uranium is
subsequently eluted with 1 M HCl (DOE, 1990 and 1997, Method U-01).
For separation on a larger scale—such as in an industrial setting—chelates are often used in a
column chromatography or filtration unit. They are immobilized by bonding to an inert matrix,
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such as polystyrene or an alumina/silica material. A solution containing the ions to be separated
is passed continuously through the column or over the filter, where the select cations are bonded
to the chelate as the other ions pass through. Washing the column or filter with a solution at
alternate pH or ionic strength will permit the elution of the bound cation.
Thorium (+4) is bound more strongly to cation exchangers than most other cations (Hyde, 1960).
The bound thorium is separated from most other ions by washing the column with mineral acids
or other eluting agents. Even the tetrapositive plutonium ion, Pu+4, and the uranyl ion, UO2+2, are
washed off with high concentrations of HCl because they form chlorocomplexes, PuCl6!2 and
UO2Cl3!1, respectively. Thorium is then removed by eluting with a suitable complexing agent
such as oxalate, which reduces the effective concentration of Th+4, reversing the exchange
process. Using oxalate, Th(C2O4)4!4 forms and the anion is not attracted to the cation exchanger.
14.3.4.3

Formation and Dissolution of Precipitates

A classical procedure for the separation and determination of nickel (Ni) is the precipitation of
Ni+2 with dimethylglyoxime, a bidentate ligand that forms a highly selective, stable chelate
complex with the ion, Ni(C4H7N2O2!1)2 (DOE, 1997, Method RP300). Uranium in the +4
oxidation state can also be precipitated from acidic solutions with a chelating agent, cupferron
(ammonium nitrosophenylhydroxylamine, C8H5(NO)O!1NH4+1) (Grindler, 1962). In another
procedure, Co+2 can be selectively precipitated from solution as K3Co(NO2)6. In this procedure,
cobalt, which forms the largest number of complexes of all the metals, forms a complex anion
with six nitrite ligands, Co(NO2)6!3 (EPA, 1973).
In radiochemical separations and purification procedures, precipitates of radionuclides are
commonly redissolved to release the metal ion for further purification or determination. In the
determination of 90Sr, Sr+2 is separated from the bulk of the solution by direct precipitation of the
sulfate, SrSO4. The precipitate is redissolved by complexation with EDTA, Sr(EDTA)!2, to
separate it from lanthanides and actinides (DOE, 1997, Method RP520). Radium also forms a
very stable complex with EDTA. Solubilization of radium, Ra+2, coprecipitated with barium
sulfate (BaSO4) is used in the 228Ra determination of drinking water by using EDTA (EPA,
1980).
14.3.4.4

Stabilization of Ions in Solution

In some radiochemical procedures, select radionuclides are separated from other elements and
other radionuclides by stabilizing the ions as complex ions, while the other substances are
precipitated from solution. In a procedure extensively used at Oak Ridge National Laboratory
(ORNL), 95Nb is determined in solutions by taking advantage of complex-ion formation to
stabilize the Nb(V) ion in solution during several steps of the procedure (Kallmann, 1964). The
niobium sample and carrier are complexed with oxalic acid in acidic solution to prevent
precipitation of the carrier and to promote interchange between the carrier and 95Nb. Niobium is
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precipitated as the pentoxide after warming the solution to destroy the oxalate ion, separating it
from the bulk of other ions in solution. Niobium is also separated specifically from zirconium by
dissolving the zirconium oxide in hydrofluoric acid.
14.3.4.5

Detection and Determination

Compleximetric titration of metal ions with EDTA using colorimetric indicators to detect the
endpoint can be used for determination procedures. Uranium does not form a selective complex
with EDTA, but this chelate has been used to titrate pure uranium solutions (Grindler, 1962). The
soluble EDTA complex of thorium is the basis of a titrimetric determination of small amounts of
thorium (Hyde, 1960).
Spectrometric determinations are also based on the formation of complex ions. Microgram
quantities of uranium are determined by the absorbance at 415 nm (a colorimetric determination)
of the uranyl chelate complex with dibenzoylmethane, C6H5-CO-CH2-CO-C6H5 (Grindler, 1962).

14.4 Solvent Extraction
14.4.1 Extraction Principles
Solvent extraction has been an important separation technique since the early days of the
Manhattan Project, when scientists extracted uranyl nitrate into diethyl ether to purify the
uranium used in the first reactors. Solvent extraction, or liquid-liquid extraction, is a technique
used both in the laboratory and on the industrial scale. However, current laboratory trends are
away from this technique, mainly because of the costs of materials and because it is becoming
more difficult and costly to dispose of the mixed waste generated from the large volumes of
solvents required. The technique also tends to be labor intensive because of the need for multiple
extractions using separatory funnels. Nonetheless, solvent extraction remains a powerful
separation technique worthy of consideration.
Solvent extraction refers to the process of selectively removing a solute from a liquid mixture
with a solvent. As a separation technique, it is a partitioning process based on the unequal
distribution of the solute (A) between two immiscible solvents, usually water (aq) and an organic
liquid (org):
Aaq W Aorg
The solute can be in a solid or liquid form. The extracting solvent can be water, a water-miscible
solvent, or a water-immiscible solvent; but it must be insoluble in the solvent of the liquid
mixture. Solutes exhibit different solubilities in various solvents. Therefore, the choice of
extracting solvent will depend upon the properties of solute, the liquid mixture, as well as other
requirements of the experimental procedure. The solvents in many applications are water and a
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nonpolar organic liquid, such as hexane or diethyl ether, but other solvent pairs are commonly
used. In general terms, the solute to be removed along with impurities or interfering analytes to
be separated are already dissolved in one of the solvents (water, for example). In this example, a
nonpolar organic solvent is added and the two are thoroughly mixed, usually by shaking in a
separatory funnel. Shaking produces a fine dispersion of each solvent in the other that will
separate into two distinct layers after standing for several minutes. The more dense solvent will
form as the bottom layer. Separation is achieved because the solute and accompanying impurities
or analytes have different solubilities in the two solvents. The solute, for example, might
preferentially remain in the aqueous phase, while the impurities or analyte selectively dissolve in
the organic phase. The impurities and analyte are extracted from the aqueous layer into the
organic layer. Alternatively, the solute might be more soluble in the organic solvent and will be
extracted from the aqueous layer into the organic layer, leaving the impurities behind in the
aqueous layer.
14.4.2 Distribution Coefficient
The different solubilities of a solute in the solvent pairs of an extraction system are described by
the distribution or partition coefficient, Kd. The coefficient is an equilibrium constant that
represents the solubility of the solute in one solvent relative to its solubility in another solvent.
Once equilibrium is established, the concentration of solute in one phase has a direct relationship
to the solute concentration in the other phase. This is expressed mathematically by:
Kd =

[A org ]
[A aq ]

where [Aorg] and [Aaq] are the concentration of the solute in the organic and aqueous phase
respectively, and Kd is a constant. The concentrations are typically expressed in units of moles/kg
(molality) or g/g; therefore, the constant is unitless. These solubilities usually represent saturated
concentrations for the solute in each solvent. Because the solubilities vary with temperature, the
coefficient is temperature-dependent, but not by a constant factor. Wahl and Bonner (1951)
contains a table of solvent extraction systems for carrier-free tracers containing laboratory
conditions and distribution coefficients.
A distribution coefficient of 90 for a solute in a hexane/water system, for example, means that
the solute is 90 times more soluble at saturation conditions in hexane than in water, but note that
some of the water still contains a small amount of the solute. Solvent extraction selectively
dissolves the solute in one solvent, but it does not remove the solute completely from the other
solvent. A larger coefficient would indicate that, after extraction, more solute would be
distributed in hexane relative to water, but a small quantity would still be in the water. Solvent
extraction procedures often use repeated extractions to extract a solute quantitatively from a
liquid mixture.
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The expression of the distribution law is only a very useful approximation; it is not thermodynamically rigorous, nor does it account for situations in which the solute is involved in a
chemical reaction, such as dissociation or association, in either phase. Consider, for example,
dimerization in the organic phase:
2Aorg W (A)2, org
where the distribution ratio, D, is an alternate form of the distribution coefficient expressed by:
D = ([Aorg]monomer + [Aorg]dimer)/[Aaq]

or

D = ([Aorg] + 2 [(A)2, org]) /[Aaq]
Because the concentration of the monomer that represents the dimeric form of the solute is twice
that of the concentration of the dimer:
[Aorg]dimer = 2 [(A)2, org]
Substitution of Kd produces:
D = Kd (1 + 2 K2 [Aorg])
where K2 is the dimerization constant, K2 = [(A)2, org]/[Aorg]2. Because dimerization decreases the
concentration of the monomer, the species that takes part directly in the phase partition, the
overall distribution increases.
14.4.3 Extraction Technique
There is extensive literature on the topic of extraction techniques, but only a few sources are
listed here. The theory of solvent extraction is covered thoroughly in Irving and Williams (1961),
Lo et al. (1983), and Dean (1995). The journal Solvent Extraction and Ion Exchange is an
excellent source for current advances in this field. A practical discussion on the basics of solvent
extraction is found in Korkisch (1969). The discussion applies to a metallic element in solution
as a cation extracted by a nonpolar solvent:
“In solvent extraction, the element which is to be separated, contained in an aqueous solution,
is converted to a compound which is soluble in an organic solvent. The organic solvent must
be virtually immiscible with water. By shaking the aqueous solution with the organic solvent
(extractant) in a separating funnel, the element is extracted into the organic phase. After
allowing the aqueous and organic phases to separate in the funnel, the organic extract is
removed from contact with the aqueous layer. This single-stage batch extraction method is
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employed when Kd is relatively large and for a simple separation it is essential that the
distribution coefficients of the metal ions to be separated be sufficiently different. As in the
case of ion exchange, the effectiveness of separation is usually expressed by means of the
separation factor which is given by the ratio of the distribution coefficients of two different
elements which were determined under identical experimental conditions. This ratio
determines the separability of two elements by liquid-liquid extraction. Separations can only
be achieved if this ratio shows a value which is different from unity and they are clean and
can be quickly and easily achieved where one of the distribution coefficients is relatively
large and the other very small (high separation factor).
“In those extractions where the separation factor approaches unity, it is necessary to employ
continuous extraction or fractionation methods. With the latter techniques distribution,
transfer and recombination of various fractions are performed a sufficient number of times to
achieve separation. In continuous extraction use is made of a continuous flow of immiscible
solvent through the solution or a continuous counter-current flow of both phases. In
continuous extraction the spent solvent is stripped and recycled by distillation, or fresh
solvent is added continuously from a reservoir. Continuous counter-current extraction
involves a process where the two liquid phases are caused to flow counter to each other.
Large-scale separations are usually performed using this technique.
“When employing liquid-liquid extraction techniques, one of the most important
considerations is the selection of a suitable organic solvent. Apart from the fact already
mentioned that it must be virtually immiscible with water, the solubility of the extracted
compound in the solvent must be high if a good separation is to be obtained. Furthermore, it
has to be selective, i.e., has to show the ability to extract one component of a solution in
preference to another. Although the selectivity of a solvent for a given component can be
determined from phase diagrams, it is a little-used procedure in analytical chemistry. The
principal difficulty is simply that too few phase diagrams exist in the literature. The result is
that the choice of an extractant is based on either experience or semi-empirical considerations. As a rule, however, polar solvents are used for the extraction of polar substances from
nonpolar media, and vice versa. Certainly the interactions of solute and solvent will have an
effect on the selectivity of the solvent. If the solute is readily solvated by a given solvent, then
it will be soluble in that solvent. Hydrogen bond formation between solute and solvent
influences solubility and selectivity.
“Almost as important as the selectivity of the extractant is the recovery of the solute from the
organic extract. Recovery can be achieved by distillation or evaporation of the solvent,
provided that the solute is nonvolatile and thermally stable. This technique is, however, less
frequently used than the principle of back extraction (stripping) which involves the treatment
of the organic extract with an aqueous solution containing a reagent which causes the
extracted solute to pass quantitatively into the aqueous layer...
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“In solvent extraction the specific gravity of the extractant in relation to the aqueous phase is
important. The greater the difference in the solvent densities, the faster will be the rate at
which the immiscible layers separate. Emulsions are more easily produced when the densities
of the two solvents are similar. Sometimes troublesome emulsions can be broken by
introducing a strong electrolyte into the system or by the addition of small quantities of an
aliphatic alcohol”
Korkisch (1969) continues:
“Liquid-liquid extraction can be applied to the analysis of inorganic materials in two different
ways.
(a) Where the element or elements to be determined are extracted into the organic phase.
(b) Where the interfering elements are removed by extraction, leaving the element or
elements to be determined in the aqueous phase.
“Solvent extraction separations are mainly dependent for their successful operation upon the
distribution ratio of the species between the organic and aqueous phase and the pH and salt
concentration of the aqueous phase. Much of the selectivity which is achieved in liquid-liquid
extraction is dependent upon adequate control of the pH of the solution. The addition of
masking agents such as EDTA and cyanide can greatly improve selectivity, but they too are
dependent upon the pH of the solution to exert their full effect. In many cases complete
extractions and separations are obtained only in the presence of salting-out agent. An
example is the extraction of uranyl nitrate. In the presence of additional nitrate, the increase
in the concentration of the nitrate ion in the aqueous solution shifts the equilibrium between
the uranyl ion and the nitrate complexes toward the formation of the latter, and this facilitates
a more complete extraction of the uranium into the organic solvent. At the same time, the
salting-out agent has another, more general, effect: as its affinity for water is large, it
becomes hydrated by the water molecules so that the substance to be extracted is really
dissolved in a smaller amount of water, and this is the same as if the concentration in the
solution were increased. As a result, the distribution coefficient between the aqueous and the
organic phases is increased. As a rule the salting-out agent also lowers the solubility of the
extractant in the aqueous phase, and this is often important in separations by extraction. The
efficiency of the salting-out action depends upon the nature and the concentration of the
salting-out agent. For the same molar concentration of the salting-out agent its action
increases with an increase in the charge and decrease in the radius of its cation.”
A hydrated metal ion will always prefer the aqueous phase to the organic phase because of
hydrogen bonding and dipole interaction in the aqueous phase. Therefore, to get the metal ion to
extract, some or all of the inner hydration sphere must be removed. The resulting complex must
be neutrally charged and organophilic. Removal of the hydration sphere is accomplished by
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coordination with an anion to form a neutral complex. Neutral complexes will generally be more
soluble in an organic phase. Larger complexing anions favor the solubility in the organic phase.
Extracting agents are thus divided into three classes: polydentate organic anions, neutral organic
molecules, and large organic cations. Many of the multidentate ligands discussed previously are
used in solvent extraction systems.
The radioanalytical procedure for uranium and thorium employs solvent extraction to separate
the analytes before alpha counting (EPA, 1984). An aqueous solution of the two is extracted with
a 10 percent solution of triisooctylamine (TIOA) in para-xylene to remove uranium, leaving
thorium in the water (Grinder, 1962). Each solution is further processed to recover the respective
radionuclides for separate counting.
14.4.4 Solvent Extraction and Radiochemical Analysis
In many purification procedures, separated solutions are used directly in further isolation steps. If
necessary, the substances can be collected by distillation or evaporation of the respective
solvents. In the uranium/thorium procedure described above, the aqueous layer containing
thorium is evaporated, and the thorium is redissolved in an alternate solution before it is purified
further. In other cases, the solution is extracted again to take up the solute in another solvent
before the next step in the procedure. Uranium in TIOA/p-xylene, for example, is extracted back
into a nitric acid solution for additional purification (EPA, 1984).
In some solvent-extraction procedures, more than one extraction step is required for the
quantitative removal of a solute from its original solvent. The solute is more soluble in one
component of the solvent pair, but not completely insoluble in the other component, so
successive extractions of the aqueous solution of the solute by the organic solvent will remove
more and more of the solute from the water until virtually none remains in the aqueous layer.
Extraction of uranium with TIOA/p-xylene, for example, requires two extractions before
quantitative removal is achieved (EPA, 1984). The organic layers containing the uranium are
then combined into one solution for additional processing.
Solvent extraction is greatly influenced by the chemical form (ionic or molecular) of the solute to
be extracted, because different forms of the solute can have different solubilities in the solvents.
In the uranium/thorium procedure described above, uranium is extracted from water by TIOA/
hydrochloric acid, but it is stripped from the amine solution when extracted with nitric acid.
Simply changing the anion of uranium and TIOA from chloride to nitrate significantly alters the
complex stability of uranium and TIOA.
Organic amines are sometimes converted to their cationic forms, which are much more soluble in
water and much less soluble in organic solvents. The amine is converted to the corresponding
ammonium salt by an acid, such as hydrochloric acid:

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RNH2 + HCl 6 RNH3+1Cl!1
Correspondingly, carboxylic acids are converted to their carboxylates that are more soluble in
water and less soluble in organic solvents. They are produced by treating the carboxylic acid with
a base, such as sodium hydroxide:
RCOOH + NaOH 6 RCOO!1Na+1 + H2O
Multidentate organic anions that form chelates are important extracting agents. These reagents,
such as the β-diketonates and thenoyltrifluoroacetone (TTA) (Ahrland, 1986), are commonly
used for extracting the actinide elements. When the aqueous solution and organic phase come
into contact with one another, the chelating agent dissolves in the aqueous phase, ionizes, and
complexes the metal ion; the resulting metal chelate subsequently dissolves in the organic phase.
A number of organophosphorus compounds are also efficient extractants because they and their
complexes are readily soluble in organic solvents. The actinide MO2+2 and actinide +4 ions are
very effectively extracted by reagents such as bis(2-ethylhexyl) phosphoric acid (HDEHP) and
dibutylphosphoric acid (HDBP) (Cadieux and Reboul, 1996).
Among the neutral compounds, alcohols, ethers, and ketones have been commonly employed as
extractants. Methyl isobutyl ketone was used in one of the early large-scale processes (the Redox
process) to recover uranium and plutonium from irradiated fuel (Choppin et al., 1995). However,
the most widely used neutral extractants are the organophosphorus compounds such as TBP
(tributylphosphate). The actinide elements thorium, uranium, neptunium, and plutonium easily
form complexes with TBP (Choppin et al., 1995). Salting-out agents such as HNO3 and Al(NO3)3
are commonly employed to increase extraction in these systems. This chemistry is the basis of
the Purex process used to reprocess spent nuclear fuel (Choppin et al., 1995).
An important addition to the Purex process is the solvent extraction procedure known as TRUEX
(Trans Uranium Extraction). This process uses the bifunctional extractant CMPO ([octyl
(phenyl)]-N,N-diisobutylcarbonylmethylphosphine oxide) to remove transuranium elements from
the waste solutions generated in the Purex process. This type of compound extracts actinides at
high acidities, and can be stripped at low acidity or with complexing agents. Many of the recent
laboratory procedures for biological waste and environmental samples are based upon this
approach (see Section 14.4.5.1, “Extraction Chromatography Columns”).
The amines, especially the tertiary and quaternary amines, are strong cationic extractants. These
strong bases form complexes with actinide metal cations. The extraction efficiency improves
when the alkyl groups have long carbon chains, such as in tri-n-octylamine (TnOA) or TIOA.
The pertechnetate ion (TcO4!1) is also extracted by these cationic extractants (Chen, 1990).
Table 14.5 lists common solvent extraction procedures for some radionuclides of interest and
includes the examples described above.
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TABLE 14.5 — Radioanalytical methods employing solvent extraction (1)
Analyte
89/90

Sr

TcO4!

99

Extraction Conditions (Reference)
From soils and sediments with dicyclohexano-18-crown-6 in trichloromethane with back
extraction with EDTA (Pimpl, 1995)
From dilute H2SO4 solutions into a 5% TnOA in xylene mixture and back extracted with NaOH
(Golchert and Sedlet, 1969; Chen, 1990); from dilute H2SO4, HNO3, and HCl solutions into a
5% TnOA in xylene (Dale et al., 1996); from HNO3 into 30% TnOA in xylene and back
extracted with NaOH (Hirano, 1989); from dilute H2SO4 solutions into TBP (Holm et al., 1984;
Garcia-Leon, 1990); the tetraphenyl arsonium complex of Tc into chloroform (Martin and
Hylko, 1987); from K2CO3 with methyl ethyl ketone (Paducah R-46); from alkaline nuclearwaste media with crown ethers (Bonnesen et al., 1995)

210

As lead bromide from bone, food, urine, feces, blood, air, and water with Aliquat-336® (DOE,
1990 and 1997, Method Pb-01; Morse and Welford, 1971)

Radium through
Californium

From soil following KF-pyrosulfate fusion and concentration by barium sulfate precipitation
with Aliquat-336® in xylene (Sill et al., 1974)

Actinides

From water following concentration by ferric hydroxide precipitation and group separation by
bismuth phosphate precipitation, uranium extracted by TOPO, plutonium and neptunium
extracted by TIOA from strong HCl, and thorium separated from americium and curium by
extraction with TOPO (EPA, 1980, Method 907.0)

Pb

And other metals from TOPO (NAS-NS 3102) and from high-molecular weight amines such as
TIOA (NAS-NS 3101).
Uranium and plutonium from HCl with TIOA (Moore, 1958)
From nitric acid wastes using the TRUEX process with CMPO (Horwitz et al., 1985 and 1987)
With various extractive scintillators followed by PERALS® spectrometry (McDowell 1986 and
1992); with HDEHP after extraction chromatography followed by PERALS® spectrometry
(Cadieux and Reboul, 1996)
Thorium

From aqueous samples after ion exchange with TTA, TIOA, or Aliquat-336® (DOE, 1997,
Method RP570)

Uranium

From waters with ethyl acetate and magnesium nitrate as salting-out agent (EPA, 1980, Method
908.1); with URAEX™ followed by PERALS® spectrometry (Leyba et al., 1995)
From soil, vegetation, fecal ash, and bone ash with Alamine-336 (DOE, 1990 and 1997,
Methods Se-01, U-03)

(1) This list is representative of the methods found in the literature. It is not an exhaustive compilation, nor does it
imply preference over methods not listed.

14.4.5 Solid-Phase Extraction
A technique closely related to solvent extraction is solid-phase extraction (SPE). SPE is a
solvent-extraction system in which one of the phases is made stationary by adsorption onto a
solid support, usually silica, and the other liquid phase is mobile. Small columns or membranes
are used in the SPE approach. Many of the same extracting agents used in solvent extraction can
be used in these systems. SPE is becoming widely accepted as an excellent substitute for liquidliquid extraction because it is generally faster, more efficient, and generates less waste.

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14.4.5.1

Extraction Chromatography Columns

Over the past decade, extraction chromatography methods have gained wide acceptance in the
radiochemistry community as new extraction chromatographic resins have become commercially
available, such as Sr, TRU®, and TEVA® resins (Eichrom Technologies, Inc., Darien, IL) (Dietz
and Horwitz, 1993; Horwitz et al., 1991, 1992a, and 1993). These resins are composed of extractant materials, such as CMPO and 4,4'(5')-bis(t-butylcyclohexano)-18-crown-6, absorbed onto an
inert polymeric support matrix. They are most frequently used in a column rather than a batch
mode.
Another example of the advances in the area is the use of fibrous discs impregnated with highmolecular-weight chelates that select for certain elements such as Cs, Sr, and Tc (Empore Discs,
3M Company, and the TEVA® Disc, Eichrom Technologies, Inc.). Many of the traditional
methods based upon repetitive precipitations, or solvent extraction in separatory funnels, have
been replaced by this strategy. This approach allows for the specificity of liquid-liquid extraction
with the convenience of column chromatography. Numerous papers detailing the determination
of radionuclides by this technique have been published recently, and examples are cited in Table
14.6.
TABLE 14.6 — Radioanalytical methods employing extraction chromatography (1)
Analyte

Ligand

Method Citations

Ni-59/63

dimethylglyoxime

Aqueous samples (DOE, 1997)

Sr-89/90

4,4'(5')-bis(t-butyl-cyclohexano)-18crown-6 in n-octanol

Biological, Environmental, and Nuclear Waste (Horwitz
et al., 1991 and 1992a); Water (ASTM, D5811-95;
DOE, 1997, Method RP500); Urine (Dietz and Horwitz,
1992; Alvarez and Navarro, 1996); Milk (Jeter and
Grob, 1994); Geological Materials (Pin and Bassin,
1992)

Sr-90

octyl(phenyl)-N,N-diisobutylcarbamoylmethylphosphine oxide
(CMPO) in tributyl phosphate

Brines (Bunzl et al., 1996)

Y-90

4,4'(5')-bis(t-butyl-cyclohexano)-18crown-6 in n-octanol

Medical applications (Dietz and Horwitz, 1992)

Tc-99

Aliquat-336N

Low-level radioactive waste (Banavali, 1995); Water
(Sullivan et al., 1993; DOE, 1997, Method RP550)

Pb-210

4,4'(5')-bis(t-butyl-cyclohexano)-18crown-6 in isodecanol

Water (DOE, 1997, Method RP280); Geological
materials (Horwitz et al., 1994; Woittiez and Kroon,
1995); complex metal ores (Gale, 1996)

Ra-228

CMPO in tributyl phosphate or HDEHP Natural waters (Burnett et al., 1995); Volcanic rocks
impregnated in Amberlite XAD-7
(Chabaux, 1994)

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Analyte

Ligand

Method Citations

Rare earths

diamyl,amylphosphonate

Actinide-containing matrices (Carney, 1995)

CMPO in tributyl phosphate and
HDEHP impregnated in Amberlite
XAD-7

Sequential separation of light rare earths, U, and Th in
geological materials (Pin et al., 1996)

CMPO in tributyl phosphate and 4,4'(5')bis(t-butyl-cyclohexano)-18-crown-6 in
n-octanol
Concomitant separation of Sr, Sm, and Nd in silicate
samples (Pin et al., 1994)
Actinides

CMPO in tributyl phosphate

Air filters (Berne, 1995); Waters (Berne, 1995); Groupscreening (DOE, 1997, Method RP725); Urine (Horwitz
et al., 1990; Nguyen et al., 1996); Acidic media
(Horwitz, 1993; DOE, 1997); Soil and sludge (Smith et
al., 1995; Kaye et al., 1995); Environmental (Bunzl and
Kracke, 1994)

diamyl,amylphosphonate

Acidic media (Horwitz et al., 1992b)

tri-n-octylphosphine oxide [TOPO] and Environmental and industrial samples (Testa et al.,
HDEHP
1995)
(1) This list is representative of the methods found in the literature. It is not complete, nor does it imply preference
over methods not listed.

14.4.5.2

Extraction Membranes

SPE membranes have also become a popular approach to sample preparation for organic
compounds in aqueous samples over the past decade. As of 1995, 22 methods employing SPE
disks have been accepted by the U.S. Environmental Protection Agency. More recently, disks
have been developed for specific radionuclides, such as technetium, strontium, and radium
(DOE, 1990 and 1997; Orlandini et al., 1997; Smith et al., 1996 and 1997).
These SPE membranes significantly reduce extraction time and reagent use in the processing of
large environmental water samples. Samples typically are processed through the membranes at
flow rates of at least 50 mL/min; a 1 L sample can be processed in as little as 20 minutes.
Moreover, these selective-membranes often can be counted directly, thereby condensing sample
preparation and counting source preparation into a single step. Many of the hazardous reagents
associated with more traditional methods are eliminated in this approach, and these membranebased extractions use up to 90 percent less solvent than liquid-liquid extractions. The sorbent
particles embedded in the membrane are extremely small and evenly distributed, thereby
eliminating the problem of channeling that is associated with columns.

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14.4.6 Advantages and Disadvantages of Solvent Extraction
14.4.6.1

Advantages of Liquid-Liquid Solvent Extraction

• Lends itself to rapid and very selective separations that are usually highly efficient.
• Partition coefficients are often approximately independent of concentration down to tracer
levels and, therefore, can be applied to a wide range of concentrations.
• Can usually be followed by back-extraction into aqueous solvents or, in some cases, the
solution can be used directly in subsequent procedures. This also provides significant preanalysis concentration of the analyte.
• Wide scope of applications—the composition of the organic phase and the nature of
complexing or binding agents can be varied so that the number of practical combinations is
virtually unlimited.
• Can be performed with simple equipment, but can also be automated.
14.4.6.2

Disadvantages of Liquid-Liquid Solvent Extraction

• Cumbersome for a large number of samples or for large samples.
• Often requires toxic or flammable solvents.
• Can be time consuming, especially if attainment of equilibrium is slow.
• Can require costly amounts of organic solvents and generate large volumes of organic waste.
• Can be affected by small impurities in the solvent(s).
• Multiple extractions might be required, thereby increasing time, consumption of materials,
and generation of waste.
• Formation of emulsions can interfere with the phase-separation process.
• Counter-current process can be complicated and can require complicated equipment.
• Alteration of chemical form can change, going from one phase to the other, thereby altering
the distribution coefficient and effectiveness of the extraction.
• Tracer-levels of analytes can form radiocolloids that cannot be extracted, dissociate into less
soluble forms, or adsorb on the container surface or onto impurities in the system.
14.4.6.3

Advantages of Solid-Phase Extraction Media

• Column/filter extraction may be unattended.
• Column/filter extraction is very selective.
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• Generates a low volume of waste, can often be applied to samples dissolved in very acidic
media.
• Requires relatively inexpensive equipment.
• In may cases can be correlated with liquid/liquid extraction.
14.4.6.4

Disadvantages of Solid-Phase Extraction Media

• Extraction columns cannot be reused—a cost factor.
• Any suspended matter may be filtered by the media, carrying contaminants into the next step
of the separation or analysis.
• Flow rate through columns are generally slow (1-3 mL/min).

14.5 Volatilization and Distillation
14.5.1 Introduction
Differences in vapor pressures of elements or their compounds can be exploited for the
separation of radionuclides. Friedlander et al. (1981), describes the process:
“The most straightforward application is the removal of radioactive rare gases from aqueous
solutions or melts by sweeping an inert gas or helium. The volatility of ... compounds ... can
be used to effect separations ... by distillation ... Distillation and volatilization methods often
give clean separations, provided that proper precautions are taken to avoid contamination of
the distillate by spray or mechanical entrapment. Most volatilization methods can be done
without specific carriers, but some nonisotopic carrier gas might be required. Precautions are
sometimes necessary to avoid loss of volatile radioactive substances during the dissolving of
irradiated targets or during irradiation itself.”
Similar precautions are also advisable during the solubilization of samples containing volatile
elements or compounds (Chapter 13, Sample Dissolution).
14.5.2 Volatilization Principles
Volatilization particularly provides a rapid and often selective method of separation for a wide
range of elements (McMillan, 1975). A list of the elements that can be separated by volatilization
and their chemical form(s) upon separation are given in Table 14.7.

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14-37

Mg

Ca

Sr
d

Ba
a

Ra

Na
a

K
a

Rb
a

Cs
a

Fr
a
Pr

Hf
d

Nb
d

V
d

Th** Pa
d

Ce*

Hf
d

La*
Ac**

Zr
d

Ti
d

Y

Sc

(From Coomber, 1975)

Key to volatile form of element:

Be

Li
a

H
abcd

U
d

Np
d

Pm

Re
cd

Tc
cd

Mn
c*

Am

Sm

Os
cd

Ru
cd

Fe
d

Cm

Eu

Ir
d

Rh
a

Co

Bk

Gd

Pt

Pd

Ni

Cf

Tb

Au
a

Ag
a

Cu

Es

Dy

Hg
ad

Cd
a

Zn

Fm

Ho

Tl
a

In
a

Ga
bd

Mv

Er

Pb

Sm
bd

Ge
bd

Si
bd

C
bcd

No

Tm

Bi
ab

Sb
bd

Yb

Po
ad

Te
bcd

As
Se
abcd bcd

Lu

At
ab

I
abd

Br
abd

Rn
ad

Xe
ad

Kr
ad

P
S
Cl
Ar
abcd abcd abcd a

N
O
F
Ne
abcd abcd abcd a

a - Element; b - Hydride; c - Oxide; c * - Permanganic acid; c+ - Boric acid; d - Halides;
d* - Chromyl chloride

Nd

W
d

Mo
d

Cr
d*

Al
d

B
bc+d

He
a

Separation Techniques

TABLE 14.7 — Elements separable by volatilization as certain species

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McMillan (1975) states:
“While many of the volatile species are commonly encountered and a large proportion can be
produced from aqueous solutions, a significant number are rarely met. The volatilization of
highly reactive materials and those with high boiling points are only used in special
circumstances, e.g., for very rapid separations. ... Many other volatile compounds have been
used to separate the elements, including sulphides, carbonyls, stable organic complexes ... ,
and fluorinated β-diketones for the lanthanides.
“Separation ... is achieved by differentiation during the volatilization process, fractionation
by transfer, and selective collection. Gaseous evolution can be controlled by making use of
differences in vapor pressure with temperature, adjustment of the oxidation state of the
element in solution or by alteration of the matrix, in order to change the chemical
combination of the element. Once gaseous, additional separation is possible and physical
processes can be adopted such as gas chromatography, zone refining, fractional distillation,
electrostatic precipitation, filtration of condensed phases and low temperature trapping.
Chemical methods used are mainly based on the selective trapping of interfering substances
by solid or liquid reagents. The methods of preferential collection of the species sought are
similar to those used in the transfer stage.”
Both solid and liquid samples can be used in volatilization separations (Krivan, 1986):
“With solid samples, there are several types of separation methods. The most important of
them are ones in which (1) the gas forms a volatile compound with only the trace elements
and not the matrix, (2) the gas forms a volatile compound with the matrix but not the trace
elements, and (3) volatile compounds are formed with both the matrix and the trace elements.
Different gases have been used in separation by volatilization, including inert gases N2, He,
and Ar and the reactive gases H2O, O2, H2, ... F2, and HF. The apparatus usually consists of
three parts: gas regulation and purification, oven with temperature programming and control,
and condensation or adsorption with temperature regulation.
“The radiotracer technique provides the best way to determine the recoveries of trace
elements in the volatilization process and to optimize the separation with respect to the
pertinent experimental parameters.”
14.5.3 Distillation Principles
Distillation is the separation of a volatile component(s) of a mixture by vaporization at the
boiling point of the mixture and subsequent condensation of the vapor. The vapor produced on
boiling the mixture is richer in the more volatile component—the component with the higher
vapor pressure (partial pressure) and correspondingly lower boiling point. The process of
distillation, therefore, essentially takes advantage of the differences in the boiling points of the
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constituents to separate a mixture into its components. It is a useful separation tool if the analyte
is volatile or can be transformed into a volatile compound. Most inorganic applications of
distillation involve batch distillation, whereas most organic applications require some type of
fractional distillation. In a simple batch distillation, the sample solution containing a single
volatile component or components with widely separated boiling points is placed in a distillation
flask, boiling is initiated, and the vapors are then continuously removed, condensed, and
collected. Mixtures containing multiple volatile components require fractional distillation, which
employs repeated vaporization-condensation cycles for separation, and is commonly performed
in a fractionation column for that purpose. The column allows the cycles to occur in one
operation, and the separated component is collected after the last condensation.
Distillation has been widely used for separating organic mixtures but this approach has less
applicability in inorganic analysis (Korkisch, 1969). Korkisch (1969) states: “Nevertheless, some
of the elements of interest to radiochemists can be very effectively separated by distillation as
their volatile chlorides, bromides, and oxides .... [T]hese elements are germanium (Ge), selenium
(Se), technetium (Tc), rhenium (Re), ruthenium (Ru), and osmium (Os).” (Also see DOE, 1997,
Method RP530). Two common analytes determined through distillation, tritium and 226Ra, by
radon emanation are discussed below.
Specific distillation principles are commonly found in chemistry reference and textbooks. For a
theoretical discussion of distillation see Peters (1974) and Perry and Weisberger (1965).
Distillation procedures are discussed for many inorganic applications in Dean (1995) and for less
common radioanalytes in DeVoe (1962) and Kuska and Meinke (1961).
14.5.4 Separations in Radiochemical Analysis
The best known use of distillation in radiochemical analysis is in the determination of tritium
(EPA, 1984; DOE, 1997). Water is the carrier as simple distillation is used to separate tritium
from water or soil samples. For determination of tritium, the aqueous sample is treated with a
small amount of sodium hydroxide (NaOH) and potassium permanganate (KMnO4), and it is then
distilled. The early distillate is discarded, and a portion of the distillate is collected for tritium
determination by liquid scintillation counting. The alkaline treatment prevents other radionuclides, such as radioiodine or radiocarbon, from distilling over with the tritium (3H), and the
permanganate (MnO4!1) treatment destroys trace organic material in the sample that could cause
quenching during the counting procedure.
Larger samples are distilled using a round-bottom flask, while a MICRO DIST® tube can be used
for smaller samples (DOE, 1997, Method RP580). The distillate can be added directly to a liquid
scintillation cocktail (EPA, 1980, Method 906.0), or further enriched by acid electrolysis (DOE,
1990 and 1997, Method 3H-01) or alkaline electrolysis (DOE, 1990 and 1997, Method 3H-02).
Iodine is separated from aqueous samples by distillation from acidic solutions into alkaline
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solutions (EPA, 1973). Iodide (I!1) is added as carrier; but nitric acid (HNO3) as part of the acid
solution, oxidizes the anion to molecular iodine as the mixture is heated for distillation.
One determination of 79Se employs an optional purification step, distillation of the metal as
selenous acid, H2SeO3 (DOE, 1997, Method RP530). The solution is maintained with excess
bromine (Br2) and hydrobromic acid (HBr) to hold the selenium in the oxyacid form during the
distillation. Technetium can be separated from other elements, or can be separated from ruthenium, osmium, or rhenium by distillation of their oxides (Friedlander et al., 1981). Metals are
sometimes distilled in their elemental form—polonium in bismuth or lead (McMillan, 1975).
Radium-226 in solution can be determined by de-emanating its gaseous progeny 222Rn into an
ionization chamber or scintillation cell. Generally, the procedure initially involves the concentration of radium by coprecipitation with barium sulfate (BaSO4). The barium sulfate is then
dissolved in an EDTA solution, transferred to a sealed bubbler, and stored to allow for the
ingrowth of 222Rn. Following sufficient in-growth, the 222Rn is de-emanated by purging the
solution with an inert gas, such as helium (He) or argon (Ar), and is transferred via a drying tube
to a scintillation cell or ionization chamber. After the short-lived 222Rn progeny have reached
secular equilibrium with the 222Rn (approximately four hours), the sample is counted to determine
alpha activity (EPA, 1980, Method 903.1; DOE, 1990 and 1997, Methods Ra-01 through Ra-07;
Sedlet, 1966; Lucas, 1990).
When processing samples containing radon, care should be taken to guard against the inadvertent
loss of the gas or contamination of the distillation apparatus. Radon can be adsorbed on, or
permeate through, materials used in its handling. Diffusion through rubber and plastic tubing or
through polyethylene bottles has been observed. Because radon is soluble in many organic
compounds, impurities, including greases used in ground-glass connections, can increase
adsorption.
14.5.5 Advantages and Disadvantages of Volatilization
14.5.5.1
•
•
•
•

Advantages

Can be very selective, producing clean separations.
Very rapid, especially with high-vacuum equipment.
Can be performed from solid or liquid samples.
Most can be performed without a specific carrier gas.

14.5.5.2

Disadvantages

• Relatively few volatile elements or inorganic compounds are available.
• Atmosphere can alter the nature of a volatile form of the tracer or surface material.
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• Effects of experimental parameters (carrier gas, gas flow, temperature, time, and recovery)
are highly variable.
• Precautions are sometimes necessary to avoid loss of volatile radionuclide substances during
subsequent procedures.
• Some systems require high-temperature, complex equipment.
• Contamination of distillate by carrier, spray, or mechanical entrapment is a potential problem.

14.6 Electrodeposition
14.6.1 Electrodeposition Principles
Radionuclides in solution as ions can be deposited (plated) by electrochemical reactions (redox
reactions) onto an electrode, either by a spontaneous process (produced by a favorable electrode
potential existing between the ion and electrode) or by a nonspontaneous process (requiring the
application of an external voltage (potential) (Section 14.2, “Oxidation-Reduction Processes”).
Spontaneous electrochemical processes are described by the Nernst equation, which relates the
electrode potential of the reaction to the activity of substances participating in a reaction:
E = E0 - RT/nF ln(a p/ar)
where E is the electrochemical potential, E0 is the standard potential for the process, R is the ideal
gas constant, T is the absolute temperature, n is the number of electrons exchanged in the redox
reaction, F is Faraday’s constant, and ap and ar are the activities of the products of the reaction
and the reactants, respectively. The activity (a) of ions in solution is a measure of their molar
concentration (c in moles/L) under ideal conditions of infinite dilution. Expressing the activities
in terms of the product of molar concentrations and activity coefficients, γ (a measure of the
extent the ion deviates from ideal behavior in solution; thus a = γ · c, where γ #1), the Nernst
equation becomes:
E = E0 - RT/nF ln(γpcp/γrcr)
For dilute solutions of electrolytes (#10!2 molar), the activity coefficient is approximately one
(γ.1; it approaches one as the solution becomes more dilute, becoming one under ideal
conditions). Then, the Nernst equation is expressed in terms of the concentrations of ions in
solution, the typical form in which the equation is found in most chemistry textbooks (see also
Section 14.8.3.1, “Solubility and Solubility Product Constant,” for an application of activity to
the solubility product constant):
E = E0 - RT/nF ln(cp/cr)
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At concentrations less than 10!6 M, electrodeposition may show considerable deviations from
behavior of macroamounts of elements whose behavior partly depends on the nature and
previous treatment of the electrode (Adolff and Guillaumont, 1993). Inconsistent behavior is the
result of heterogeneity of the surface metal, a very important consideration when electrodepositing radionuclides at very low concentrations. The spontaneity predicted by the Nernst equation
for macroconcentrations of ions in solution at controlled potential is not always observed for
microconcentrations (Choppin et al., 1995). The activity of radionuclide ions is usually unknown
at low concentrations even if the concentration is known, because the activity coefficient (γ) is
dependent on the behavior of the mixed electrolytic system. In addition, the concentration might
not be accurately known because ions might adsorb on various surfaces, form complexes with
impurities, or precipitate on the electrode, for example. (See Section 14.9.3.7, “Oxidation and
Reduction,” for another application of the Nernst equation.) Separation is limited partly because
electrodeposition from very dilute solutions is slow, but it is also limited because it rarely leads
to complete separation of one element from many others (Coomber, 1975). Overall, the behavior
of an element during an electrochemical process is determined by its electrochemical potential,
which depends on the nature of the ion; its chemical form, its concentration, the general
composition of the electrolyte, the current density, material and design of the electrode, and
construction features of the electrochemical cell (Zolotov, 1990).
Often, trace elements are deposited on a solid cathode, but large separation factors between
micro- and macro-components are required. This condition is met when electrochemically active
metals are the main components or when the analyzed matrix does not contain macrocomponents that will separate on the cathode (Zolotov, 1990). Deposition of heavy metals and
actinides can be more difficult to control, for example, because of the decomposition of water
and reactions of cations and anions at electrodes (Adolff and Guillaumont, 1993). In some cases,
deposition of matrix components can be avoided by selection of a suitable medium and
composition of the electrolyte. Overall, the effectiveness of electrodeposition of trace
components depends on the electrode potential, electrode material and its working surface area,
duration of electrolysis, properties of the electrolyte (composition and viscosity), temperature,
and mixing rate (Zolotov, 1990). Even so, published data are empirical for the most part, and
conditions for qualitative reproducible separation are determined for each case. It is difficult,
therefore, to make general recommendations for selecting concentration conditions. It is
advisable to estimate and account for possible effects of different electrolysis factors when
developing separation or concentration methodologies (Zolotov, 1990).
14.6.2 Separation of Radionuclides
Although electrodeposition is not frequently used as a radiochemical separation technique,
several radionuclides [including iron (Hahn, 1945), cadmium (Wright, 1947), and technetium
(Flagg, 1945)] have been isolated by electrodeposition on a metal electrode. Electrodeposition is,
however, the standard separation technique for polonium, copper, and platinum. Polonium is
isolated through deposition on nickel from a strong hydrochloric acid (DOE, 1990 and 1997,
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Method Po-01). This separation is very specific, and, therefore, can be accomplished in the
presence of many other radionuclides. Electrodeposition at a mercury cathode has also been used
to separate technetium from fission products and for group separation of fission products
(Coomber, 1975). Numerous metals have been deposited on thin metal films by electrolysis with
a magnesium cathode. According to Coomber, “Electrodeposition of metals can be sensitive to
the presence of other substances” (Coomber, 1975). Deposition of polonium on silver is inhibited
by iron unless a reducing agent is present; and the presence of fluoride (F!1), trace amounts of
rare earths, can inhibit the deposition of americium. “In many cases the uncertainties of yield can
be corrected by the use of another radioisotope as an internal standard” (Coomber, 1975).
14.6.3 Preparation of Counting Sources
Electrodeposition is primarily used to prepare counting sources by depositing materials uniformly
in an extremely thin layer. Because of potential self-absorption effects, this approach is ideal for
the preparation of alpha sources. Numerous methods have been published for the electrodeposition of the heavy metals, e.g., the Mitchell method from hydrochloric acid (Mitchell,
1960), the Talvitie method from dilute ammonium sulfate [(NH4)2SO4] (Talvitie, 1972), and the
Kressin method from sodium sulfate-sodium bisulfate media (Kressin, 1977).
Sill and Williams (1981) and Hindman (1983, 1986) contend that coprecipitation is the preferred
method for preparation of sources for alpha spectrometry and that it should be assessed when
electrodeposition is being considered. Also see Section 14.8.4, “Coprecipitation.”
14.6.4 Advantages and Disadvantages of Electrodeposition
14.6.4.1

Advantages

• Highly selective in some cases.
• Deposits material in an extremely thin uniform layer resulting in excellent spectral resolution.
• One of the common methods for preparing actinides for alpha spectrometry.
14.6.4.2 Disadvantages
• Not applicable to many radionuclides.
• Sensitive to the presence of other substances.
• For tracer-level quantities, the process is relatively slow, it seldom leads to complete
separation of one element from many others, and there is usually no direct comparison of
concentration in solution to deposited activity.
• Takes longer than microprecipitation, because it requires evaporation of solutions after
column separation and ashing to remove all organic residue.
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• Subject to interference from such metals as Fe or Ti.
• Subject to interference from such ions as fluoride.

14.7 Chromatography
14.7.1 Chromatographic Principles
Chromatography is a separation technique that is based on the unequal distribution (partition) of
substances between two immiscible phases, one moving past the other. A mixture of the
substances (the analytical mixture) in the mobile phase passes over the immobile phase. Either
phase can be a solid, liquid, or gas, but the alternate phase cannot be in the same physical state.
The two most common phase pairs are liquid/solid and gas/liquid. Separation occurs as the
components in the mixture partition between the two phases because, in a properly designed
chromatographic system, the phases are chosen so that the distribution of the components
between the phases is not equal.
With the broad range of choices of phase materials, the number of techniques employed to
establish differential distributions of components between the phases, and the various practical
laboratory methods used to cause the mobile phases to pass over the immobile phases, there are
many chromatographic techniques available in separation chemistry. The names of the
chromatographic techniques themselves partially identify the methods or principles employed
and suggest the variety of applications available using this approach to separation. They include
paper chromatography, ion-exchange chromatography, adsorption chromatography, gas
chromatography, high-pressure liquid chromatography, and affinity chromatography. Each aspect
of chromatography used in separation chemistry will be described below, including the phases
commonly employed, the principles used to establish differential distributions, and the laboratory
techniques employed to run a chromatographic separation.
The most common phase pairs used in chromatography are a mobile liquid phase in contact with
a solid phase. The liquid phase can be a pure liquid, such as water or an organic solvent, or it can
be a solution, such as methyl alcohol, sodium chloride in water, or hexane in toluene. The solid
phase can be a continuous material such as paper, or a fine-grained solid such as silica, powdered
charcoal, or alumina. The fine-grained solid can also be applied to a supporting material, such as
paper, plastic, or glass, to form a coat of continuous material. Alternatively, gas/liquid phase
systems can consist of an inert gas, such as nitrogen or helium, in conjunction with a high-boiling
point liquid polymer coated on the surface of a fine-grained inert material, such as firebrick. This
system is called gas-liquid phase chromatography (GLPC), or simply gas chromatography (GC).
In each system, both phases play a role in the separation by offering a physical or chemical
characteristic that will result in differential distribution of the components of the analytical
mixture being separated. Liquid-liquid phase systems are similar to gas/liquid phase systems in
that one of the liquid phases is bound to an inert surface and remains stationary. These systems
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are often referred to as liquid-partition chromatography or liquid-phase chromatography (LPC),
because they are essentially liquid-liquid extraction systems with one mobile and one immobile
phase (Section 14.4, “Solvent Extraction”).
Differential distributions are established between the separating phases by the combination of
physical and chemical properties of the two phases in combination with those of the components
of the analytical mixture. The properties that are most commonly exploited by separation
chromatography are solubility, adsorption, ionic interactions, complementary interactions, and
selective inclusion. One or more of these properties is acting to cause the separation to occur.
14.7.2 Gas-Liquid and Liquid-Liquid Phase Chromatography
In gas-liquid phase chromatography, the components of the analytical mixture are first converted
to a vapor themselves and added to the flowing gas phase. They are then partitioned between the
carrier gas and liquid phases primarily by solubility differences of the components in the liquid
phase. As the gas-vapor mixture travels over the liquid phase, the more soluble components of
the mixture spend more time in the liquid. They travel more slowly through the chromatography
system and are separated from the less soluble, and therefore faster moving, components. Liquidliquid phase chromatography provides separation based on the same principle of solubility in the
two liquid phases, but the separation is performed at ambient temperatures with the components
of the analytical mixture initially dissolved in the mobile phase. Partitioning occurs between the
two phases as the mobile phase passes over the stationary liquid phase.
Gas chromatography has been used to concentrate tritium, and to separate krypton and xenon
fission products and fission-produced halogens (Coomber, 1975). A large number of volatile
metal compounds could be separated by gas chromatography, but few have been prepared.
Lanthanides and trivalent actinides have been separated on glass capillary columns using volatile
double halides formed with aluminum chloride (Coomber, 1975).
14.7.3 Adsorption Chromatography
Adsorption chromatography partitions components of a mixture by means of their different
adsorption characteristics onto the surface of a solid phase and their different solubilities in a
liquid phase. Adsorption phenomena are primarily based on intermolecular interactions between
the chemical components on the surface of the solid and the individual components of the
mixture. They include van der Waals forces, dipole-dipole interactions, and hydrogen bonds.
Silica is a useful adsorption medium because of the ability of its silyl OH groups to hydrogen
bond or form dipole-dipole interactions with molecules in the mixture. These forces compete
with similar intermolecular interactions—between the liquid phase and the components of the
mixture—to produce the differential distribution of the components. This process causes
separation to occur as the liquid phase passes over the solid phase.
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Many separations have been performed using paper and thin-layer chromatography. Modified
and treated papers have been used to separate the various valence states of technetium (Coomber,
1975).
14.7.4 Ion-Exchange Chromatography
14.7.4.1

Principles of Ion Exchange

Since the discovery by Adams and Holmes (1935) that synthetic resins can have ion-exchanging
properties, ion exchange has become one of the most popular, predominant, and useful techniques for radiochemical separations, both with and without carriers. There are many excellent
references available in the literature, e.g., Dean (1995), Dorfner (1972), Korkisch (1989), Rieman
and Walton (1970), and NAS monographs (listed in the references, under the author’s name).
The journal, Ion Exchange and Solvent Extraction, reports recent advances in this field of
separation.
Ion-exchange methods are based on the reversible exchange of metal ions between a liquid
phase, typically water, and a solid ionic phase of opposite charge, the resin. The resin competes
with the ion-solvent interactions in the liquid phase, primarily ion-dipole interactions and
hydrogen bonding, to produce the selective partition of ions, causing separation. The solid phase
consists of an insoluble, but permeable, inert polymeric matrix that contains fixed charged groups
(exchange sites) associated with mobile counter-ions of opposite charge. It is these counter-ions
that are exchanged for other ions in the liquid phase. Resins are either naturally occurring substances, such as zeolites (inorganic silicate polymers) or synthetic polymers. The synthetic resins
are organic polymers with groups containing the exchange sites. The exchange sites are acid or
base groups (amines, phenols, and carboxylic or sulfonic acids) used over a specific pH range
where they are in their ionic form. Typical exchange groups for cations (K+1, Ca+2, and UO2+2) are
the sulfonate anion, RSO3!1, or the carboxylate anion, RCOO!1. The quaternary-amine cation,
RNH3+1, or its derivative, is a common exchange group for anions (Cl !1, OH!1, and UO2(SO4)3!4).
In a practical description of ion-exchange equilibria, the weight distribution coefficient, Kd, and
the separation factor, α, are significant. The weight distribution coefficient is defined as:
Kd =

[C1 / g resin ]
[C 2 / mLsolution ]

where C1 is the weight of metal ion adsorbed on 1 g of the dry resin, and C2 is the weight of
metal that remains in 1 mL of solution after equilibrium has been reached. The separation factor
refers to the ratio of the distribution coefficients for two ions that were determined under
identical experimental conditions:

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Separation factor (α) =

[K d,a ]
[K d,b ]

where a and b refer to a pair of ions. This ratio determines the separability of the two ions;
separation will only be achieved if α … 1. The more that α deviates from unity, the easier it will
be to obtain separation.
An example of the separation process is the cation-exchange resin. It is usually prepared for
separation procedures as a hydrogen salt of the exchange group. Separation occurs when an
aqueous solution of other cation (e.g., Na+1, Ca+2, Al+3, or Cs+1) comes in contact with the resin. .
Different ions bond selectively to the exchange group, depending on the separation conditions,
displacing the counter-ion that is present in the prepared resin as follows:
ResinSO3!1 H+1 + Cs+1 6 ResinSO3!1Cs+1 + H+1
Diffusion is an important process during ion exchange; the solute ions must penetrate the pores
of the spherical resin beads to exchange with the existing ions. Equilibrium is established
between each ion in the analyte solution and the exchange site on the resin. The ion least tightly
bonded to the exchange site and most solvated in solution spends more time in solution. Selective bonding is a factor of the size and charge of the ion, the nature of the exchange group, and
the pH and ionic strength of the media. The order of strength of bonding at low acid concentrations for group 1 cations is H+1 or Li+1 < Na+1 < K+1 < Rb+1 < Cs+1 (Showsmith, 1984). Under the
appropriate conditions, for example, Cs+1 will bond exclusively, or Cs+1 and Rb+1 will bond,
leaving the remaining cations in solution. The process can be operated as a batch operation or via
continuous-flow with the resin in an ion-exchange column. In either case, actual separation is
achieved as the equilibrated solution elutes from the resin, leaving select ions bonded to the resin
and others in solution. The ion that spends more time in solution elutes first. The ability to “hold”
ionic material is the resin capacity, measured in units of mg or meq per gram of resin. Eventually,
most of the exchange groups are occupied by select ions. The resin is essentially saturated, and
additional cations cannot bond. In a continuous-flow process, breakthrough will then occur. At
this time, added quantities of select cations (Cs+1 or Cs+1 and Rb+1 in this example) will pass
through the ion-exchange column and appear in the output solution (eluate). No further separation can occur after breakthrough, and the bonded ions must be remove to prepare the column for
additional separation. The number of bed volumes of incoming solution (eluant) that passes
through a column resin before breakthrough occurs provides one relative measure of the treatment capacity of the resin under the conditions of column use. The bonded cations are displaced
by adjusting the pH of the medium to change the net charge on the exchange groups. This change
alters the ability of the exchange groups to attract ions, thereby replacing the bonded cations with
cations that bond more strongly. More commonly, the resin is treated with a more concentrated
solution of the counter-ion—H+1 in this example. Excess H+1 favors the equilibrium that produces
the initial counter-ion form of the exchange group. This process that returns the column to its
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original form is referred to as “regeneration.”
Overall, selectivity of the exchange resin determines the efficiency of adsorption of the analyte
from solution, the ease with which the ions can be subsequently removed from the resin, and the
degree to which two different ions of like charge can be separated from each other. The
equilibrium distribution of ions between the resin and solution depends on many factors, of
which the most important are the nature of the exchanging ions, the resin, and the solution:
• In dilute solutions, the stationary phase will show preference for ions of higher charge.
• The selectivity of ion exchangers for ions increases with the increase of atomic number
within the same periodic group, i.e., Li+ < Na+ < K+ < Rb+ < Cs+.
• The higher the polarizability and the lower the degree of solvation (favored by low charge
and large size), the more strongly an ion will be adsorbed.
• Resins containing weakly acidic and weakly basic groups are highly selective towards H+ and
OH! ions. Ion-exchange resins that contain groups capable of complex formation with
particular ions will be more selective towards those ions.
• As cross-linking is increased (see discussion of resins below), resins become more selective
in their behavior towards ions of different sizes.
• No variation in the eluent concentration will improve the separation for ions of the same
charge; however, for ions of different net charges, the separation does depend on the eluent
concentration.
14.7.4.2

Resins

The most popular ion-exchange resins are polystyrenes cross-linked through divinylbenzene
(DVB). The percentage of DVB present during polymerization controls the extent of crosslinking. Manufacturers indicate the degree of cross-linking by a number following an X, which
indicates the percentage of DVB used. For instance, AG 1-X8 and AG 1-X2 are 8 percent and 2
percent cross-linked resins, respectively. As this percentage is increased, the ionic groups effectively come into closer proximity, resulting in increased selectivity. However, increases in crosslinking decrease the diffusion rate in the resin particle. Because diffusion is the rate-controlling
step in column operations, intermediate cross-linking in the range of 4 to 8 percent is commonly
used.
Particle diameters of 0.04-0.3 mm (400 – 50 mesh) are commonly used, but larger particles give
higher flow rates. Difficult separations can require 200 – 400 mesh resins. Decreasing the particle
size reduces the time required for attaining equilibrium; but at the same time, it decreases flow
rate. When extremely small particle sizes are used, pressure must be applied to the system to
obtain acceptable flow rates (see discussion of high pressure liquid chromatography in Section
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14.7.7, “Chromatographic Methods”).
Ion-exchange resins are used in batch operations, or more commonly, in column processes in the
laboratory. Columns can be made in any size desired. The diameter of the column depends on the
amount of material to be processed, and the length of the column depends primarily on the
difficulty of separations to be accomplished. Generally, the ratio of column height to diameter
should be 8:1. Higher ratios lead to reduced flow rate; lower ratios might not provide effective
separations.
Some other factors should be considered when using ion-exchange resins:
• Resins should not be allowed to dry out, especially during analysis. Rehydration of dried
resins will result in cracking; these resins should not be used.
• Nonionic and weakly ionic solutes may be absorbed (not exchanged) by the resin. These
materials, if present during analysis, can alter the exchange characteristics of the resin for
certain ions.
• Particulate matter present in the analyte solution may be filtered by the resin. This material
will have several undesired effects, such as decreased flow rate, reduced capacity, and
ineffective separation.
• Organic solvents suspended in the analyte solution from previous separation steps can be
adsorbed by the resin creating separation problems.
Ion exchangers are classified as cationic or anionic (cation exchangers or anion exchangers,
respectively), according to their affinity for negative or positive counter-ions. They are further
subdivided into strongly or weakly ionized groups. Most cation exchangers (such as Dowex-50™
and Amberlite IR-100™) contain free sulfonic acid groups, whereas typical anion exchangers
(such as AG-1™ and Dowex-1™) have quaternary amine groups with replaceable hydroxyl ions
(Table 14.8).
TABLE 14.8 — Typical functional groups
of ion-exchange resins
Cation Exchangers

Anion Exchangers

- SO3H

- NH2

- COOH

- NHR

- OH

- NR2

- SH

- NR3+

R=alkyl group

The sulfonate resins are known as strong acid cation (SAC) resins because the anion is derived
from a strong sulfonic acid (RSO3H). Likewise, the carboxylate resins are known as weak acid
cation (WAC) resins because the anion is derived from a weak carboxylic acid (RCOOH). R in
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the formulas represents the inert matrix. The quaternary-amine cation (RNH3+1) or its derivatives,
represents the common exchange group for anions. Other functional groups can be used for
specific purposes.
Several examples from the literature illustrate the use of ion-exchange chromatography for the
separation of radionuclides. Radium is separated from other alkaline-earth cations (Be +2, Mg+2,
Ca+2, Sr+2, and Ba+2) in hydrochloric solutions on sulfonated polystyrene resins (Kirby and
Salutsky, 1964), or converted to an anionic complex with citrate or EDTA and separated on a
quaternary ammonium polystyrene resin (Sedlet, 1966).
Anion-exchange resins separate anions by an analogous process beginning with a prepared resin,
usually in the chloride form (RNH3+1Cl!1), and adding a solution of ions. Anion-exchange
chromatography is used in one step of a procedure to isolate thorium for radioanalysis by alpha
counting (EPA, 1984). Thorium cations (Th+4) form anionic nitrate complexes that bind to an
anion-exchange resin containing the quaternary complex, R-CH2-N(CH3)3+1. Most metal ion
impurities do not form the complex and, as cations, they do not bind to the exchanger, but remain
with the liquid phase. Once the impurities are removed, thorium itself is separated from the resin
by treatment with hydrochloric acid (HCl) that destroys the nitrate complex, leaving thorium in
its +4 state, which will not bind to the anionic exchanger. A selection of commercially available
resins commonly employed in the radiochemistry laboratory is given in Table 14.9.
TABLE 14.9 — Common ion-exchange resins (*)
Resin type &
nominal %
cross-link

Density
Minimum
(nominal)
wet
Description
g • mL!1
capacity
meq • mL!1
Anion-exchange resins — gel type — strongly basic — quaternary ammonium functionality
™
Dowex , AG™
1.0
0.70
Strongly basic anion exchanger with S-DVB matrix for separation
or Eichrom™
of organic acids, nucleotides, and other anions. Molecular weight
1- X 4
exclusion < 1400.
Dowex, AG or
1.2
0.75
Strongly basic anion exchanger with S-DVB matrix for separation
Eichrom
of inorganic and organic anions with molecular weight exclusion
1- X 8
< 1000. 100–200 mesh is standard for analytical separations.
Anion-exchange resins — gel type — intermediate basicity
1.1
0.70
Intermediate basic anion exchanger with primary tertiary amines
Bio-Rex™ 5
on an polyalkylene-amine matrix for separation of organic acids.
Anion-exchange resins — gel type — weakly basic — polyamine functionality
Dowex or AG
0.8
0.7
Weakly basic anion exchanger with tertiary amines on an acrylic
4- X 4
matrix. Suitable for use with high molecular weight organic
compounds.
™
1.6
1.06
Acrylic-DVB with unusually high capacity for large organic
Amberlite
IRA-68
molecules.
Cation-exchange resins - gel type - strongly acidic - sulfonic acid functionality

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Resin type &
nominal %
cross-link

Minimum
wet
capacity
meq • mL!1
1.1

Density
(nominal)
g • mL!1

1.7

0.80

Amberlite
IR-120

1.9

1.26

Duolite™ GT73
Amberlite
IRA-743A
Amberlite
IRC-718
Chelex® 100

1.3

Selective ion-exchange resins
1.30
Removal of Ag, Cd, Cu, Hg, and Pb.

0.6

1.05

Boron-specific.

1.0

1.14

Removal of transition metals.

0.4

0.65

Dowex, AG or
Eichrom
50W- X4
Dowex, AG or
Eichrom
50W- X8

0.80

Description
Strongly acidic cation exchanger with S-DVB matrix for
separation of amino acids, nucleosides and cations. Molecular
weight exclusion is < 1400.
Strongly acidic cation exchanger with S-DVB matrix for
separation of amino acids, metal cations, and cations. Molecular
weight exclusion is < 1000. 100–200 mesh is standard for
analytical applications.
8% styrene-DVB type; high physical stability.

Weakly acidic chelating resin with S-DVB matrix for heavy metal
concentration.
Chelating ion-exchange resin containing geminally substituted
Eichrom
diphosphonic groups chemically bonded to a styrenic-based
Diphonix®
polymer matrix. Extraordinarily strong affinity for actinides in the
tetra- and hexavalent oxidation states from highly acidic media.
Anion exchanger — macroreticular type — strongly basic — quaternary ammonium functionality
AG MP-1
1.0
0.70
Strongly basic macroporous anion exchanger with S-DVB matrix
for separation of some enzymes, and anions of radionuclides.
Cation-exchange resin — macroreticular type — sulfonic acid functionality
AG MP-50
1.5
0.80
Strongly acidic macroporous cation exchanger with S-DVB
matrix for separation of cations of radionuclides and other
applications.
Microcrystalline exchanger
AMP-1
4.0
Microcrystalline ammonium molybophosphate with cation
exchange capacity of 1.2 meq/g. Selectively exchanges larger
alkali-metal ions from smaller alkali-metal ions, particularly
cesium.
* Dowex is the trade name for Dow resins; AG and Bio-Rex are the trade names for Bio-Rad Laboratories resins;
Amberlite is the trade name of Rohm & Haas resins. MP is the acronym for macroporous resin; S-DVB is the
acronym for styrene-divinylbenzene.

The behavior of the elements on anion- and cation-exchange resins is summarized for several
resins in Faris and Buchanan (1964), Kraus and Nelson (1956), and Nelson et al. (1964). The
behavior in concentrated HCl is illustrated for cations on cation-exchange resins in Figure 14.3
(Dorfner, 1972) and for anions on anion-exchange resins in Figure 14.4 (Dorfner, 1972).
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FIGURE 14.3 — The behavior of elements in concentrated
hydrochloric acid on cation-exchange resins
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Figure 14.4 —The behavior of elements in concentrated
hydrochloric acid on anion-exchange resins
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14.7.5 Affinity Chromatography
Several newer types of chromatography are based on highly selective and specific attractive
forces that exist between groups chemically bound to an inert solid matrix (ligands) and molecular or ionic components of the analytical mixture. Affinity chromatography is an example of this
separation technique, which is used in biochemistry to isolate antigenic materials, such as
proteins. The proteins are attracted to their specific antibody that is bonded to a solid matrix.
These attractive forces are often called complementary interactions because they are based on a
lock-and-key type of fit between the two constituents. The interaction is complementary because
the two components match (fit) each other in size and electrical nature.
Crown ethers bonded to solid matrices serve as ligands in a chromatographic separation of
radium ions from aqueous solutions containing other cations (see Section 14.4.5.1, “Extraction
Chromatography Columns”). Even other alkaline-earth cations with the same +2 charge, such as
Sr+2 and Ba+2, offer little interference with radium binding because the cyclic nature of the crown
ether creates a ring structure with a cavity that complements the radius of the radium ion in
solution. In addition, the oxygen atoms of the cyclic ether are inside the ring, allowing these
electron-dense atoms to form effective ion-dipole interactions through water molecules with the
radium cation. Radionuclides analyzed by this method include 89/90Sr, 99Tc, 90Y, and 210Pb.
14.7.6 Gel-Filtration Chromatography
Another physical property that is used to separate molecules by a chromatographic procedure is
the effective size (molecular weight) of the molecule. High molecular-weight ions can also be
separated by this procedure. The method is known by several names, including gel-filtration
chromatography, molecular-sieve filtration, exclusion chromatography, and gel-permeation
chromatography. This technique is primarily limited to substances such as biomolecules with
molecular weights greater than 10,000 daltons (1.657 × 10-20 g). In similar types of solutions
(similar solutes and similar concentrations), the molecules or ions have a similar shape and
molecular weight that is approximately proportional to the hydrodynamic diameter (size) of the
molecule or ion. The solid phase consists of a small-grain inert resin that contains microscopic
pores in its matrix that will allow molecules and ions up to a certain diameter, called included
particles, to enter the resin. Larger particles are excluded. Of the included particles, the smaller
ones spend more time in the matrices. Separation of the molecules or ions is based on the fact
that those substances that are excluded are separated in a batch from the included substances,
while those that are included are separated by size. The log of the molecular weight of the
included molecules or ions is approximately inversely proportional to the time the particles spend
in the matrix.

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14.7.7 Chromatographic Laboratory Methods
Chromatographic separations are achieved using a variety of laboratory techniques. Some are
actually quite simple to perform, while others require sophisticated instrumentation. Paper
chromatography employs a solid-liquid phase system that separates molecules and ions with filter
paper or similar material in contact with a developing solvent. The analytical mixture in solution
is spotted at the bottom of the paper and allowed to dry, leaving the analytes on the paper. The
paper is suspended so that a small part of the bottom section is in a solvent, but not so deep that
the dry spots enter the solvent. By capillary action, the solvent travels up the paper. As the
solvent front moves up, the chromatogram is produced with the components of the mixture
partitioning between the liquid phase and the paper. Thin-layer chromatography is similar, but
the paper is replaced by a thin solid phase of separatory material (silica gel, alumina, cellulose,
etc.) coated on an inert support, such as plastic or glass.
Column chromatography can accommodate a larger quantity of both phases and can, therefore,
separate greater quantities of material by accepting larger loads or provide more separating power
with an increased quantity of solid phase. In the procedure, a solid phase is packed in a glass or
metal column and a liquid phase is passed through the column under pressure supplied by gravity
or low-pressure pumping action. For this reason, gravity flow (or pumping the liquid phase under
pressures similar to those generated by gravity flow) is often referred to as low-pressure chromatography. The liquid phase is usually referred to as the eluent and the column is eluted with the
liquid. Column chromatography is the common method used in ion-exchange chromatography.
With column chromatography, separation depends on: (1) type of ion-exchange resin used (i.e.,
cationic, anionic, strong, or weak); (2) eluting solution (its polarity affects ion solubility, ionic
strength affects displacement of separating ions, and pH affects net charge of exchange groups or
their degree of ionization in solution); (3) flow rate, grain size, and temperature, which affect
how closely equilibrium is approached (generally, low flow rate, small grain size, and high
temperature aid the approach to equilibrium and, therefore, increase the degree of separation);
and (4) column dimensions (larger diameter increases column capacity, while increased length
increases separation efficiency by increasing distance between ion bands as they travel through
the column) (Wahl and Bonner, 1951).
Metal columns can withstand considerably more pressure than glass columns. High-pressure
liquid chromatography (HPLC) employs stainless steel columns and solid phases designed to
withstand high pressures without collapsing. The method is noted for its rapid separation times
because of relatively high flow rates under high pressures (up to almost 14 MPa). For this reason,
the acronym HPLC alternatively represents high-performance liquid chromatography. HPLC is
often performed with a liquid-partition technique between an aqueous phase and organic phase,
but gel filtration, ion exchange, and adsorption methods are also employed. In the case of liquidpartition separations, either a stationary aqueous phase or stationary organic phase is selected.
The former system is referred to as normal phase chromatography and the latter as reversed phase
chromatography, a holdover from the first applications of the technique that employed a
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stationary aqueous phase. The aqueous phase is made stationary by adsorption onto a solid
support, commonly silica gel, cellulose powder, or polyacrylamide. An organic stationary phase
is made from particles of a polymer such as polyvinyl chloride or Teflon™. Reversed phase
HPLC has been used to separate individual elements of the lanthanides and actinides and
macroquantities of actinides (Choppin et al., 1995).
Gas/liquid phase systems are also used. During gas-liquid phase chromatography (GLPC—or
simply, gas chromatography [GC]), the gas phase flows over the liquid phase (coated onto an
inert solid) as an inert carrier gas—commonly helium or nitrogen—flows through the system at
low pressure. The carrier gas is supplied from a tank of the stored gas.
14.7.8 Advantages and Disadvantages of Chromatographic Systems
Ion-exchange chromatography is by far the predominant chromatographic method used for the
separation of radionuclides. Its advantages and disadvantages is presented exclusively in this
section.
14.7.8.1
•
•
•
•
•
•
•
•

Advantages

Highly selective.
Highly efficient as a preconcentration method.
Works as well with carrier-free tracer quantities as with weighable amounts.
Produces a high yield (recovery).
Can separate radionuclides from interfering counter-ions.
Simple process requiring simple equipment.
Wide scope of applications.
Can handle high volumes of sample.

14.7.8.2

Disadvantages

• May require high volume of eluent.
• Usually a relatively slow process, but rapid selective elution processes are known.
• Requires narrow pH control.

14.8 Precipitation and Coprecipitation
14.8.1 Introduction
Two of the most common and oldest methods for the separation and purification of ions in radioanalytical chemistry are precipitation and coprecipitation. Precipitation is used to isolate and
collect a specific radionuclide from other (foreign) ions in solution by forming an insoluble
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compound. Either the radionuclide is precipitated from solution itself, or the foreign ions are
precipitated, leaving the radionuclide in solution. Sometimes a radionuclide is present in solution
at sub-micro concentrations, i.e., levels so low that the radionuclide will not form an insoluble
compound upon addition of a counter-ion. In these cases, the radionuclide can often be brought
down from solution by coprecipitation, associating it with an insoluble substance that precipitates
from solution. This phenomenon is especially important in gravimetric analysis and radiochemistry. In gravimetric analysis, carrying down of impurities is a problem. For radiochemists,
coprecipitation is a valuable tool.
14.8.2 Solutions
Precipitation and coprecipitation provide an analytical method that is applied to ions in solution.
Solutions are simply homogeneous mixtures (a physical combination of substances), which can
be solids, liquids, or gases. The components of a solution consist of a solute and a solvent. The
solute is generally defined as the substance that is dissolved, and the solvent is the substance that
dissolves the solute. In an alternative definition, particularly suitable for liquid components when
it is not clear what is being dissolved or doing the dissolving, the solute is the minor constituent
and the solvent is the major constituent. In any event, the solute and solvent can consist of any
combinations of substances, so long as they are soluble in each other. However, in this chapter,
we are generally referring to aqueous solutions in which a solute is dissolved in water. The terms
below further describe solutions:
• Solubility is defined as the concentration of solute in solution that exists in equilibrium with
an excess of solute; it represents the maximum amount of solute that can dissolve in a given
amount of the solvent. The general solubilities of many of the major compounds of concern
are described in Table 14.10.
• An unsaturated solution is one in which the concentration of the solute is less than the
solubility. When additional solute is added to an unsaturated solution, it dissolves.
• A saturated solution is one that is in equilibrium with an excess of the solute. The
concentration of a saturated solution is equal to the solubility of the solute. When solute is
added to the saturated solution, no more solute dissolves.
• A supersaturated solution is a solution in which the concentration of solute is temporarily
greater than its solubility—an unstable condition. Therefore, when additional solute is added
to a supersaturated solution, solute comes out of solution as solid until the concentration
decreases to that of the saturated solution.

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TABLE 14.10 — General solubility behavior of some cations of interest (1)
The Common Cations
Na+1, K+1, NH4+1, Mg+2, Ca+2, Sr+2, Ba+2, Al+3, Cr+3, Mn+2, Fe+2, Fe+3,
Co+2, Ni+2, Cu+2, Zn+2, Ag+1, Cd+2, Sn+2, Hg2+2, Hg+2, and Pb+2
There are general rules of solubilities for the common cations found in most basic chemistry texts
(e.g., Pauling, 1970).
Under the class of mainly soluble substances:
• All nitrates (NO3!) are soluble.
• All acetates (C2H3O2!) are soluble.
• All chlorides (Cl!), bromides (Br!), and iodides (I!) are soluble, except for those of silver,
mercury, and lead. PbCl2 and PbBr2 are sparingly soluble in cold water, and more soluble in hot
water.
• All sulfates (SO4!2) are soluble, except those of barium, strontium, and lead. CaSO4, Ag2SO4, and
Hg2SO4 are sparingly soluble.
• Most salts of sodium (Na), potassium (K), and ammonium (NH4+) are soluble. Notable exceptions
are NaSb(OH)6, K3Co(NO2)6, K2PtCl6, (NH4)2PtCL6, and (NH4)3Co(NO2)6.
Under the class of mainly insoluble substances:
• All hydroxides (OH!1) are insoluble, except those of the alkali metals (Li, Na, K, Rb, and Cs),
ammonium, and barium (Ba). Ca(OH)2 and Sr(OH)2 are sparingly soluble.
• All normal carbonates (CO3!2) and phosphates (PO4!3) are insoluble, except those of the alkali
metals and ammonium. Many hydrogen carbonates and phosphates are soluble, i.e., Ca(HCO3)2,
Ca(H2PO4)2.
• All sulfides (S!2), except those of the alkali metals, ammonium, and the alkaline-earth metals (Be,
Mg, Ca, Sr, Ba, and Ra), are insoluble. Both aluminum- and chromium sulfide are hydrolyzed by
water, resulting in the precipitation of Al(OH)3 and Cr(OH)3.
• Some cations, such as Ba+2, Pb+2, and Ag+1, form insoluble chromates (CrO4!2), which can be used
as a basis for separation.
Actinide Elements
The solubility properties of the actinide M+3 ions are similar to those of the trivalent lanthanide ions,
while the behavior of the actinide M+4 ions closely resembles that of Ce+4.
• The fluorides (F!), oxalates (C2O4!2), hydroxides (OH!), and phosphates are insoluble.
• The nitrates, halides (except fluorides), sulfates, perchlorates (ClO4!), and sulfides are all soluble.
(1) Solubility data for specific compounds can be found in the CRC Handbook of Chemistry and Physics (CRC,
1999) and in the NAS-NS monographs.

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14.8.3 Precipitation
Precipitation is accomplished by combining a selected ion(s) in solution with a suitable counterion in sufficient concentrations to exceed the solubility of the resulting compound and produce a
supersaturated solution. Nucleation occurs and growth of the crystalline substance then proceeds
in an orderly manner to produce the precipitate (see Section 14.8.3.1, “Solubility and the
Solubility Product Constant, Ksp”). The precipitate is collected from the solvent by a physical
method, such as filtration or centrifugation. A cation (such as Sr+2, for example) will precipitate
from an aqueous solution in the presence of a carbonate anion, forming the insoluble compound,
strontium carbonate (SrCO3), when sufficient concentrations of each ion are present in solution
to exceed the solubility of SrCO3. The method is used to isolate and collect strontium from water
for radioanalysis (EPA, 1984).
A precipitation process should satisfy three main requirements:
• The targeted species should be precipitated quantitatively.
• The resulting precipitate should be in a form suitable for subsequent handling; it should be
easily filterable and should not creep.
• If it is used as part of a quantitative scheme, the precipitate should be pure or of known purity
at the time of weighing for gravimetric analysis.
Precipitation processes are useful in several different kinds of laboratory operations, particularly
gravimetric yield determinations—as a separation technique and for preconcentration—to
eliminate interfering ions, or for coprecipitation.
14.8.3.1

Solubility and the Solubility Product Constant, Ksp

Chemists routinely face challenges in the laboratory as a result of the phenomenon of solubility.
Examples include keeping a dissolved component in solution and coprecipitating a trace-level
analyte from solution.
Solubility equilibrium refers to the equilibrium that describes a solid (s) dissolving in solution
(soln), such as strontium carbonate dissolving in water, for example:
SrCO3(s) º Sr+2 (soln) + CO3!2 (soln)
or, alternately, a solid forming from solution, with the carbonate precipitating:
Sr+2 (soln) + CO3!2 (soln) 6 SrCO39 (s)
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The solubility product constant, Ksp, is the equilibrium constant for the former process, a solid
dissolving and forming ions in solution. Leussing (1959) explains K sp in general terms:
“For an electrolyte, MmNn, which dissolves and dissociates according to the equation:
MmNn(s) » MmNn(soln) » mM+n(soln.) + nN-m(soln)
“The equilibrium conditions exists that:
aMmNn(s) = aMmNn(soln) = amM+n(soln) · anN!m(soln.)
“[The value a is the activity of the ions in solution, a measure of the molar concentration
(moles/L) of an ion in solution under ideal conditions of infinite dilution.] (Also see Section
14.6.1, “Principles of Electrodeposition,” for a discussion of activity as applied to the Nernst
equation.) [This equation] results in the familiar solubility product expression since the
activity of a solid under given conditions is a constant. Expressing the activities in terms of
the product of molar concentrations and activity coefficients, γ [a measure of the extent the
ion deviates from ideal behavior in solution; thus a = γ · c where γ #1], [this] equation
becomes...
[M+n]m [N-m]n γmM+n γnN-m = a constant = Ksp ”
For dilute solutions of electrolytes (#10!2 molar), the activity coefficient is approximately one
(γ.1; it approaches one as the solution becomes more dilute, becoming one under the ideal
conditions of infinite dilution). Then, the solubility product constant is expressed in terms of the
concentrations of ions in solution, the typical form in which the equation is found in most
chemistry textbooks:
Ksp=[M+n]m [N-m]n
For strontium carbonate, Ksp is defined in terms of the concentrations of Sr+2 and CO3!2:
Ksp = [Sr+2][CO3!2] = 1.6×10!9
In order for the carbonate to precipitate, the product of the concentration of the ions in solution
representing the ions in the equilibrium expression, the common ions, must exceed the value of
the Ksp. The concentration of each common ion does not have to be equal. For example, if [Sr +2]
is 1×10!6 molar, then the carbonate ion concentration must be greater than 0.0016 molar for
precipitation to occur because (1×10!6) × (0.0016) = 1.6×10!9.
At higher concentrations ($10!2 molar), where the ions in solution deviate from ideal behavior,
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the value of the activity coefficient decreases, and the concentrations of the ions do not
approximate their activities. Under these conditions, the concentrations do not reflect the
behavior of the dissolution equilibrium, and the equation cannot be used for precipitation or
solubility calculations. More complex estimations of activity coefficients must be made and
applied to the general equation (Birkett et al., 1988). Generally, radiochemical separations use an
excess of a precipitating agent. The exact solution concentrations do not need to be known but
they should be high to ensure complete reaction. Practical radiochemical separations performed
based on solubility (either Ksp or coprecipitation phenomenon) are best described by Salutsky
(1959).
Analysts often need to know if a precipitate will form when two solutions are mixed. For
example:
“If a chemist mixes 100 mL of 0.0050 M NaCl with 200 mL of 0.020 M Pb(NO3)2, will lead
chloride precipitate? The ion product, Q, must be calculated and compared to Ksp for the
process:
PbCl2(s) º Pb+2(soln) + 2 Cl!(soln)
“After the two solutions are mixed, [Pb+2] = 1.3×10!2 M (0.2 L × 2.0×10!2 M/0.3 L), and
[Cl!] = 1.7×10!3 M (0.1 L × 5.0×10!3 M/0.3 L). The value for the ion product is calculated
from the expression
Q = [Pb+2] [Cl!]2 or [1.3×10!2] [1.7×10!3]2
Q = 3.8×10!8
“The numerical value for Ksp is 1.6×10!5. Because the ion product Q is less than Ksp, no
precipitate will form. Only when the ion product is greater than Ksp will a precipitate form.”
Conditions in the solution phase can affect solubility. For example, the solubility of an ion is
lower in an aqueous solution containing a common ion, one of the ions comprising the
compound, than in pure water because a precipitate will form if the Ksp is exceeded. This
phenomenon is known as the common ion effect and is consistent with LeChatelier’s Principle.
For example, the presence of soluble sodium carbonate (Na2CO3) in solution with strontium ions
can cause the precipitation of strontium carbonate, because carbonate ions from the sodium salt
contribute to their overall concentration in solution and tend to reverse the solubility equilibrium
of the “insoluble” strontium carbonate:
Na2CO3(s) º 2 Na+1(soln) + CO3!2(soln)
SrCO3(s) º Sr+2(soln) + CO3!2(soln)
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Alternatively, if a complexing agent or ligand is available that can react with the cation of a
precipitate, the solubility of the compound can be markedly enhanced. An example from Section
14.3.4.3, “Formation and Dissolution of Precipitates,” provides an illustration of this
phenomenon. In the determination of 90Sr, Sr+2 is separated from the bulk of the solution by direct
precipitation of the sulfate (SrSO4). The precipitate is redissolved by forming a complex ion with
EDTA, Sr(EDTA)!2, to separate it from lanthanides and actinides (DOE, 1997, Method RP520):
SrSO4(s) 6 Sr+2(soln) + SO4!2(soln)
Sr+2(soln) + EDTA!4 6 Sr(EDTA)!2(soln)
Additionally, many metal ions are weakly acidic and hydrolyze in solution. Hydrolysis of the
ferric ion (Fe+3) is a classical example of this phenomenon:
Fe+3 + H2O 6 Fe(OH)+2 + H+1
When these metal ions hydrolyze, producing a less soluble complex, the solubility of the salt is a
function of the pH of the solution, increasing as the pH decreases. The minimum solubility is
found under acidic conditions when the concentrations of the hydrolyzed species become
negligible. As demonstrated by Leussing, the solubility of a salt also depends upon the activity of
the solid phase. There are a number of factors that affect the activity of the solid phase (Leussing,
1959):
• Polymorphism is the existence of a chemical substance in two or more crystalline forms. For
example, calcium carbonate can have several different forms; only one form of a crystal is
stable at a given temperature. At ordinary pressures and temperatures, calcite with a solubility
of 0.028 g/L, is the stable form. Aragonite, another common form of calcium carbonate
(CaCO3), has a solubility of 0.041g/L at these conditions. It is not necessarily calcite that
precipitates when solutions of sodium carbonate and calcium nitrate are mixed. Extremely
low concentrations of large cations, such as strontium, barium, or lead, promote the
precipitation of aragonite over calcite (Wray and Daniels, 1957). On aging, the more soluble
aragonite converts to calcite.
• Various possible hydrates of a solid have different solubilities. For instance, at 25 EC, the
molar solubility of gypsum (CaSO4.2H2O) is 0.206 and that of anhydrite (CaSO4) is 0.271.
• The solid phase can undergo a reaction with a salt in solution.
• Particle size of a solid can affect its solubility. It has been demonstrated that the solubility of
smaller particles is greater than that of larger particles of the same material.
• Age of a precipitate can affect solubility. For example, Biederman and Schindler (1957) have
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demonstrated that the solubility of precipitated ferric hydroxide [Fe(OH) 3] undergoes a fourfold decrease to a steady state after 200 hours.
• Exchange of ions at the surface of the crystal with ions in the solution can affect the solubility
of a solid. This effect is a function of the amount of surface available for exchange and is,
therefore, greater for a finely divided solid. For example, Kolthoff and Sandell (1933)
observed that calcium oxalate (CaC2O4) can exchange with either sulfate or barium ions:
CaC2O4(s) + SO4!2(soln) 6 CaSO4(s) + C2O4!2(soln)
CaC2O4(s) + Ba+2(soln) 6 BaC2O4(s) + Ca+2(soln)
The excess of common ions that appears on the right-hand side of the equations represses the
solubility of calcium oxalate according to the laws of mass action.
Ideally, separation of common ions from foreign ions in solution by precipitation will result in a
pure solid that is easy to filter. This method should ensure the production of a precipitate to meet
these criteria as closely as possible. The physical process of the formation of a precipitate is quite
complex, and involves both nucleation and crystal growth. Nucleation is the formation within a
supersaturated solution of the smallest particles of a precipitate (nuclei) capable of spontaneous
growth. The importance of nucleation is summarized by Salutsky (1959):
“The nucleation processes govern the nature and purity of the resulting precipitates. If the
precipitation is carried out in such a manner as to produce numerous nuclei, precipitation will
be rapid, individual crystals will be small, filtration and washing difficult, and purity low. On
the other hand, if precipitation is carried out so that only a few nuclei are formed, precipitation will be slower, crystals larger, filtration easier, and purity higher. Hence, control of
nucleation processes is of considerable significance in analytical chemistry.”
Once the crystal nuclei are formed, crystal growth proceeds through diffusion of the ions to the
surface of the growing crystal and deposition of those ions on the surface. This crystal growth
continues until supersaturation of the precipitating material is eliminated and equilibrium
solubility is attained.
Thus, the goal is to produce fewer nuclei during precipitation so that the process will occur
slowly, within reasonable limits, and larger crystals will be formed. Impurities result from three
mechanisms: (1) inclusion, either by isomorphous replacement (isomorphic inclusion),
replacement of a common ion in the crystal structure by foreign ions of similar size and charge to
form a mixed crystal, or by solid solution formation (nonisomorphic inclusion), simultaneous
crystallization of two or more solids mixed together; (2) surface absorption of foreign ions; and
(3) occlusion, the subsequent entrapment of adsorbed ions as the crystal grows. Slow growth
gives the isomorphous ion time to be replaced by a common ion that fits the crystal structure
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perfectly, producing a more stable crystal. It also promotes establishment of equilibrium
conditions for the formation of the crystal structure so that adsorbed impurities are more likely to
desorb and be replaced by a common ion rather than becoming entrapped. In addition, for a given
weight of the solid that is forming, a small number of large crystals present an overall smaller
surface area than a large number of small crystals. The large crystals provide less surface area for
impurities to adsorb.
14.8.3.2

Factors Affecting Precipitation

Several factors affect the nature and purity of the crystals formed during precipitation. A
knowledge of these factors permits the selection and application of laboratory procedures that
increase the effectiveness of precipitation as a technique for the separation and purification of
ions, and for the formation of precipitates that are easily isolated. These factors, summarized
from Berg (1963) and Salutsky (1959), include the following:
• Rate of precipitation. Formation of large, well-shaped crystals is encouraged through slow
precipitation because fewer nuclei form and they have time to grow into larger crystals to the
detriment of smaller crystals present. Solubility of the larger crystals is less than that of
smaller crystals because smaller crystals expose more surface area to the solution. Larger
crystals also provide less surface area for the absorption of foreign ions. Slow precipitation
can be accomplished by adding a very dilute solution of the precipitant gradually, with
stirring, to a medium in which the resulting precipitate initially has a moderate solubility.
• Concentration of Ions and Solubility of Solids. The rate of precipitation depends on the
concentration of ions in solution and the solubility of the solids formed during the
equilibrium process. A solution containing a low concentration of ions, but sufficient
concentration to form a precipitate, will slow the process, resulting in larger crystal
formation. At the same time, increasing the solubility of the solid, either by selecting the
counter-ion for precipitation or by altering the precipitating conditions, will also slow
precipitation. Many radionuclides form insoluble solids with a variety of ions, and the choice
of precipitating agent will affect the solubility of the precipitate. For example, radium sulfate
(RaSO4) is the most insoluble radium compound known. Radium carbonate (RaCO3) is also
insoluble, but its Ksp is greater than that of radium sulfate (Kirby and Salutsky, 1964).
• Temperature. Precipitation at higher temperature slows nucleation and crystal growth
because of the increased thermal motion of the particles in solution. Therefore, larger crystals
form, reducing the amount of adsorption and occlusion. However, most solids are more
soluble at elevated temperatures, effectively reducing precipitate yield; an optimum
temperature balances these opposing factors.
• Digestion. Extremely small particles, with a radius on the order of one micron, are more
soluble than larger particles because of their larger surface area compared to their volume
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(weight). Therefore, when a precipitate is heated over time (digestion) the small crystals
dissolve and larger crystals grow (“Ostwald ripening”). Effectively, the small crystals are
recrystallized, allowing the escape of impurities (occluded ions) and growth of larger crystals.
This process reduces the surface area for adsorption of foreign ions and, at the same time,
replaces the impurities with common ions that properly “fit” the crystal lattice. Recrystallization perfects the crystal lattice, producing a purer precipitate (see Reprecipitation on page
14-68). Digestion is used in an 131I determination to increase the purity of the lead iodide
(PbI2) crystals (EPA, 1984).
• Degree of Supersaturation. A relatively high degree of supersaturation is required for
spontaneous nucleation, and degree of supersaturation is the main factor in determining the
physical character of a precipitate. Generally, the higher the supersaturation required, the
more likely a curdy, flocculated colloid will precipitate because more nuclei form under
conditions of higher supersaturation and crystal growth is faster. In contrast, the lower the
supersaturation required, the more likely a crystalline precipitate will form because fewer
nuclei form under these conditions and crystal growth is slower. Most perfect crystals are
formed, therefore, from supersaturated solutions that require lower ion concentrations to
reach the necessary degree of supersaturation and, as a result, inhibit the rate of nucleation
and crystal growth. Degree of supersaturation ultimately depends on physical properties of
the solid that affect its formation. Choice of counter-ion will determine the type of solid
formed from a radionuclide, which, in turn, determines the degree of saturation required for
precipitation. Many radionuclides form insoluble solids with a variety of ions, and the choice
of precipitating agent will affect the nature of the precipitate.
• Solvent. The nature of the solvent affects the solubility of an ionic solid (precipitate) in the
solvent. The polarity of water can be reduced by the addition of other miscible solvents such
as alcohols, thereby reducing the solubility of precipitates. Strontium chromate (SrCrO4) is
soluble in water, but it is insoluble in a methyl alcohol (CH3OH)-water mixture and can be
effectively precipitated from the solution (Berg, 1963). In some procedures, precipitation is
achieved by adding alcohol to an aqueous solution, but the dilution effect might reduce the
yield because it lowers the concentration of ions in solution.
• Ion Concentration. The common-ion effect causes precipitation to occur when the
concentration of ions exceeds the solubility-product constant. In some cases, however, excess
presence of common ions increases the solubility of the precipitate by decreasing the activity
of the ions in solution, as they become more concentrated in solution and deviate from ideal
behavior. An increase in concentration of the ions is necessary to reach the activity of ions
necessary for precipitate formation.
• Stirring. Stirring the solution during precipitation increases the motion of particles in solution
and decreases the localized buildup of concentration of ions by keeping the solution
thoroughly mixed. Both of these properties slow nucleation and crystal growth, thus
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promoting larger and purer crystals. This approach also promotes recrystallization because
the smaller crystals, with their net larger surface area, are more soluble under these
conditions. Virtually all radiochemical laboratories employ stirring with a magnetic stirrer
during precipitation reactions.
• Complex-Ion Formation. Formation of complex ions can be used to hold back impurities
from precipitating by producing a more soluble form of a solid. The classical example of this
phenomenon is the precipitation of lead (Pb+2) in the presence of silver ions (Ag+1). Chloride
ion (Cl!1) is the precipitating agent that produces insoluble lead chloride (PbCl2). In an excess
of the agent, silver chloride (AgCl) is not formed because a soluble salt containing the
complex ion, AgCl2!1 is formed. Complex-ion formation is also used to form precipitates (see
Section 14.3, “Complexation”).
• pH Effect. Altering the pH of aqueous solutions will alter the concentration of ions in the
precipitation equilibrium by the common-ion effect, if the hydrogen ion (H+1) or hydroxide
ion (OH!1) is common to the equilibrium. For example, calcium oxalate (CaC2O4) can be
precipitated or dissolved, depending on the pH of the solution, as follows:
Ca+2 + C2O4!2 6 CaC2O4
Because the oxalate concentration is affected by the hydrogen-ion concentration,
H+1 + C2O4!2 6 HC2O4!1,
increasing the hydrogen-ion concentration (lowering the pH) decreases the oxalate ion
concentration by forming bioxalate, which makes the precipitate more soluble. Therefore,
decreasing the hydrogen-ion concentration (raising the pH), therefore, aids precipitation.
Similar effects are obtained with carbonate precipitates:
Sr+2 + CO3!2 6 SrCO3
H+1 + CO3!2 6 HCO3!1
Many metal sulfides are formed in a solution of hydrogen sulfide by generating the sulfide
ion (S!2) at suitable pH:
H2S 6 H+1 + HS!1
HS!1 6 H+1 + S!2
Pb+2 + S!2 6 PbS
The pH can also influence selective formation of precipitates. Barium chromate will
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precipitate in the presence of strontium at pH 4 to 8, leaving strontium in solution. Sodium
carbonate is added and strontium precipitates after ammonia (NH3) is added to make the
solution more alkaline. This procedure is the basis for the separation of radium from
strontium in the radioanalysis of strontium in drinking water (EPA, 1980).
• Precipitation from Homogeneous Solution. Addition of a precipitating agent to a solution of
ions causes a localized excess of the reagent (higher concentrations) to form in the mixture.
The excess reagent is conducive to rapid formation of a large number of small crystals,
producing a precipitate of imperfect crystals that contains excessive impurities. The
precipitate formed under these conditions is sometimes voluminous and difficult to filter.
Localized excesses can also cause precipitation of more soluble solids than the expected
precipitate.
These problems largely can be avoided if the solution is homogenous in all stages of
precipitate formation, and if the concentration of precipitating agent is increased, as slowly as
practical, to cause precipitation from the most dilute solution possible. This increase in
concentration is accomplished, not by adding the precipitating agent directly to the solution,
but rather by generating the agent throughout the solution, starting with a very small concentration and slowly increasing the concentration while stirring. The precipitating agent is
generated indirectly as the result of a chemical change of a reagent that produces the precipitating agent internally and homogeneously throughout the solution. The degree of supersaturation is low because the concentration of precipitating agent in solution is always
uniformly low enough for nucleation only. This method produces larger crystals with fewer
impurities.
Table 14.11 (Salutsky, 1959) summarizes methods used for precipitate formation from
homogeneous solution. Descriptions of these methods can be found in Gordon et al. (1959).
Some agents are generated by decomposition of a compound in solution. Hydrogen sulfide,
for example, is produced from thioacetamide:
CH3CSNH2 + 2 H2O 6 CH3COO!1 + H2S + NH4+1
Copper sulfide (CuS) coprecipitates technetium from a homogeneous medium by the
generation of hydrogen sulfide by this method (EPA, 1973). Other agents alter the pH of the
solution (see “pH Effect” on the previous page). Hydrolysis of urea, for example, produces
ammonia, which raises the pH of a solution:
H2NCONH2 + H2O 6 CO2 + 2 NH3

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TABLE 14.11 — Summary of methods for utilizing precipitation
from homogeneous solution
Precipitant

Reagent

Element Precipitated

Hydroxide

Urea
Acetamide
Hexamethylenetetraamine
Metal chelate and H2O2
Triethyl phosphate
Trimethyl phosphate
Metaphosphoric acid
Urea
Dimethyl oxalate
Diethyl oxalate
Urea and an oxalate
Dimethyl sulfate
Sulfamic acid
Potassium methyl sulfate
Ammonium persulfate
Metal chelate and persulfate
Thiocetamide

Al, Ga, Th, Fe+3, Sn, and Zr
Ti
Th
Fe+3
Zr and Hf
Zr
Zr
Mg
Th, Ca, Am, Ac, and rare earths
Mg, Zn, and Ca
Ca
Ba, Ca, Sr, and Pb
Ba, Pb, and Ra
Ba, Pb, and Ra
Ba
Ba
Pb, Sb, Bi, Mo, Cu, and As, Cd, Sn, Hg,
and Mn
Th and Zr
Th and Fe+3

Phosphate

Oxalate

Sulfate

Sulfide
Iodate

Arsenate

Iodine and chlorate
Periodate and ethylene diacetate
(or ß-hydroxy acetate)
Ce+3 and bromate
Trichloroacetate
Urea and dichromate
Potassium cyanate and dichromate
Cr+3 and bromate
Acetamide
Silver ammonia complex
and ß-hydroxyethyl acetate
Arsenite and nitrite

Zr

Tetrachlorophthalate

Tetrachlorophthalic acid

Th

Dimethylglyoxime
8-Hydroxyquinoline
Fluoride

Urea and metal chelate
Urea and metal chelate
Fluoroboric acid

Ni
Al
La

Carbonate
Chromate

Periodate
Chloride

Ce+4
Rare earths, Ba, and Ra
Ba and Ra
Ba, Ra
Pb
Pb
Ag

Source: Salutsky, 1959.

• Reprecipitation. This approach increases the purity of precipitates. During the initial
precipitation, crystals collected contain only a small amount of foreign ions relative to the
common ions of the crystal. When the precipitate is redissolved in pure solvent, the foreign
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ions are released into solution, producing a concentration of impurities much lower than that
in the original precipitating solution. On reprecipitation, a small fraction of impurities is
carried down with the precipitate, but the relative amount is much less than the original
because their concentration in solution is less. Nevertheless, foreign ions are not eliminated
because absorption is greater at lower, rather than at higher, concentrations. On balance,
reprecipitation increases the purity of the crystals. Reprecipitation is used in the procedure to
determine Am in soil (DOE, 1990 and 1997, Method Am-01). After americium is coprecipitated with calcium oxalate (CaC2O4), the precipitate is reprecipitated to purify the solid.
14.8.3.3

Optimum Precipitation Conditions

There is no single, fixed rule to eliminate all impurities during precipitation (as discussed in the
section above), but over the years, a number of conditions have been identified from practical
experience and theoretical considerations that limit these impurities (Table 14.12). Precipitations
are generally carried out from dilute solutions adding the precipitant slowly with some form of
agitation to a hot solution. Normally, the precipitant is then allowed to age before it is removed
by filtration and washed. Reprecipitation is then commonly performed. Reprecipitation is one of
the most powerful techniques available to the analyst because it increases purity, regardless of the
form of the impurity. Table 14.12 highlights the optimum precipitation conditions to eliminate
impurities.
TABLE 14.12 — Influence of precipitation conditions on the purity of precipitates
Form of Impurity*
Condition

Mixed
Crystals

Surface
Adsorption

Occlusion and
Inclusion

Postprecipitation

Dilute solutions

"

+

+

"

Slow precipitation

+

+

+

-

Prolonged digestion

-

+

+

-

High temperature

-

+

+

-

Agitation

+

+

+

"

Washing the precipitate

"

+

"

"

Reprecipitation
+
+
+
*Symbols: +, increased purity; -, decreased purity; ", little or no change in purity
Source: Salutsky, 1959.

"

14.8.4 Coprecipitation
In many solutions, especially those of environmental samples, the concentration of the radionuclide of interest is too low to cause precipitation, even in the presence of high concentrations of its
counter-ion, because the product of the concentrations does not exceed the solubility product.
Radium in most environmental samples, for example, is not present in sufficient concentration to
cause its very insoluble sulfate (RaSO4) to precipitate. The radionuclide can often be brought
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down selectively and quantitatively from solution during precipitation of an alternate insoluble
compound by a process called coprecipitation. The insoluble compound commonly used to
coprecipitate radium isotopes in many radioanalytical procedures is another insoluble sulfate,
BaSO4 (EPA, 1984, Method Ra-01; EPA, 1980, Method 900.1). The salt is formed with barium,
also a member of the alkaline earth family of elements with chemical properties very similar to
those of radium. Alternatively, a different salt that is soluble for the radionuclide can be used to
cause coprecipitation. Radium can be coprecipitated with lanthanum fluoride, even though
radium fluoride is soluble itself. For trace amounts of some radionuclides, other isotopic forms of
the element are available that can be added to the solution to bring the total concentration of all
forms of the element to the level that will result in precipitation. For example, to determine 90Sr
in environmental samples, stable strontium (containing no radioisotopes of strontium) is added to
increase the concentration of total strontium to the point that the common ion effect causes
precipitation. The added ion that is present in sufficient concentration to cause a precipitate to
form is called a carrier (Section 14.9, “Carriers and Tracers”). Barium, lanthanum, and stable
strontium, respectively, are carriers in these examples (DOE, 1997, Method RP5001; DOE, 1990
and 1997, Method Sr-02; EPA, 1984, Sr-04). The term carrier is also used to designate the
insoluble compound that causes coprecipitation. Barium sulfate, lanthanum fluoride (LaF3), and
strontium carbonate are sometimes referred to as the carrier in these coprecipitation procedures.
See Wahl and Bonner (1951) for additional examples of tracers and their carriers used for
coprecipitation.
The common definition of coprecipitation is, “the contamination of a precipitate by substances
that are normally soluble under the conditions of precipitation” (Salutsky, 1959). In a very broad
sense, coprecipitation is alternately defined as the precipitation of one compound simultaneously
with one or more other compounds to form mixed crystals (Berg, 1963). Each is present in macro
concentrations (i.e., sufficient concentrations to exceed the solubility product of each). As the
term is used in radiochemistry, coprecipitation is the simultaneous precipitation of one
compound that is normally soluble under the conditions of precipitation with one or more other
compounds that form a precipitate under the same conditions. Coprecipitation of two or more
rare earths as oxalates, barium and radium as sulfates, or zirconium and hafnium as phosphates
are examples of this broader definition (Salutsky, 1959). By either definition, coprecipitation
introduces foreign ions into a precipitate as impurities that would normally be expected to remain
in solution; and precipitation techniques, described in the previous section, are normally used to
maximize this effect while minimizing the introduction of true impurities. As a method to
separate and collect radionuclides present in solution at very low concentration, coprecipitation is
performed in a controlled process to associate the ion of choice selectively with a precipitate,
while excluding other foreign ions that would interfere with the analytical procedure.
14.8.4.1

Coprecipitation Processes

In order to choose the best conditions to coprecipitate an ion selectively, two processes should be
considered. First is precipitation itself and the appropriate techniques employed to minimize
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association of impurities (see Section 14.8.3). Second is coprecipitation mechanisms and the
controlling factors associated with each. Three processes (described above in Section 14.8.3.1,
“Solubility and the Solubility Product Constant”) are responsible for coprecipitation, although
the distinction between these processes is not always clear (Hermann and Suttle, 1961). They
consist of: (1) inclusion, i.e., uptake from solution of an ion similar in size and charge to the solid
forming the precipitate in order to form a mixed crystal or solid solution; (2) surface adsorption;
and (3) occlusion (mechanical entrapment).
Inclusion. If coprecipitation is accomplished from a homogeneous solution allowing the crystals
to form slowly in an orderly manner, then inclusion contributes to the coprecipitation process.
Under these conditions, the logarithmic distribution law applies, which represents the most
efficient coprecipitation method that involves mixed crystals (Salutsky, 1959):
log(Ii/If) = λ log(Pi/Pf)
In the equation, I i is the concentration of impurity in solution at the start of crystallization and I f
is the concentration at the end. P represents the corresponding concentration of the primary ion in
solution. Lambda, λ, is the logarithmic distribution coefficient and is a constant. Values of λ for
some tracers distributed in solid carriers can be found in Wahl and Bonner (1951). Lambda
values greater than one represent removal of a foreign ion by inclusion during coprecipitation.
The larger the value of lambda, the more effective and selective the process for a specific ion.
Lambda is also inversely proportional to the rate of precipitation. Slow precipitation, as
accomplished by homogeneous precipitation, results in larger values and more efficient
coprecipitation. For example, “Actinium [Ac] has been selectively removed from solutions
containing iron and aluminum [Al] through slow oxalate precipitation by the controlled
hydrolysis of dimethyl oxalate” (Hermann and Suttle, 1961). Also, as described in Section
14.8.3.2, “Factors Affecting Precipitation,” technetium is coprecipitated with copper sulfide
(CuS) carrier produced by the slow generation of hydrogen sulfide (H2S) as thioacetamide is
hydrolyzed in water (EPA, 1973).
Generally, λ decreases as the temperature increases; thus, coprecipitation by inclusion is favored
by lower temperature.
Digestion of the precipitate at elevated temperature over lengthy time periods—a process that
promotes recrystallization and purer crystals—will often cause mixed crystals to form by an
alternate mechanism (i.e., homogeneous distribution) that is not as efficient, but which is often as
successful as logarithmic distribution. The equilibrium distribution law is represented by
(Salutsky, 1959):
(I/P)ppt. = D (I/P)soln.
where I represents the amount of impurity and P the amount of primary substance forming the
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precipitate. The symbol D is the homogeneous distribution coefficient. Values of D greater than
one represent removal of a foreign ion by inclusion during coprecipitation. Some values of D can
be found in Wahl and Bonner (1951). According to Hermann and Suttle (1961):
“Homogeneous distribution is conveniently obtained at ordinary temperatures by rapid
crystallization from supersaturated solutions with vigorous stirring. Under such conditions
the precipitate first formed is very finely divided, the recrystallization of the minute crystals
is rapid, and each molecule [sic] passes many times between solution and precipitate. If this
process is repeated often enough, an equilibrium between solid and solution is obtained, and
all the resulting crystals grow from a solution of constant composition.”
In either case, optimal results are obtained through inclusion when the precipitate contains an ion
with chemical properties similar to those of the foreign ion, although it is not necessary for the
similarity to exist in every successful coprecipitation. Barium sulfate is very successful in
coprecipitating Ra+2, primarily because radium is in the same chemical family as barium, and has
the same charge and a similar ionic radius. For best results, the radius of the foreign ion should
be within approximately 15 percent of that of one of the common ions in the precipitate
(Hermann and Suttle, 1961).
Surface Adsorption. During surface adsorption, ions are adsorbed from solution onto the surfaces
of precipitated particles. The conditions leading to surface adsorption are described by Salutsky
(1959):
“The surface of a precipitate is particularly active. Ions at the surface of a crystal (unlike
those within the crystal) are incompletely coordinated and, hence are free to attract other ions
of opposite charge from solution.”
Adsorption involves a primary adsorption layer that is held very tightly, and a counter-ion layer
held more loosely. Ions common to the precipitate are adsorbed most strongly at the surface to
continue growth of the crystal. During precipitation of BaSO4, barium ions (Ba+2) and sulfate ions
(SO4!2) are the primary ions adsorbed. If only one of the common ions remains in solution, then
foreign ions of the opposite charge are adsorbed to maintain electrical neutrality. When barium
sulfate is precipitated from a solution containing excess barium ions, for example, foreign ions
such as Cl!1, if present, are adsorbed after sulfate ions are depleted in the precipitation process.
Foreign ions of the same charge, such as Na+1, are repelled from the surface. Surface adsorption
can be controlled, therefore, by controlling the concentration of ions during precipitation or by
the addition of ions to alter the concentration. A precipitate of silver chloride (AgCl) in excess
Ag+1 repels 212Pb+2, but in a solution containing an equal quantity of the common silver and
chloride ions, approximately 2 percent of 212Pb is adsorbed (Salutsky, 1959). In contrast, almost
86 percent of 212Pb is adsorbed if an iodide solution is added to precipitate the silver ions as silver
iodide (AgI), thereby reducing the concentration of silver ions and making the chloride ion in
excess in the solution. According to the Paneth-Fajans-Hahn adsorption rule, the ion most
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adsorbed will be the one that forms the least soluble compound with an ion of the precipitate. For
example, barium sulfate in contact with a solution containing excess sulfate ions will adsorb ions
of Pb > Ca > K > Na, which reflects the order of solubility of the respective sulfates: thus, PbSO 4
< CaSO4 < K2SO4 < Na2SO4 (Salutsky, 1959).
“Because adsorption is a surface phenomenon, the larger the surface area of a precipitate, the
greater the adsorption of impurities” (Salutsky, 1959). For that reason, colloidal crystals
exhibit a high degree of nonspecific adsorption. When a colloid is flocculated by the addition
of an electrolyte, the electrolyte can be adsorbed as an impurity. This interference largely can
be eliminated by aging the precipitate, thereby growing larger crystals and reducing the
surface area. Additionally, nonvolatile impurities can be replaced on the particle by washing
the colloidal precipitate with a dilute acid or ammonium salt solution. Well-formed large
crystals exhibit much less adsorption, and adsorption is not a significant factor in
coprecipitation with these solids. The tendency for a particular ion to be adsorbed depends
on, among other factors, charge and ionic size (Berg, 1963). Large ions with a high charge
exhibit high adsorption characteristics: a high ionic charge increases the electrostatic
attraction to the charged surface, and an ion with a large radius is less hydrated by the
solution and not as attracted to the solution phase.
“The amount of adsorption is also affected by prolonged standing of the precipitate in contact
with the solution. The fraction adsorbed is higher for some tracer ions, while the fraction is
lower for others. Recrystallization occurring during standing decreases the surface area so
that the fraction of tracer carried will decrease unless the tracer is trapped in the growing
crystals ... in which case the fraction carried may increase (Wahl, 1951).”
Adsorption also depends on the concentration of an ion in solution (Berg, 1963). A high
concentration of impurity increases the probability of solute interaction at the solid surface and
favors adsorption. Salutsky (1959) comments on the percent adsorption:
“Generally, the percent adsorption is much greater at low concentrations than at high
concentrations. At very high concentrations of impurity, adsorption reaches a maximum
value, i.e., the adsorption is saturated.”
Occlusion. Occlusion of an impurity within a precipitate results when the impurity is trapped
mechanically by subsequent crystal layers. For that reason, occluded impurities cannot be
physically removed by washing. Occlusion is more prevalent with colloidal precipitates than with
large crystals because of the greater surface area of colloidal solids. Freshly prepared hydroxides
and sulfides commonly contain occluded impurities, but most of them are released upon aging of
the precipitate.
Mechanical entrapment occurs particularly when the precipitating agent is added directly to a
solution. Because of the localized high concentrations of precipitant, impurities are precipitated
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that become occluded by the subsequent precipitation of the primary substance. The speed of the
precipitation process also affects the extent of occlusion. Occlusion can be reduced, therefore, by
homogeneous precipitation. Coprecipitation of strontium by barium sulfate, for example, is
accomplished by the homogeneous generation of sulfate by the hydrolysis of dimethylsulfate,
(CH3)2SO4 (Hermann and Suttle, 1961). Digestion also eliminates occluded particles as the solid
is recrystallized. Considerable occlusion occurs during nucleation, and, therefore, reducing the
precipitation rate by lowering the temperature and reducing the number of nuclei formed reduces
the initial coprecipitation by occlusion.
This type of coprecipitation is not limited to solid impurities. Sometimes the solvent and other
impurities dissolved in the solvent become trapped between layers of crystals. This liquid
occlusion is common in numbers of minerals such as quartz and gypsum.
14.8.4.2

Water as an Impurity

In addition to other impurities, all precipitates formed from aqueous solutions contain water
(Salutsky, 1959). This water might be essential water, present as an essential part of the chemical
composition (e.g., MgNH4PO4 @ 6H2O, Na2CO3 @ H2O), or it might be nonessential water.
Nonessential water can be present in the precipitate as hygroscopic water, surface water, or
included water. Hygroscopic water refers to the water that a solid adsorbs from the surrounding
atmosphere. Many colloidal precipitates are highly hygroscopic because of their large surface
areas. Moreover, water can be adsorbed to the surface of the precipitate or included within the
crystal matrix, as described previously.
14.8.4.3

Postprecipitation

Postprecipitation results when a solution contains two ions, one that is rapidly precipitated and
another that is slowly precipitated by the precipitating agent (Kolthoff et al., 1969). The first
precipitate is usually contaminated by the second one. For example, calcium oxalate is a
moderately insoluble compound that can be precipitated quantitatively with time. Because the
precipitation tends to be slow, the precipitate is allowed to remain in contact with the solution for
some time before filtering. Magnesium oxalate is too soluble to precipitate on its own under
normal conditions. As long as the solution contains a predominance of calcium ions, very little
magnesium precipitates. However, as the precipitation of calcium approaches quantitative levels,
the competition of calcium and magnesium ions for adsorption at the surface becomes more
intense. As time progresses, the magnesium oxalate adsorbed on the surface acts as seed to
induce the post-precipitation of a second solid phase of magnesium oxalate (MgC2O4). Once
precipitated, the magnesium oxalate is only slightly soluble and does not redissolve.

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14.8.4.4

Coprecipitation Methods

Selective coprecipitation of a radionuclide with an insoluble compound is primarily accomplished by the judicious selection of the compound that forms the precipitate and the concentration
of solutions used in the precipitate’s formation. Using good precipitation technique minimizes
the coprecipitation of impurities. The compound, then, should maximize coprecipitation of the
select radionuclide while providing a well-formed solid that attracts a minimum of other foreign
ions as impurities. In general, conditions that favor precipitation of a substance in macroamounts
also favor the coprecipitation of the same material from tracer concentrations (i.e., too low for
precipitate formation) with a foreign substance (Friedlander et al., 1981). Wahl and Bonner
(1951) provide a useful summary for coprecipitation of a tracer by a carrier:
“In general a tracer is efficiently carried by an ionic precipitate if: (1) the tracer ion is
isomorphously incorporated into the precipitate, or (2) the tracer ion forms a slightly soluble
or slightly dissociated compound with the oppositely charged lattice ion and if the precipitate
has a large surface with charge opposite to that of the tracer ion (i.e., presence of excess of
the oppositely charged lattice ion).”
Considering the principles of precipitation and coprecipitation, radium is coprecipitated quantitatively with barium sulfate using excess sulfate in solution because: (1) radium forms the least
soluble sulfate of the other elements in the alkaline earth family (Paneth-Fajans-Hahn adsorption
rule); (2) the radium ion carries the same charge as the barium ion and is very similar in size
(inclusion); and 3) an excess of sulfate preferentially creates a common-ion layer on the crystalline solid of sulfate ions that attracts barium ions and similar ions such as radium (absorption).
For example, in a procedure to determine 226Ra in water samples, radium is coprecipitated as
barium sulfate using 0.36 moles of sulfate with 0.0043 moles of barium, a large excess of sulfate
(EPA, 1984, Method Ra-03).
The isolation of tracers often occurs in two steps: first the tracer is separated by coprecipitation
with a carrier, and then it is separated from the carrier (Hermann and Suttle, 1961). Use of
carriers that can be easily separated from the tracer is helpful, therefore, coprecipitation by
inclusion is not generally used. Coprecipitation by surface adsorption on unspecific carriers is the
most common method employed. Manganese dioxide MnO2, sulfides (MnS), and hydroxides
[Mn(OH)2] are important nonspecific carriers because of their high surface areas. Ferric
hydroxide [Fe(OH)3] is very useful for adsorbing cations, because it forms a very finely divided
precipitate with a negative charge in excess hydroxide ion. Ferric hydroxide is used, for example,
to collect plutonium in solution after it has been isolated from tissue (DOE, 1990 and 1997,
Method Pu-04). Tracers can be separated by dissolving the solid in acid and extracting the iron in
ether (Hermann and Suttle, 1961).
“The amount of ion adsorbed depends on its ability to compete with other ions in solution.
Ions capable of displacing the ions of the radioelements are referred to as holdback carriers
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[see Section 14.9.2.4, “Holdback Carriers”]. Highly charged ions, chemical homologs, and
ions isotopic with the radioelement are among the most efficient displacers. Thus, the
addition of a little inactive strontium makes it possible to precipitate radiochemically pure
radiobarium as the nitrate or chloride in the presence of radiostrontium.”
Tables 14.13 and 14.14 provide more details about common coprecipitating agents for
radionuclides.
TABLE 14.13 — Common coprecipitating agents for radionuclides(1)
Radionuclide
Am

Oxidation
State
+3

Cs

+1

Co

+2

Fe

+3

I
Ni
Nb
Np
Po

!1
+2
(V)
+4
+4

Pu

+3
+4
(VI)

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Coprecipitate
hydroxide
iodate
fluoride, oxalate, phosphate,
hydroxide
oxalate
acetate
fluoride, sulfate
acetate
phosphomolybdate,
chloroplatinate, bismuth
nitrate, silicomolybdate
hydroxide
potassium cobalt nitrate
1-nitroso-2-napthol
sulfide
hydroxide
ammonium pyrouranate
iodide
dimethylglyoxime hydroxide
hydroxide, phosphate
phosphate
tellurium
tellurate
selenium
dioxide
hydroxide
sulfide
fluoride
sulfate
fluoride
oxalate, iodate
phosphate
sodium uranylacetate

14-76

Carrier(2)
Am+3, Fe+3
Ce+4, Th+4, Zr+4
La+3, Ce+3, Nd+3, Bi+3

Notes

Ca+2
Am+4
La+3
UO2+2
Cs+1

Co+2
Co+2
Co+2
Co+2
Fe+3
Fe+3
+2
+1
Pb , Ag , Pd+2, Cu+2
Ni+2
Nb(V)
Ca+2
Te
Pb+2
Se or Se!2
Mn+4
+3
Fe , Al+3, La+3
Cu+2, Bi+2, Pb+2
La+3, Nd+3, Ce+3, Ca+2
La+3(K+1)
La+3, Nd+3, Ce+3
Th+4
+2
Zr , Bi+3
UO2+2

Tellurate reduced with
SnCl2

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Radionuclide
Ra

Sr

Tc

Oxidation
State
+2

+2

+4
(VII)

Th

+4

U

+4

(V)
(VI)

Coprecipitate
hydroxide
sulfate, chromate, chloride,
bromide
oxalate, phosphate
fluoride
carbonate
nitrate
chromate
sulfate
phosphate
hydroxide
hydroxide
chlorate, iodate,
perruthenate,
tetrafluoroborate
sulfide
hydroxide
fluoride
iodate
phosphate, peroxide
sulfate
oxalate
cupferron, pyrophosphate,
phosphate, iodate, sulfate,
oxalate
fluoride
phosphate
sulfate
cupferron
pyrouranate

Carrier(2)
Fe+3
Ba+2
Th+4, Ca+2, Ba+2
La+3
+2
Sr , Ba+2, Ca+2
Sr+2, Ba+2
Ba+2
+2
Sr , Ca+2, Pb+2
Sr+2
Fe+3
+4
Tc , Fe+3, Mn+2
(Phenyl)4As+1

Alkaline pH

Tc+7, Re+7, Cu+2, Cd+2
Th+4, La+3, Fe+3, Zr+3,
Ac+3, Zn+2
+4
Th , La+3, Nd+3, Ce+3
Th+4, Zr+3
Th+4, Bi+3
Ba+2
Ca+2
U+4

La+3, Nd+3
Zr+3
Ca+2
U(VI)
U(VI)

U(VI), Al+3
U(VI)

phosphate
peroxide

Notes

Neutral solution
From aqueous NH3, many
ions stay in solution as
NH3 complex
Th+4, Zr+3 also
coprecipitate
Without carbonate

hydroxide
Fe+3
fluoride
Th+4
Zr
+4
hydroxide
Fe+3
(1) Compiled from: Anders, 1960; Booman and Rein, 1962; Cobble, 1964; EPA, 1973; 1980; 1984; DOE, 1990,
1995, 1997; Finston and Kinsley, 1961, Grimaldi, 1961; Grindler, 1962; Hyde, 1960; Kallmann, 1961;
Kallmann, 1964; Kirby and Salutsky, 1964; Metz and Waterbury, 1962; Sedlet, 1964; Sundermann and
Townley, 1960; and Turekian and Bolter, 1966.
(2) If the radionuclide itself is listed as the carrier, a different isotope would be used to assess recovery.

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TABLE 14.14 — Coprecipitation behavior of plutonium and neptunium
Carrier Compound
Pu+3
Pu+4
Pu(VI)
Np+4
Np(V)
Np(VI)
Hydroxides
C
C
C
C
C
C
Calcium fluoride
C
C
C
Lanthanum fluoride
C
C
NC
C
C
NC
Barium sulfate
C
C
NC
C
NC
NC
Phosphates:
Calcium phosphate
C
C
C
Bismuth phosphate
C
C
C
NC
NC
Zirconium phosphate
NC
C
NC
C
NC
NC
Thorium pyrophosphate
NC
C
NC
Thorium hypophosphate
C
NC
C
NC
U+4 hypophosphate
Oxalates:
Lanthanum oxalate
C
C
NC
NC
Bismuth oxalate
C
C
NC
Thorium oxalate
C
C
NC
C
C
C
NC
U+4 oxalate
Iodates:
Zirconium iodate
C
NC
C
Ceric iodate
C
NC
C
Thorium iodate
C
NC
C
NC
Sodium uranyl acetate
NC
NC
C
NC
Poor
C
Zirconium phenylarsenate
NC
C
NC
C
Poor
NC
Thorium peroxide
C
C
Bismuth arsenate
C
NC
C
“C” indicates nearly quantitative coprecipitation under proper conditions; “NC” indicates that
coprecipitation can be made less than 1–2 percent under proper conditions. [Data compiled from
Seaborg and Katz, Korkisch (1969), and the NAS-NS 3050, 3058 and 3060 monographs.]

14.8.5 Colloidal Precipitates
Many precipitates exhibit colloidal properties, especially when freshly formed (Salutsky, 1959).
The term “colloid state” refers to the dispersion of one phase that has colloidal dimensions (less
than one micrometer, but greater than one nanometer) within a second phase. A colloidal solution
is a colloid in which the second phase is a liquid (also known as a sol). However, in radiochemistry, a colloid refers to the dispersion of solid particles in the solution phase. The mixture is not a
true solution: particles of the dispersed phase are larger than typical ions and molecules, and can
often be viewed by a light microscope. Colloidal precipitates are usually avoided in analytical
procedures because they are difficult to filter and to wash. Moreover, the purity of the precipitate
is controlled by the tremendously large surface area of the precipitate and by the localized
electrical character of the colloidal surface.
The stability of colloidal solutions and suspensions is governed by two major forces, one of
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attraction between the particles (van der Waals) and one of repulsion (electrical double layer)
(Salutsky, 1959). This repulsive force is a result of the adsorptive capacity of the colloidal
particles for their own ions. For instance, when silver chloride is precipitated in the presence of
excess silver ions, the particles adsorb silver ions and become positively charged. Then counterions of opposite charge (in this case, nitrate ions) tend to adsorb to the particles to form a second
electrical layer, as illustrated in Figure 14.5.
Counter ions
Adsorbed Layer
(Primary Layer)

NO3-

NO3Ag+

Cl -

Ag+
Ag+

Cl-

NO3-

NO3-

Ag+
Ag+

Cl-

NO3-

Ag+
Ag+

Cl -

Ag+
Ag +

Cl-

Ions in surface
FIGURE 14.5 — The electrical double layer: A schematic representation of adsorption of
nitrate counter-ions onto a primary adsorbed layer of silver ions at the surface of a silver
chloride crystal (Peters et al., 1974).

In a similar fashion, in the presence of a slight excess of alkali chloride, the silver chloride
particles would adsorb chloride ions and become negatively charged. Therefore, precipitates
brought down in the presence of an excess of one of the lattice ions tend to be contaminated with
ions of the opposite charge. Moreover, because all of the particles have the same charge, they
repel each other. If these repulsive forces exceed the attractive van der Waals’ forces, a stable
colloid results, and the tightness with which the counter-ions are held in and with the water layer,
or the completeness with which they cover the primary adsorbed ion layer, determines the
stability of the colloid.
Such adsorption of ions upon the surface of solids in solution is largely, but not entirely, based
upon electrical attraction, otherwise adsorption would not be selective. Recall that there are four
other factors, in addition to magnitude of charge, that affect the preferential adsorption by a
colloid (see Surface Adsorption on page 14-72).
• The Paneth-Fajans-Hahn Law dictates that when two or more types of ions are available for
adsorption, the ion that forms the least soluble compound with one of the lattice ions will be
adsorbed preferentially.
• The ion present in the greater concentration will be adsorbed preferentially.
• Ions with a large radius will be adsorbed more readily than ions with a smaller radius because
the larger ion is less hydrated by the solution and not as attracted to the solution phase.
• The ion that is closer to the same size as the lattice ion will be adsorbed preferentially. For
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example, radium ions are adsorbed tightly onto barium sulfate, but not onto calcium sulfate;
radium ions are close in size to barium ions, but are much larger than calcium ions.
If an excess of electrolyte is added to the colloidal solution, the electrical double layer is
destroyed and the particles can agglomerate to form larger particles that can settle to the bottom
of the container, a process known as flocculation (or coagulation). For example, Smith et al.
(1995) used polyethylene glycol to remove colloidal silica from a dissolved-soil solution before
the addition of the sample to an ion-exchange resin. Alternatively, the process whereby
coagulated particles pass back into the colloidal state is known as deflocculation, (or peptization). Special precautions should be taken during the washing of coagulated precipitates to assure
that deflocculation does not occur. When coagulation is accomplished through charge
neutralization, deflocculation would occur if the precipitate was washed with water. A solution
containing a volatile electrolyte such as nitric acid should be used instead.
There are two types of colloidal solutions (Salutsky, 1959):
• Hydrophobic colloids show little or no attraction for water. These solutions have a low
viscosity, can be easily flocculated by the addition of an appropriate electrolyte, and yield
precipitates that are readily filterable.
• Hydrophilic colloids have a high affinity for water and are often highly viscous. They are
more difficult to flocculate than hydrophobic colloids, and relatively large amounts of
electrolytes are necessary to cause precipitation. The flocculate keeps water strongly adsorbed
and tends to form jellylike masses that are difficult to filter.
Colloidal precipitations can be a useful separation technique. Because of their great adsorption
capacity, colloidal precipitates are excellent scavengers (collectors) for concentrating trace
substances (Salutsky, 1959). Unspecific carriers such as manganese dioxide, sulfides and
hydrated oxides are frequently used as scavengers. For example, protactinium can be efficiently
scavenged and concentrated on manganese dioxide that is precipitated by adding a manganous
salt to a solution containing permanganate. Ferric hydroxide is commonly used to scavenge
cations (Section 14.8.4.4, “Coprecipitation Methods”). Moreover, scavenging precipitations can
sometimes be used to remove interferences. For example, a radionuclide that is capable of
existing in two oxidation states can be effectively purified by precipitation in one oxidation state,
followed by scavenging precipitations for impurities, while the element of interest is in another
oxidation state. A useful procedure for cerium purification involves repeated cycles of ceric
iodate precipitation, reduction to Ce+3, zirconium iodate [Zr(IO3)4] precipitation to remove
impurities (with Ce+3 staying in solution), and reoxidation to Ce+4.

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14.8.6 Separation of Precipitates
The process of precipitation chemically separates an analyte from contaminants or other analytes.
Precipitation generally is followed by one of two techniques that physically separates the
precipitate: centrifugation or filtration.
Centrifugation is a technique that can be used for precipitates of many different physical forms.
The best way to demonstrate the utility of centrifugation in radiochemical analyses is by
example:
Example of Centrifugation
A method of radium analysis coprecipitates radium with barium using sulfuric acid to isolate
the radium from its progeny. When the precipitation is completed, the mixture is centrifuged.
The supernatant solution contains contaminants and radium progeny and is decanted. The
precipitate is washed, in situ, with an isotonic sulfuric acid solution to maintain the
insolubility of the precipitate, and to further enhance the removal of the contaminants. The
mixture is re-centrifuged and the supernate again decanted.

This example demonstrates that centrifugation separates and purifies the precipitate without
disturbing the mechanical flow of the separation process, and it minimizes the introduction of
new contaminants by using the same glassware. It is noteworthy that there are several instances
of using centrifugation to discard the precipitate and retain the supernate (e.g., the separation of
barium from strontium using chromate). Separation by filtration at this point (not the final
analytical step) would involve transfer onto and subsequent removal from the filter media.
Filtration would be time consuming and risk low yield for the analysis. The speed and capacity of
the centrifuge is dictated by the type of precipitate (e.g., gelatinous, crystalline, amorphous etc.),
the sample size being processed, and the ancillary procedural steps to purify the precipitate.
The final separation of the analyte immediately preceding counting techniques is generally best
suited by using filtration techniques. The physical nature of a precipitate not only affects the
purity of the precipitate, but also the filterability of the precipitate. Large, well-formed crystals
are desirable because they tend to contain fewer impurities, and are also easier to filter and wash.
Many coagulated colloidal precipitates, such as hydrous oxides or sulfides, tend to form slimy
aggregates and to clog the filter during filtration. There are several approaches that can be taken
to improve the physical form of the precipitate (Salutsky, 1959):
• A trace quantity of a hydrophilic colloid can be added to produce complete and rapid
flocculation. For example, gelatin has been used as a sensitizer in the precipitation of zinc
sulfide, hydrous silica, and various other hydrous oxides, as well-coagulated, filterable
precipitates (Salutsky, 1959).

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• The slow precipitation techniques described in Section 14.8.3.2, “Factors Affecting
Precipitation,” can be used to produce good precipitates.
• Aging the precipitate can result in a precipitate more amenable to filtration. During aging,
small particles with a larger solubility go into solution, and larger particles grow at the cost of
the smaller ones (see “Digestion” under Section 14.8.3.2, “Factors Affecting Precipitation”).
Ostwald ripening results in a decrease in the number of particles and, therefore, a decrease in
surface area. The speed of aging generally increases with temperature and with the increasing
solubility of the precipitate in the aging media. Shaking can sometimes promote aging,
perhaps by allowing particles to come into contact and to cement together.
14.8.7 Advantages and Disadvantages of Precipitation and Coprecipitation
14.8.7.1

Advantages

•
•
•
•
•
•

Provides the only practical method of separation or concentration in some cases.
Can be highly selective and virtually quantitative.
High degree of concentration is possible.
Provides a large range of scale (mg to industrial).
Convenient, simple process.
Carrier can be removed and procedure continued with tracer amounts of material (e.g., carrier
iron separated by solvent extraction).
• Not energy- or resource-intensive compared to other techniques (e.g., solvent extraction).

14.8.7.2
•
•
•
•
•

Disadvantages

Can be time consuming to digest, filter, or wash the precipitate.
Precipitate can be contaminated by carrying of ions or postprecipitation.
Large amounts of carrier might interfere with subsequent separation procedures.
Coprecipitating agent might contain isotopic impurities of the analyte radionuclide.
Scavenger precipitates are not as selective and are more sensitive to changes in separation
procedures.

14.9 Carriers and Tracers
14.9.1 Introduction
Radiochemical analysis frequently requires the radiochemist to separate and determine radionuclides that are present at extremely small quantities. The amount can be in the picomole range or
less, at concentrations in the order of 10!15 to 10!11 molar. Analysis of radionuclides using
counting techniques, such as alpha spectrometry, liquid scintillation, proportional counting, or
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gamma spectrometry, allows activities of radionuclides to be determined easily, even though the
number of atoms (and mass percent) of these materials is vanishingly small. Table 14.15 identifies the number of atoms and mass present in several radionuclides, based on an activity of 500
dpm (8.33 Bq).
TABLE 14.15 — Atoms and mass of select radionuclides equivalent to 500 dpm
Radionuclide
Half-life*
Number of Atoms
Mass (g)
11
2.3 × 10!10
Radium-226
1,600 y
6.0 × 10
8
Polonium-210
138.3 d
1.5 × 10
5.0 × 10!14
Lead-212
10.6 h
4.5 × 105
1.6 × 10!16
Thallium-208
3.1 min
2.3 × 103
8.0 × 10!19
*
Half-lives taken from Brookhaven National Laboratory, National Nuclear Science Database (www.nndc.bnl.gov/).

Considering the minute masses of these analytes and their subsequently low concentration in
solution, it is obvious why conventional techniques of analysis, such as gravimetry, spectrophotometry, titrimetry, and electrochemistry, cannot be used for their quantitation. However, it is
not immediately obvious why these small quantities might present other analytical difficulties.
As described below, the behavior of such small quantities of materials can be seriously affected
by macro constituents in an analytical mixture in a way that may be unexpected chemically.
14.9.2 Carriers
The key to radiochemical analysis of samples with multiple radionuclides is effective separation
of the different analytes. Separations are most easily accomplished when performed on a macro
scale. As described above, however, the analytes are frequently at levels that challenge the
analyst and the conventional methods to perform the separations. The use of a material that is
different in isotopic make-up to the analyte and that raises the effective concentration of the
material to the macro level is referred to as a carrier. In many cases, the carrier is a nonradioactive isotope of the analyte. Some carriers are stable isotopes of chemically similar elements.
A distinction exists between traditional and radiochemical analyses when referring to macro
amounts. Generally, carriers are present in quantities from a few tenths to several hundred
milligrams of material during the progress of the radiochemical separation.
14.9.2.1

Isotopic Carriers

An isotopic carrier is usually a stable isotope of the analyte. Stable strontium (consisting of
naturally occurring 84Sr, 86Sr, 87Sr, and 88Sr) is frequently used as the carrier in the analysis of 89Sr
and 90Sr. Regardless of the stability of the isotope, the number of protons in the nucleus
ultimately governs the chemical properties of the isotope. Thus, all nuclei that have 38 protons
are strontium and react as strontium classically does.
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The purpose of adding a carrier is to raise the chemical concentration of the analyte to the point
where it can be separated using conventional techniques, but for the carrier to perform properly,
it must have the same oxidation state and chemical form as the analyte. It is important then to add
the carrier to the sample as early as possible in chemical process. For example, in the determination of 131I in milk, the radioiodine might be present as I!1, IO3!1, CH3I, or I2. The analyst should
assume that all states are present, and treat the sample so that all atoms are brought to a common
oxidation state and chemical form during some step in the procedure, before any separation takes
place. If the final step is precipitation of AgI and the carrier is in the IO3!1 form, no precipitate
will form because AgIO3 that forms when Ag+1 is added is relatively soluble compared to AgI.
Furthermore, if separations of other radioisotopes are performed before this step, there is the
possibility that quantities of the radioiodine could be trapped in the precipitate with other
separated analytes. When concentrations of these materials are very small, even small losses are
significant. The carrier also functions to prevent losses of the analyte during the separation of
other radionuclides or interfering macro-contaminants. This is another reason that it is essential
to add the carrier prior to any chemical treatment of the sample.
The laws of equilibrium for precipitation, distillation, complexation, and oxidation-reduction will
apply to the entire chemical form of analyte in solution, both carrier and radioisotope. If, for
example, 99.995 percent of all strontium is determined to be precipitated during a radiochemical
procedure, then the amount of stable strontium remaining in solution will be 0.005 percent,
which means that 0.005 percent of the radiostrontium still remains in the solution as well. Losses
such as this occur during any chemical process. Frequently then, carriers are used in radiochemical analyses not only to raise the chemical concentration of the element, but also to determine the
yield of the process. In order to determine the exact amount of radionuclide that was originally
present in the sample, the yield (sometimes called the recovery) of the radionuclide collected at
the end of the procedure should be known. However, because the amount of analyte at the start of
the procedure is the unknown, the yield should be determined by an alternate method. The mass
of the radioanalyte is insignificant in comparison to the carrier, and measuring the yield of the
carrier (gravimetrically, for example) will allow the calculation of the yield of the analyte.
14.9.2.2

Nonisotopic Carriers

Nonisotopic carriers are materials that are similar in chemical properties to the analyte being
separated, but do not have the same number of protons in their nucleus. Usually these carriers
will be elements in the same family in the periodic table. In the classical separation of radium by
the Curies, the slight difference in solubility of radium chloride versus barium chloride allowed
the tedious fractional crystallization of radium chloride to take place (Hampel, 1968). When
barium is present in macro-quantities and the radium in femtogram quantities, however, the two
may be easily precipitated together as a sulfate.
For several elements, nonisotopic carriers are chosen from a different family of elements, but
they have the same ionic charge or similar crystalline morphology as the analyte. Lanthanum and
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neodymium as +3 ions are frequently used as nonisotopic carriers for U+4 and Pu+4 in their final
separation as insoluble fluorides by the process of coprecipitation (Metz and Waterbury, 1962)
(see also Section 14.8, “Precipitation and Coprecipitation”). The chemical form of the uranium
and plutonium is particularly important for this process; the +4 oxidation state will coprecipitate,
but the (VI) form will not. Uranium(VI) is present in solution as UO2+2 and, therefore, will not be
coprecipitated with lanthanum fluoride. However, it is very important to note that even though
the precipitation of LaF3 may be quantitative (i.e., >99.995 percent may be precipitated), there is
no measure of how much uranium will also be coprecipitated. Because uranium and lanthanum
are not chemically equivalent, the laws of solubility product constant for lanthanum cannot be
applied to uranium. For these types of processes, separate methods, usually involving a tracer
isotope of the analyte, should be used to determine the chemical yield of the process.
For alpha counting, rare-earth fluorides (such as NdF3) are frequently used to coprecipitate the
transuranic elements (Hindman, 1983 and 1986; Sill and Williams, 1981).
Another group of nonisotopic carriers can be described as general scavengers. Substances with
high surface areas, or the ability to occlude contaminants in their floc, can be used to effect gross
separation of all radionuclides from macro quantities of interfering ions. Ferric hydroxide,
manganese dioxide (MnO2) and sulfides (MnS), and hydrated oxides [Mn(OH)x] are examples of
these nonspecific carriers that have been used in many radiochemical separations to eliminate
gross quantities of interfering substances.
14.9.2.3

Common Carriers

Carriers for specific analytes are discussed below.
Alkaline Earths
STRONTIUM AND BARIUM. Radioisotopes of Sr+2 and Ba+2 will coprecipitate with ferric hydroxide
[Fe(OH)3], while Ca+2 exhibits the opposite behavior and does not coprecipitate with ferric
hydroxide. Lead sulfate (PbSO4) will also carry strontium and barium.
Frequently, inactive strontium and barium are used as carriers for the radionuclides in order to
facilitate separation from other matrix constituents and from calcium. The precipitates used most
frequently in radiochemical procedures are the chromates (CrO4!2), nitrates (NO3!1), oxalates
(C2O4!2), sulfates (SO4!2), and barium chloride (BaCl2). Several different methods of separation
are identified here:
• Chromate precipitation is used in the classical separation of the alkaline earths. Barium
chromate (BaCrO4) is precipitated from a hot solution buffered to a pH of 4 to minimize
strontium and calcium contamination of the barium precipitate. Ammonium ion (NH4+1) is
then added to the solution, and strontium chromate (SrCrO4) is precipitated.
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• Barium and strontium can be separated from calcium as the nitrates. Fuming nitric acid is
used to increase the nitric acid concentration to 60 percent, conditions at which barium and
strontium nitrate [Ba(NO3)2 and Sr(NO3)2] precipitate and calcium does not.
• Oxalate precipitation does not separate one alkaline earth from another, but it is usually used
to produce a weighable and reproducible form suitable for radioassay. The precipitation is
accomplished from a basic solution with ammonium oxalate [(NH4)2C2O4].
• Barium sulfate (BaSO4) precipitation is generally not used in separation procedures. It is
more common as a final step to produce a precipitate that can be readily dried, weighed, and
mounted for counting. Barium is readily precipitated by slowly adding dilute sulfuric acid
(H2SO4) to a hot barium solution and digesting the precipitate. For the precipitation of
strontium or calcium sulfate (SrSO4 or CaSO4), a reagent such as alcohol should be added to
lower the solubility, and the precipitant must be coagulated by heat.
• Insolubility of barium chloride (BaCl2) in strong hydrochloric acid solution (HCl) is the basis
of the method to separate barium from calcium, strontium, and other elements. The
precipitation is performed either by adding an ether-hydrochloric acid solution or by bubbling
dry hydrogen chloride gas into the aqueous solution.
RADIUM. Radium yields the same types of insoluble compounds as barium: sulfates, chromates,
carbonates (CO3!2), phosphates (PO4!3), oxalates, and sulfites (SO3!2). Hence, Ra coprecipitates
with all Ba compounds and, to a lesser extent, with most Sr and Pb compounds. Barium sulfate
and barium chromate are most frequently used to carry radium. Other compounds that are good
carriers for radium include ferric hydroxide when precipitated at moderately high pH with
sodium hydroxide (NaOH), barium chloride when precipitated from a cold mixed solvent of
water and alcohol saturated with hydrochloric acid, barium iodate (BaIO3) and various insoluble
phosphates, fluorides and oxalates (e.g., thorium phosphate [Th3(PO4)], lanthanum fluoride
(LaF3), and thorium oxalate [Th(C2O4)].
Rare Earths, Scandium, Yttrium, and Actinium
Ferric hydroxide and calcium oxalate (CaC2O4) will coprecipitate radioisotopes of the rare earths
without difficulty.
The rare earths will coprecipitate one with another in almost all of their reactions; one rare earth
can always be used to coprecipitate another. The rare earth hydroxides, fluorides, oxalates, and 8hydroxyquinolates in ammoniacal solution are insoluble. Conversely, the rare earth hydroxides
will carry a number of elements that are insoluble in basic solution; the rare earth oxalate will
coprecipitate calcium; and the rare earth fluorides tend to carry Ba and Zr. In the absence of
macro quantities of rare earths, actinium will carry on barium sulfate and lead sulfate (PbSO4).
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Lead
Ferric hydroxide and aluminum hydroxide [Al(OH)3] carry lead very effectively from ammonium
solutions under a variety of conditions. Lead is carried by barium or radium chloride, but not
carried by barium or radium bromide (BaBr2 or RaBr2). This behavior has been used to separate
radiolead isotopes from radium salts. Lead is also carried by barium carbonate (BaCO3), barium
sulfate, radium sulfate, radium chloride, lanthanum carbonate [La2(CO)3], barium chloride, and
silver chromate (Ag2CrO4). Calcium sulfate in the presence of alcohol has also been used to
coprecipitate lead.
Polonium
Trace quantities of polonium are carried almost quantitatively by bismuth hydroxide [Bi(OH)3]
from ammoniacal solution. Ferric, lanthanum, and aluminum hydroxides have also been used as
carriers for polonium in alkaline solutions. Colloidal platinum and coagulated silver hydroxide
(AgOH) and ferric hydroxide sols have been used to carry polonium. Because of the high
oxidation state of polonium, it is susceptible to being a contaminant in almost any precipitate.
Removal of polonium by electrodeposition on nickel metal is recommended prior to final
precipitation for any gross counting technique (proportional counting and liquid scintillation, for
example).
Actinides
THORIUM. Thorium will coprecipitate with ferric, lanthanum [La(OH)3], and zirconium
hydroxide [Zr(OH)4]. These hydroxide carriers are nonspecific, and therefore, will only remove
thorium from a simple group of contaminants or as a group separation. The ferric hydroxide
precipitation is best carried out at pH 5.5 to 6.
Thorium will coprecipitate quantitatively with lanthanum fluoride from strongly acidic solutions,
providing an effective means to remove small quantities of thorium from uranium solutions.
However, the rare earths will also carry quantitatively, and zirconium and barium radioisotopes
will carry unless macro quantities of these elements are added as holdback carriers (see Section
14.9.2.4, “Holdback Carriers”).
Precipitation of thorium with barium sulfate is possible from strongly acidic solutions containing
high concentrations of alkali metal sulfates; however, this coprecipitation is nonspecific. Other
actinides, lead, strontium, rare earths, bismuth, scandium (Sc), and yttrium will also carry.
Coprecipitation of thorium on hydrogen hypophosphate (HPO3!2) or phosphate carriers can be
performed from rather strongly acidic solutions. Zirconium phosphate [Zr3(PO4)4] serves as a
good carrier for trace levels of thorium. Moreover, thorium also will carry quantitatively on
zirconium iodate from a strongly acidic solution. If coprecipitation is performed from a strongly
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acidic solution and the precipitate is washed with a solution containing iodate, the rare earths and
actinium are eliminated. Cesium(+4) must be reduced to Ce+3 before precipitation so that it does
not carry.
PROTACTINIUM. Protactinium will be carried quantitatively on hydroxide, carbonate, or
phosphate precipitates of tantalum, zirconium, niobium, hafnium, and titanium. It is also carried
by adsorption onto flocculent precipitates of calcium hydroxide [Ca(OH)2)] or ferric hydroxide,
and it is carried by manganese dioxide, which is produced by addition of potassium
permanganate (KMnO4) to a dilute nitric acid (HNO3) solution containing manganese nitrate.
However, titanium and zirconium are also carried under these conditions.
URANIUM. Trace concentrations of uranium can be coprecipitated with any of the common
insoluble hydroxides. When coprecipitating U(VI) with hydroxides at pH 6 to 7, the ammonium
used must be free of carbonate or some of the uranium will remain in solution as the stable
anionic carbonate complex. Hydroxide precipitation is nonspecific, and many other metals will
carry with the uranium.
Uranium(+4) can be coprecipitated as the fluoride or phosphate [UF4 or U3(PO4)4] from relatively
strong acid media; however, U(VI) phosphate [(UO2)3(PO4)2] is precipitated only from very weak
acid solutions (pH 5 to 6) by the addition of carbonate-free ammonium. The rare earths, and other
metals can also coprecipitate under these conditions.
In general, U+4 should behave similarly to Pu+4 and Np+4, and should be carried by lanthanum
fluoride, ceric and zirconium iodates [Ce(IO4)3 and Zr(IO3)4], cesium and thorium oxalates
[Th(C2O4)2], barium sulfate, zirconium phosphate [Zr3(PO4)4], and bismuth arsenate (BiAsO4).
However, U(VI) does not carry with these agents as long as the concentration of either carrier or
that of uranium is not too high.
PLUTONIUM AND NEPTUNIUM. Classically, plutonium and neptunium in their ter- and tetravalent
oxidation states have been coprecipitated with lanthanum fluoride in the method most widely
used for the isolation of femtograms of plutonium. However, large amounts of aluminum
interfere with coprecipitation of plutonium, and other insoluble fluorides, such as the rare earths,
calcium, and U+4, coprecipitate.
AMERICIUM AND CURIUM. Bismuth phosphate (BiPO4), which historically has been used to
precipitate plutonium, will also carry americium and curium from 0.1–0.3 M nitric acid.
Impurities such as calcium and magnesium are not carried under these conditions.
Lanthanum fluoride provides a convenient carrier for Am+3 and Cm+3. A lanthanum fluoride
precipitation is not totally specific, but it can provide a preliminary isolation from the bulk of the
fission products and uranium. Additionally, a lanthanum fluoride precipitation can be used to
separate americium from curium. Am+3 is oxidized to Am(V) in dilute acid with persulfate, and
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fluoride is added to precipitate Cm+3 on lanthanum fluoride.
14.9.2.4

Holdback Carriers

It is often necessary to add holdback carriers to analytical mixtures to prevent unwanted radionuclides from being carried in a chemical process. Coprecipitation of a radionuclide with ferric
hydroxide carries other ions in addition to the analyte, because of its tendency to adsorb other
ions and occlude them in its crystal matrix. The addition of a holdback carrier, a highly charged
ion, such as Co+3, represses counter-ion exchange and adsorption to minimize the attraction of
foreign ions. The amount of a given substance adsorbed onto a precipitate depends on its ability
to compete with other ions in solution. Therefore, ions capable of displacing the radionuclide
ions (the hold-back carrier) are added to prohibit the coprecipitation of the radionuclide. Highly
charged ions, chemical homologs, and ions isotopic with the radionuclide are among the most
efficient holdback carriers. Hence, the addition of inactive strontium makes it possible to precipitate radiochemically pure radiobarium as the nitrate or chloride in the presence of radiostrontium.
Actinium and the rare earth elements can be separated from zirconium and radium by lanthanum
fluoride coprecipitation with the addition of zirconium and barium holdback carriers. Holdback
carriers are used in other processes as well. The extraction of lutetium from water employs
neodymium ions (Nd+3) to avoid adsorption loses (Choppin et al., 1995).
14.9.2.5

Yield of Isotopic Carriers

The use of an isotopic carrier to determine the chemical yield (recovery) of the analyte is a
critical step in the plan of a radiochemical analysis. The analytical method being used to
determine the final amount of carrier will govern the method of separation. If a gravimetric
method is to be used for the final yield determination, the precipitate must have all the
characteristics that would be used for macro gravimetric analysis—easily dried, definite
stoichiometry, nonhygroscopic, etc.
Similarly, the reagent used as source of carrier at the beginning of the analysis must be of
primary-standard quality to ensure that the initial mass of carrier added can be determined very
accurately. For a gravimetric yield determination, the equation would be the following:
 mass of carrier in final separation step 
Percent Yield = 
 × 100
mass of carrier added


It should be recognized that the element of interest is the only quantity used in this formula. For
example, if strontium nitrate is used as the primary standard and strontium sulfate is the final
precipitate, both masses should be corrected, using a gravimetric factor, so that only the mass of
strontium is used in the equation in both the numerator and denominator.

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Other methods to determine the yield of the carrier include atomic absorption spectrometry, ultraviolet/visible spectrometry, titrimetry, and potentiometry.
14.9.3 Tracers
The term “tracer”was used classically to express the concentration of any pure radionuclide in
solution that had a mass too small to be measured by an analytical balance (<10!5 to 10!6 g).
More recently, the definition of a tracer has become more pragmatic. The current definition of a
tracer is a known quantity of a radioisotope that is added to a solution of a chemically equivalent
radioisotope of unknown concentration so that the yield of the chemical separation can be
monitored. In general, a tracer is not a carrier, and a carrier is not a tracer.
The analysis of 241Am in an environmental sample provides an example of a radioisotope
employed in a manner consistent with the recent use of the term tracer. In the analytical
procedure, no stable isotope of americium exists to act as a carrier. Femtogram quantities of
243
Am can be produced, however, with accurately known activities. If a known quantity of 243Am
in solution is added to the unknown sample containing 241Am at the beginning of the separation
procedure, and if the resulting activity of 243Am can be determined at the end of the procedure,
then the yield of 241Am can be determined accurately for the process. Americium-243 added to
the sample in this example is used as a tracer. A measurable mass of this element was not used,
but a known activity was added through addition of the solution. During the course of the
radiochemical separation, lanthanides may have been used to help carry the americium through
analysis. However, they are not used to determine the yield in this example and would be
considered, therefore, a nonisotopic carrier.
When using a tracer in an analytical method, it is important to consider the availability of a
suitable isotope, its chemical form, its behavior in the system, the amount of activity required, the
form in which it should be counted, and any health hazards associated with it (McMillan, 1975).
Perhaps the most important property of the tracer is its half-life. It is preferable to select an
isotope with a half-life that is long compared to the duration of the experiment. By doing so, one
avoids the problems of having to handle high levels of activity at the beginning of the experiment
and of having to make large decay corrections.
Purity of the tracer is of critical importance. Radionuclide and radiochemical impurities are the
two principal types of impurities encountered. Radionuclide impurity refers to the presence of
radionuclides other than those desired. For instance, it is very difficult to obtain 236Pu tracer that
does not contain a very small quantity of 239Pu. This impurity should be taken into account when
calculating the 239Pu activity levels of samples. Radiochemical impurity refers to the nuclide of
interest being in an undesired chemical form. This type of impurity has its largest effects in
organic tracer studies, where the presence of a tracer in the correct chemical form is essential. For
example, the presence of 32P-labeled pyrophosphate in an orthophosphate tracer could lead to
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erroneous results in an orthophosphate tracer study.
Tracer solutions can also contain other forms of radiochemical impurities. Many tracers are
actinides or other isotopes that have progeny that are radioactive. Tracer solutions are purchased
with known specific activities for the isotopes listed in the solutions. However, from the time of
production of the tracer, ingrowth of progeny radioisotopes occurs. Plutonium-236 is used as a
tracer for 239/240Pu analysis, for example. Plutonium-236 has a half-life of 2.9 years and decays to
232
U, which has a half-life of 72 years. After solutions of 236Pu have been stored for about three
years, half of the radionuclide will be converted to 232U. If the solution is then used as a tracer in
a procedure for analysis of uranium and plutonium in soil, erroneously high results would be
produced for the content of uranium if a gross-counting technique is used. Thus, it is important to
consider chemical purification of a tracer solution prior to use to remove unwanted radioactive
progeny.
Tracer analysis is very dependent upon the identical behavior of the tracer and the analyte.
Therefore, tracers should be added to the system as early as possible, and complete isotopic
exchange should be ensured as discussed previously (see Section 14.10, “Analysis of Specific
Radionuclides”). Obvious difficulties arise when a tracer is added to a solid sample, especially if
the sample is subdivided. Unless complete dissolution and isotopic exchange is ensured, results
should be interpreted carefully.
Isotopes selected for tracer work should be capable of being easily measured. Gamma-emitting
isotopes are ideal because they can easily be detected by gamma spectroscopy without being
separated from other matrix constituents. Alpha- and beta-emitting tracers require separation
before counting. Some common tracers are listed below:
• Strontium-85 has a 514 keV gamma ray that can be used to monitor the behavior of strontium
in a system, or for yield determination in a 89Sr/90Sr procedure, as long as the gamma is
accounted for in the beta-counting technique.
• Technetium-99m, with a half-life of 6.02 h and a 143 keV gamma ray, is sometimes used as a
yield monitor for 99Tc determinations. Samples are counted immediately to determine the
chemical recovery, then the 99mTc is allowed to decay before analysis of the 99Tc.
• Europium-152 and 145Sm are frequently used in the development of a new method to estimate
the behavior of the +3 actinides and lanthanides.
• Tritium, 14C, 32P, and 36Cl are frequently used in biological studies. In some of these studies,
the radionuclide is covalently bonded to a molecule. As a result, the chemical behavior of the
radionuclide will follow that of the molecule, not the element.
• Thorium-229 is used for Th determinations, both in alpha spectroscopy and inductively
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coupled plasma-mass spectroscopy (ICP-MS).
• Uranium-232 is commonly used as a tracer in alpha spectroscopy, whereas 233U is used
commonly for ICP-MS determinations. It should be noted that 232U decays to 228Th and
therefore this needs to be taken into account when determining other alpha emitters.
• Plutonium-242 and 236Pu are both used as tracers in Pu analyses. However, 236Pu decays to
232
U, which needs to be taken into account when analyzing both Pu and U in the same sample
aliquant.
• Americium-243 is employed in the analysis of 241Am and Cm by alpha spectroscopy. It is
assumed that Am and Cm are displaying similar chemical behavior.
14.9.3.1

Characteristics of Tracers

The behavior of tracers is often different from that of elements in normal concentrations. The
chemical form of a radionuclide predominant at normal concentrations, for example, might not
be the primary form at tracer concentrations. Alternatively, a shift in the equilibrium that is partly
responsible for a radionuclide’s chemical behavior might increase or reduce its concentration as a
result of the low tracer concentration. Hydrolysis reactions are influenced particularly by changes
in concentration because water is one of the species in the equilibrium. For example, hydrolysis
of the uranyl ion is represented by (Choppin et al., 1995):
m @ UO2+2 + p· H2O 6 (UO2)m(OH)p2m!p + p· H+1
At tracer quantities, the equilibrium will shift to the left as the amount of the uranyl ion
decreases. At 10!3 molar (pH 6), the uranyl ion is 50 percent polymerized; at 10!6 molar, there is
negligible polymerization.
Interactions of radionuclides with impurities present special problems at low concentration.
Difficulties include adsorption onto impurities such as dust, silica, or colloidal or suspended
material, or adsorption onto the walls of the container. Generally, 10 !8 to 10!7 moles are needed
to cover a container’s walls; but at tracer concentrations, much less is present (Choppin et al.,
1995). Adsorption depends on (see Surface Adsorption on page 14-72):
• Concentration. A larger percentage is adsorbed at lower tracer concentrations than at higher
concentrations, because a larger surface area is available compared to the amount of tracer
present. Dilution with carrier decreases the amount of tracer adsorbed because the carrier is
competing for adsorption, and the relative amount of tracer interacting with the walls is much
less.
• Chemical State. Adsorption increases with charge on the ion.
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• Nature of the Surface Material. Surfaces that have a negative charge or that contain hydroxyl
groups can interact with cations through electrostatic attraction and hydrogen bonding,
respectively.
• pH. Generally, adsorption decreases with a lower pH (higher hydrogen ion concentration)
because the ions interact with negatively charged surfaces, and hydrogen bonding decreases
their ability to interact with metal ions.
All these processes will reduce the quantity of analyte available for radiochemical procedures
and, therefore, the yield of a procedure. The amount measured by the detection process will be
correspondingly lower, introducing additional uncertainty that would go undetected at normal
concentrations.
However, the adsorption process has been shown to be useful in some instances. For example,
carrier-free Y+3 is quantitatively adsorbed onto filter paper from basic strontium solutions at
concentrations at which yttrium hydroxide, Y(OH)3, will not precipitate. Also, carrier-free Nb has
been adsorbed on glass fiber filters for a fast specific separation technique (Friedlander et al.,
1981).
Specific behavior characteristics of compounds in separation techniques are further described
below. Additional discussion can also be found in the respective sections found earlier in this
document that describe each separation technique.
14.9.3.2

Coprecipitation

Often, the concentration of tracer is so low that precipitation will not occur in the presence of a
counter-ion that, at normal concentrations, would produce an insoluble salt. Under these
conditions, carriers are used to coprecipitate the tracer (coprecipitation is described in
Section 14.8.4).
14.9.3.3

Deposition on Nonmetallic Solids

Radionuclides can be deposited onto preformed ionic solids, charcoal, and ion-exchange resins
(Wahl and Bonner, 1951). The mechanisms of adsorption onto preformed ionic solids are similar
to those responsible for coprecipitation: counter-ion exchange and isomorphous exchange
(Section 14.8, “Precipitation and Coprecipitation”). Adsorption is favored by a large surface area,
charge of the solid and radionuclide, solubility of compound formed between the solid and the
radionuclide, and time of contact; however, it depends, to a large extent, on whether or not the
radionuclide ion can fit into the crystal lattice of the precipitate. Similarly, adsorption onto
charcoal depends on the amount of charcoal and its surface area, time of contact, and nature of
the surface, because it can be modified by the presence of other ions or molecules.
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Adsorption of radionuclides, with and without carriers (Friedlander et al., 1981), onto ionexchange resins, followed by selective elution, has been developed into a very efficient
separation technique (Wahl and Bonner, 1951) (see Section 14.7.4, “Ion-Exchange
Chromatography”). Friedlander et al. (1981) illustrates this phenomenon:
“Ion-exchange separations generally work as well with carrier-free tracers as with weighable
amounts of ionic species. A remarkable example was the original isolation of mendelevium at
the level of a few atoms ...The transuranium elements in the solution were ... separated from
one another by elution ... through a cation-exchange column.”
14.9.3.4

Radiocolloid Formation

At the tracer level, a radionuclide solution is not necessarily truly homogeneous, but can be a
microparticle (colloid) of variable size or aggregation (Adolff and Guillaumont, 1993). Carrierfree tracers can become colloidal by two mechanisms:
1. Sorption onto a preexisting colloidal impurity (approximately 0.001 to 0.5 µm), such as
dust, cellulose fibers, glass fragments, organic material, and polymeric metal hydrolysis
products (Adolff and Guillaumont, 1993; Choppin et al., 1995).
2. Polycondensation of a monomeric species consisting of aggregates of 103 to 107
radioactive atoms (Adolff and Guillaumont, 1993).
The presence of radiocolloids in solution can be detected by one or more of the following
characteristics of the solution, which is not typical behavior of a true solution (Adolff and
Guillaumont, 1993):
• The radionuclide can be separated from solution by a physical method such as ultrafiltration
or ultracentrifugation.
• The radionuclide does not follow the laws of a true solution when a chemical gradient
(diffusion, dialysis, isotopic exchange) or electrical gradient (electrophoresis, electrolysis,
electrodialysis) is applied.
• Adsorption on solid surfaces and spontaneous deposition differ from those effects observed
for radionuclides in true solution.
• Autoradiography reveals the formation of aggregates of radioactive atoms.
Several factors affect the formation of radiocolloids (Wahl and Bonner, 1951):
• Solubility of the Tracer. The tendency of the tracer radionuclide to hydrolyze and form an
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insoluble species with another component of the solution favors radiocolloid formation,
while the presence of ligands that form soluble complexes hinders formation; low pH tends
to minimize hydrolysis of metallic radionuclides.
• Foreign Particles. The presence of foreign particles provides sites for the tracer to adsorb
onto their surfaces. Ultrapure water prepared with micropore filters reduces the amount of
foreign particles. However, the preparation of water that is completely free of suspended
particles is difficult.
• Electrolytes. Electrolytes affect the nature (species) of the tracer ions in solution (see Section
14.10, “Analysis of Specific Radionuclides”), as well as the charge on both the radiocolloid
and the foreign particle from which the colloid might have been derived.
• Solvent. Polar and nonpolar solvents can favor the formation of radiocolloids, depending on
the specific radiocolloid itself.
• Time. The amount of radiocolloidal formation generally increases with the age of solution.
14.9.3.5

Distribution (Partition) Behavior

Distribution (partition) coefficients, which reflect the behavior of solutes during solvent
extraction procedures (Section 14.4, “Solvent Extraction”), are virtually independent of
concentration down to tracer concentrations (Friedlander et al., 1981). Whenever the radioactive
substance itself changes into a different form, however, the coefficient naturally changes,
affecting the distribution between phases during extraction or any distribution phenomena, such
as ion-exchange or gas-liquid chromatography (Section 14.7, “Chromatography”). Several
properties of tracer solutions can alter the physical or chemical form of the radionuclide in
solution and alter its distribution behavior (Wahl and Bonner, 1951):
• Radiocolloid formation might concentrate the radionuclide in the alternate phase or at the
interface between the phases.
• Shift in equilibrium during complex-ion formation or hydrolysis reactions can alter the
concentration of multiple radionuclide species in solution (Section 14.9.3.1, “Characteristics
of Tracers”).
14.9.3.6

Vaporization

Radioisotope concentrations that challenge the minimum detectable concentration (MDC) can be
vaporized from solid surfaces or solution (Section 14.5, “Volatilization and Distillation”). Most
volatilization methods of these trace quantities of radionuclides can be performed without
specific carriers, but some nonisotopic carrier gas might be required (Friedlander et al., 1981).
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Vaporization of these amounts of materials from solid surfaces differs from the usual process of
vaporization of macroamounts of material, because the surface of the solid is usually not
completely covered with the radionuclide (Wahl and Bonner, 1951). Carrier-free radionuclides at
the surface are bonded with the surface particles instead of with themselves, and the bonds
broken during the process are between the solid and the radioisotope, rather than between the
radioisotope particles themselves. Additionally, the nature of the radioisotope can be altered by
trace quantities of gases such as oxygen and water present in the vacuum. Therefore, the identity
of the radionuclide species vaporizing might be uncertain, and the data from the procedure can be
hard to interpret. The rate of vaporization of radioisotopes also decreases with time, because the
number of radioisotope particles available on the solid surface decreases with time.
Radioisotopes near the MDC and macroquantities of radionuclide solutes should behave very
similarly in vaporization experiments from solution, however, because both are present as a
small fraction of the solution. They are, therefore, surrounded and bonded to solvent molecules
rather than to other solute particles (Wahl and Bonner, 1951). The nature of the solvent, the pH,
and the presence of electrolytes generally affect the solubility of the solute and its vaporization
behavior.
14.9.3.7

Oxidation and Reduction

Some radionuclides exist in only one oxidation state in solution, but others can exist in several
stable states (Tables 14.1 and 14.2). If multiple states are possible, it might be difficult to
ascertain in which state the radionuclide actually exists because the presence of trace amounts of
oxidation or reduction (redox) impurities might convert the radionuclide to a state other than the
one in which it was prepared (Wahl and Bonner, 1951). Excess redox reagents can often be
added to the solution to convert the forms to a fixed ratio and keep the ratio constant during
subsequent procedures.
For a redox equilibrium such as:
PuO2+2 + 4 H+1 + Hg 6 Pu+4 + Hg+2 + 2 H2O
the Nernst equation is used to calculate the redox potential, E, from the standard potential, E0:
E = E0 ! kT ln([Pu+4][Hg+2]/[PuO2+2][H+1]4)
where k is a constant for the reaction (R/2F, containing the ideal gas constant, R, and Faraday’s
constant, F) and T is the absolute temperature. Water and metallic mercury (Hg) do not appear in
the equation, because their activity is one for a pure substance. Minute concentrations of ions in
solution exhibit the same redox potential as macroquantities of ions, because E depends on the
ratio of ion concentrations and not their total concentration.
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Electrolysis of some solutions is used for electrodeposition of a carrier-free metal on an electrode
(Choppin et al., 1995) or other substance, leaving the impurities in solution (Friedlander et al.,
1981). The selectivity and efficiency, characteristic of deposition of macroquantities of ions at a
controlled potential, is not observed, however, for these metals. The activity of the ion is not
known, even if the concentration is, because the activity coefficient is dependent on the behavior
of the mixed electrolytic system. In addition, the concentration of the metal in solution might not
be known because losses may occur through adsorption or complexation with impurities.
Electrolytic deposits are usually extremely thin—a property that makes them useful for alpha,
beta, or proportional counting measurements (Wahl and Bonner, 1951).
Deposition by electrochemical displacement is sometimes used for the separation of tracer from
bulk impurities (Friedlander et al., 1981). Polonium and lead spontaneously deposit from a
solution of hydrochloric acid onto a nickel disk at 85 EC (Blanchard, 1966). Alpha and beta
counting is then used to determine 210Po and 210Pb. The same technique is frequently used in lowlevel analysis of transuranic elements to remove lead and polonium so that they do not interfere
with the subsequent alpha analysis of the elements. Wahl and Bonner (1951, Table 6F) contains
electrochemical methods used for the oxidation and reduction of carrier-free tracers.

14.10 Analysis of Specific Radionuclides
14.10.1

Basic Principles of Chemical Equilibrium

Radiochemical analysis is based on the assumption that an element reacts the same chemically,
whether or not it is radioactive. This assumption is valid when the element (analyte) and the
carrier/tracer are in the same oxidation state, complex, or compound. The atomic weight of most
elements is great enough that the difference in atomic weight between the radionuclide of interest
and the carrier or tracer will not result in any chemical separation of the isotopes. This assumption might not be valid for the very lightest elements (e.g., H, Li, Be, and B) when mass
fractionation or measuring techniques are used.
It is important to note that “chemical equilibrium” and “radioactive equilibrium” are two distinct
phenomena that come together when performing chemical separations of radionuclides. See
Attachment 14A at the end of this chapter for a thorough discussion of the phenomenon of
“radioactive equilibrium.”
Most radiochemical procedures involve the addition of one of the following:
• A carrier of natural isotopic composition (i.e., the addition of stable strontium carrier to
determine 89/90Sr; EPA, 1980, Method 905.0).
• A stable isotope tracer (i.e., enriched 18O, 15N, and 13C, are frequently used in mass
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spectroscopy studies).
• A radionuclide tracer (i.e., the addition of a known quantity of 236Pu tracer to determine 239Pu
by alpha spectroscopy; DOE, 1990 and 1997, Method Pu-02).
To achieve quantitative yields, there must be complete equilibration (isotopic exchange) between
the added isotope and all the analyte species present. In the first example, isotopic exchange of
the carrier with the radiostrontium is achieved and a weighable, stoichiometric compound of the
carrier and radionuclide are produced. The chemical recovery from the separation technique is
determined gravimetrically. Alternatively, a known quantity of a radioactive strontium isotope
(i.e., 85Sr) could be added and determined by the method appropriate for that analysis.
Carriers and tracers are added as soon in the sample preparation process as possible, usually after
the bulk sample is dried and homogenized, but before sample decomposition to ensure that the
chemistry of the carriers or tracers is truly representative of the radioisotope of interest. Thus,
losses occurring during sample preparation steps, before decomposition, are not quantified and
might not be detected, although losses during these earlier steps are usually minimized. Having
the carriers and tracers present during the sample decomposition provides an opportunity to
equilibrate the carrier or tracer with the sample so that the carrier, tracer, and analyte are in the
identical chemical form. While this can initially appear to be rather easy, in some cases it is
extremely difficult. The presence of multiple valence states and the formation of chemical
complexes are two conditions that introduce a host of equilibration problems (Section 14.2.2,
“Oxidation-Reduction Reactions”; Section 14.2.3, “Common Oxidation States”; and Section
14.2.4, “Oxidation State in Solution”). Crouthamel and Heinrich (1971) has an excellent
discussion of the intricacies and challenges associated with attaining true isotopic exchange:
“Fortunately, there are many reactions which have high exchange rates. This applies even
to many heterogeneous systems, as in the heterogeneous catalysis of certain electron
transfer reactions. In 1920, Hevesy, using ThB (212Pb), demonstrated the rapid exchange
between active lead nitrate and inactive lead chloride by the recrystallization of lead
chloride from the homogeneously mixed salts. The ionization of these salts leads to the
chemically identical lead ions, and a rapid isotopic exchange is expected. Similar
reversible reactions account for the majority of the rapid exchange reactions observed at
ordinary temperatures. Whenever possible, the analyst should conduct the isotope
exchange reaction through a known reversible reaction in a homogeneous system. The
true homogeneity of a system is not always obvious, particularly when dealing with the
very low concentrations of the carrier-free isotopes. Even the usually well-behaved alkalimetal ions in carrier-free solutions will adsorb on the surfaces of their containment
vessels or on colloidal and insoluble material in the solution. This is true especially in the
heavier alkali metals, rubidium and cesium. Cesium ions in aqueous solution have been
observed to absorb appreciably to the walls of glass vessels when the concentrations were
below 10!6 g/mL.”
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The reaction described above can be written as follows:
212

Pb(NO3)2(s) + PbCl2(s) 6 Pb(NO3)2 + 212PbCl2

Any of the following techniques may be employed to achieve both chemical and isotopic
equilibration:
• Careful adding, mixing, stirring, shaking, etc., to assure a homogeneous solution and prevent
layering.
• Introducing the carrier or tracer in several different chemical forms or oxidation states,
followed by oxidation or reduction to a single state.
• Treating the carrier or tracer and sample initially with strong oxidizing or reducing agents
during decomposition (e.g., wet ashing or fusion).
• Carrying out repeated series of oxidation-reduction reactions.
• Requiring that, at some point during the sample decomposition, all the species be together in
a clear solution.
Once a true equilibration between carrier or tracer and sample occurs, the radiochemistry
problem shifts from one of equilibration to that of separation from other elements, and ultimately
a good recovery of the radionuclide of interest.
Crouthamel and Heinrich (1971) summarize the introduction to equilibration (isotopic
exchange):
“Probably the best way to give the reader a feeling for the ways in which isotopic
exchange is achieved in practice is to note some specific examples from radiochemical
procedures. The elements which show strong tendencies to form radiocolloids in many
instances may be stabilized almost quantitatively as a particular complex species and
exchange effected. Zirconium, for example, is usually exchanged in strong nitric acidhydrofluoric acid solution. In this medium, virtually all the zirconium forms a ZrF6!2
complex. Niobium exchange is usually made in an oxalate or fluoride acid medium. The
exchange of ruthenium is accomplished through its maximum oxidation state, Ru(VIII)
which can be stabilized in a homogeneous solution and distilled as RuO4. Exchange may
also be achieved by cycling the carrier through oxidation and reduction steps in the
presence of the radioactive isotope. An iodine carrier with possible valence states of !1 to
+7 is usually cycled through its full oxidation-reduction range to ensure complete
exchange. In a large number of cases, isotopic exchange is not a difficult problem;
however, the analyst cannot afford to relax his attention to this important step. He must
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consider in each analysis the possibility of both the slow exchange of certain chemical
species in homogenous solution and the possible very slow exchange in heterogeneous
systems. In the latter case, this may consist simply of examining the solutions for
insoluble matter and taking the necessary steps to either dissolve or filter it and to assay
for possible radioactive content.”
Also see the discussion of equilibration of specific radionuclides in Section 14.10.9, “Review of
Specific Radionuclides.”
14.10.2

Oxidation State

Some radionuclides exist in solution in one oxidation state that does not change, regardless of the
kind of chemical treatment used for analysis. Cesium (Cs), radium, strontium, tritium ( 3H), and
thorium are in the +1, +2, +2, +1, and +4 oxidation states, respectively, during all phases of
chemical treatment. However, several radionuclides can exist in more than one state, and some
are notable for their tendency to exist in multiple states simultaneously, depending on the other
components present in the mixture. Among the former are cobalt, iron, iodine, and technetium,
and among the latter are americium, plutonium, and uranium. To ensure identical chemical
behavior during the analytical procedure, the radionuclide of interest and its carriers and/or
tracers in solution must be converted to identical oxidation states. The sample mixture containing
the carriers and/or tracer is treated with redox agents to convert each state initially present to the
same state, or to a mixture with the same ratio of states. Table 6E in Wahl and Bonner (1951)
provides a list of traditional agents for the oxidation and reduction of carrier-free tracers that is a
useful first guide to the selection of conditions for these radioequilibrium processes.
14.10.3

Hydrolysis

All metal ions (cations) in aqueous solution interact extensively with water, and, to a greater or
lesser extent, they exist as solvated cations (Katz et al., 1986):
Ra+2 + x@H2O 6 Ra(H2O)x+2
The more charged the cation, the greater is its interaction with water. Solvated cations, especially
those with +4, +3, and small +2 ions, tend to act as acids by hydrolyzing in solution. Simply
stated, hydrolysis is complexation where the ligand is the hydroxyl ion. To some extent, all metal
cations in solution undergo hydrolysis and exist as hydrated species. The hydrolysis reaction for a
metal ion is represented simply as (Choppin et al., 1995):
M+n + m@ H2O 6 M(OH)m+(n!m) + m@ H+1
Hydrolysis of the ferric ion (Fe+3) is a classical example:
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Fe+3 + H2O 6 Fe(OH)+2 + H+1
Considering the hydrated form of the cation, hydrolysis is represented by:
M(OH2)x+n 6 M(OH2)x!1(OH)(n!1)+ + H+1
In the latter equation, the hydrated complex ion associated with the hydroxide ion, is known as
the aquo-hydroxo species (Birkett et al., 1988). As each equation indicates, hydrolysis increases
the acidity of the solution, and the concentration of the hydrogen ion (pH) affects the position of
equilibrium. An increase in acidity (increase in H+1 concentration; decrease in pH) shifts the
position of equilibrium to the left, decreasing hydrolysis, while a decrease in acidity shifts it to
the right, increasing hydrolysis. The extent of hydrolysis, therefore, depends on the pH of the
solution containing the radionuclide. The extent of hydrolysis is also influenced by the radius and
charge of the cation (charge/radius ratio). Generally, a high ratio increases the tendency of a
cation to hydrolyze. A ratio that promotes hydrolysis is generally found in small cations with a
charge greater than one (Be+2, for example). The Th+4 cation, with a radius three times the size of
the beryllium ion but a +4 charge, is hydrolyzed extensively, even at a pH of four (Baes and
Mesmer, 1976). It is not surprising, therefore, that hydrolysis is an especially important factor in
the behavior of several metallic radionuclides in solution, and is observed in the transition,
lanthanide, and actinide groups. For the actinide series, the +4 cations have the greatest charge/
radius ratio and undergo hydrolysis most readily. Below pH 3, the hydrolysis of Th4+ is
negligible, but at higher pH, extensive hydrolysis occurs. Uranium (+4) undergoes hydrolysis in
solution at a pH above 2.9 with U(OH)3+ being the predominant hydrolyzed species. Neptunium
ions undergo hydrolysis in dilute acid conditions with evidence of polymer formation in acidic
solutions less than 0.3 M. The hydrolysis of plutonium is the most severe, often leading to
polymerization (see Section 14.10.4, “Polymerization”). In summary, the overall tendency of
actinides to hydrolyze decreases in the order (Katz et al., 1986):
An+4 > AnO2+2 > An+3 > AnO2+1
where “An” represents the general chemical symbol for an actinide.
For some cations, hydrolysis continues past the first reaction with water, increasing the number
of hydroxide ions (OH!1) associated with the cation in the aquo-hydroxo species:
U+4 + H2O 6 U(OH)+3 + H+1
U(OH)+3 +H2O 6 U(OH)2+2 + H+1
This process can, in some cases, conclude with the precipitation of an insoluble hydroxide, such
as ferric hydroxide. “Soluble hydrolysis products are especially important in systems where the
cation concentrations are relatively low, and hence the range of pH relatively wide over which
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such species can be present and can profoundly affect the chemical behavior of the metal” (Baes
and Mesmer, 1976).
Solutions containing trace concentrations of metallic radionuclides qualify as an example of
these systems. The form of hydrolysis products present can control important aspects of chemical
behavior such as (Baes and Mesmer, 1976):
•
•
•
•
•
•

Adsorption of the radionuclide on surfaces, especially on mineral and soil particles.
Tendency to coagulate colloidal particles.
Solubility of the hydroxide or metal oxide.
Extent of complex formation in solution.
Extent of extraction from solution by various reagents.
Ability to oxidize or reduce the radionuclide to another oxidation state.

Thus, a knowledge of the identity and stability of radionuclide ion hydrolysis products is
important in understanding or predicting the chemical behavior of trace quantities of radionuclides in solution (Baes and Mesmer, 1976). As the equilibrium equation indicates, H+1 is
produced as cations hydrolyze. Undesirable consequences of hydrolysis can, therefore, be
minimized or eliminated by the addition of acid to the analytical mixture to reverse hydrolysis or
prevent it from occurring. Numerous steps in radioanalytical procedures are performed at low pH
to eliminate hydrolytic effects. It is also important to know the major and minor constituents of
any sample, because hydrolysis effects are a function of pH and metal concentration. Thus,
maintaining the pH of a high iron-content soil sample below pH 3.0 is important, even if iron is
not the analyte.
14.10.4

Polymerization

The hydrolysis products of radionuclide cations described in the preceding section are
monomeric—containing only one metal ion. Some of these monomers can spontaneously form
polymeric metal hydroxo polymers in solution, represented by formation of the dimer (Birkett
et al., 1988):
2 M(H2O)x!1(OH)+(n!1) 6 [(H2O)x!2M(OH)2M(H2O)x!2]+2(n!1) + 2 H2O
The polymers contain -OH-bridges between the metal ions that, under high temperature,
prolonged aging, and/or high pH, can convert to -O-bridges, leading eventually to precipitation of
hydrated metal oxides. Birkett et al. (1988) states that:
“Formation of polymeric hydroxo species has been reported for most metals, although in
some cases, the predominant species in solution is the monomer. Some metals form only
dimers or trimers, while a few form much larger, higher-molecular-weight polymeric species.
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“Increasing the pH of a metal ion solution, by shifting the position of hydrolysis
equilibrium ..., results in an increased concentration of hydrolyzed species ..., which in turn
causes increased formation of polymeric species ... . Diluting a solution has two opposing
effects on the formation of polymeric species:
“(1)

Because dilution of acidic solutions causes a decrease in H+1 concentration (i.e.,
an increase in pH), it causes a shift in the hydrolyzed equilibrium toward
formation of hydrolyzed species.

“(2)

On the other hand, dilution decreases the ratio of polymeric to monomeric
complexes in solution. For metals that form both monomeric and polymeric
complexes, this means that monomeric species predominate beyond a certain level
of dilution.”

Because this type of polymerization begins with hydrolysis of a cation, minimizing or
eliminating polymerization can be achieved by the addition of acid to lower the pH of the
analytical solution to prevent hydrolysis (Section 14.10.3, “Hydrolysis”).
14.10.5

Complexation

Many radionuclides exist as metal ions in solution and have a tendency to form stable complex
ions with molecules or anions present as analytical reagents or impurities. The tendency to form
complex ions is, to a considerable extent, an expression of the same properties that lead to
hydrolysis; high positive charge on a +3 or +4 ion provides a strong driving force for the
interaction with ligands (Katz et al., 1986) (Section 14.3, “Complexation”).
Complex-ion formation by a radionuclide alters its form, introducing in solution additional
species of the radionuclide whose concentrations depend on the magnitude of the formation
constant(s). Alternate forms have different physical and chemical properties, and behave
differently in separation techniques, such as extraction or partition chromatography. The behavior
of alternate forms of radionuclides can present problems in the separation scheme that should be
avoided if possible or addressed in the protocol. Some separation schemes, however, take
advantage of the behavior of alternate radionuclide species formed by complexation, which can
alter the solubility of the radionuclides in a solvent or their bonding to an ion-exchange resin
(Section 14.3.4.2, “Separation by Solvent Extraction and Ion-Exchange Chromatography”).
14.10.6

Radiocolloid Interference

The tendency of some radionuclides in solution, particularly tracer levels of radionuclides, to
form radiocolloids, alters the physical and chemical behavior of those radionuclides (see Section
14.9.3.4, “Radiocolloid Formation”). Radioanalytical separations will not perform as expected in
solutions containing radiocolloids, particularly as the solubility of the radionuclide species
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decreases.
Solutions containing large molecules, such as polymeric metal hydrolysis products, are more
likely to form radiocolloids (Choppin et al., 1995). “If the solution is kept at sufficiently low pH
and extremely free of foreign particles, sorption and radiocolloid formation are usually avoided
as major problems” (Choppin et al. 1995). If tracer levels of radionuclides are present, trace
impurities become especially significant in the radiochemical procedure, and should be
minimized or avoided whenever possible (Crouthamel and Heinrich, 1971).
Crouthamel and Heinrich (1971) provide some specific insight into radiocolloidal interference in
the equilibration problem:
“The transition metals tend to form radiocolloids in solution, and in these heterogeneous
systems the isotopic exchange reaction between a radiocolloid and inactive carrier added to
the solution is sometimes slow and, more often, incomplete. Elements which show a strong
tendency to form radiocolloids, even in macro concentrations and acid solutions, are titanium,
zirconium, hafnium, niobium, tantalum, thorium, and protactinium, and, to a lesser degree,
the rare earths. Other metals also may form radiocolloids, but generally offer a wider choice
of valence states which may be stabilized in aqueous solutions”
14.10.7

Isotope Dilution Analysis

The basic concept of isotope dilution analysis is to measure the changes in specific activity of a
substance upon its incorporation into a system containing an unknown amount of that substance.
Friedlander et al. (1981), define specific activity:
“Specific activity is defined as the ratio of the number of radioactive atoms to the total
number of atoms of a given element in the sample (N*/N). In many cases where only the
ratios of specific activities are needed, quantities proportional to N */N, such as activity/mole,
are referred to as specific activity.”
Isotope dilution analysis uses a known amount of radionuclide to determine an unknown mass of
stable nuclide of the same element. For example, isotope dilution can be used to determine the
amount of some inactive material A in a system (Wang et al., 1975). To the system containing x
grams of an unknown weight of the inactive form of A, y grams of active material A* of known
activity D is added. The specific activity of the added active material, S1, is given by:
S1 = D/y
After ensuring isotopic exchange, the mixture of A and A* is isolated, but not necessarily
quantitatively, and purified. The specific activity, S2, is measured. Due to the conservation of
matter,
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S2 = D / (x + y)
and by substituting for S1y for D and rearranging, the amount x of inactive A is given as
x = y (S1/S2 ! 1)
However, this equation is valid only if complete isotopic exchange has occurred, a task not
always easy to achieve.
14.10.8

Masking and Demasking

Masking is the prevention of reactions that are normally expected to occur through the presence
or addition of a masking reagent. Masking reactions can be represented by the general reversible
equation:
A + Ms 6 A @ Ms
where A is the normal reacting molecule or ion, and Ms is the masking agent. The decreased
concentration of A at equilibrium determines the efficiency of masking. An excess of masking
agent favors the completeness of masking, as expected from LeChatelier’s Principle. Feigl (1936)
has described masking reagent and the masking of a reaction:
“... the concentration of a given ion in a solution can be so diminished by the addition of
substances which unite with the ion to form complex salts that an ion product sufficient to
form a precipitate or cause a color reaction is no longer obtained. Thus we speak of the
masking of a reaction and call the reagent responsible for the disappearance of the ions
necessary for the reaction, the masking reagent.”
The concepts of masking and demasking are discussed further in Perrin (1979) and in Dean
(1995).
Masking techniques are frequently used in analytical chemistry because they often provide
convenient and efficient methods to avoid the effects of unwanted components of a system
without having to separate the interferent physically. Therefore, the selectivity of many analytical
techniques can be increased through masking techniques. For example, copper can be prohibited
from carrying on ferric hydroxide at pH 7 by the addition of ammonium ions to complex the
copper ions. Fe3+ and Al3+ both interfere with the extraction of the +3 actinides and lanthanides in
some systems, but Fe3+ can be easily masked through reduction with ascorbic acid, and Al3+ can
be masked through complexation with fluoride ion (Horwitz et al., 1993 and 1994). In another
example, uranium can be isolated on a U/TEVA® column (Eichrom Technologies, Inc., Darien,
IL) from nitric acid solutions by masking the tetravalent actinides with oxalic acid; the tetravalent
actinides are complexed and pass through the column, whereas uranium is extracted (SpecNews,
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1993). Strontium and barium can be isolated from other metals by cation exchange from a solution of water, pyridine, acetic acid and glycolic acid. The other metals form neutral or negative
complexes and pass through the cation column, while strontium and barium are retained
(Orlandini, 1972). Masking phenomena are present in natural systems as well. It has been
demonstrated that humic and fulvic acids can complex heavy metals such that they are no longer
bioavailable and are, therefore, not taken up by plants. Tables 14.16 and 14.17 list common
masking agents.
TABLE 14.16 — Masking agents for ions of various metals
Metal Masking Agent
Ag
Al
As
Au
Ba
Be
Bi
Ca
Cd
Cs
Co
Cr
Cu

Fe

Ga
Ge
Hf
Hg
In
Ir

Br!, citrate, Cl!, CN!, I!, NH3, SCN! S2 O3!2, thiourea, thioglycolic acid, diethyldithiocarbamate,
thiosemicarbazide, bis(2!hydroxyethyl)dithiocarbamate
Acetate, acetylacetone, BF4!, citrate, C2O4!2, EDTA, F!, formate, 8-hydroxyquinoline-5-sulfonic acid,
mannitol, 2,3-mercaptopropanol, OH!, salicylate, sulfosalicylate, tartrate, triethanolamine, tiron
Citrate, 2,3-dimercaptopropanol, NH2OH.HCl, OH!, S2!2, tartrate
Br!, CN!, NH3, SCN!, S2O3!2, thiourea
Citrate, cyclohexanediaminetetraacetic acid, N,N-dihydroxyethylglycine, EDTA, F!, SO4!2, tartrate
Acetylacetone, citrate, EDTA, F!, sulfosalicylate, tartrate
Citrate, Cl!, 2,3-dimercaptopropanol, dithizone, EDTA, I!, OH!, Na5P3O10, SCN!, tartrate, thiosulfate,
thiourea, triethanolamine
BF4!, citrate, N,N-dihydroxyethylglycine, EDTA, F!, polyphosphates, tartrate
Citrate, CN!, 2,3-dimercaptopropanol, dimercaptosuccinic acid, dithizone, EDTA, glycine, I!, malonate,
NH3, 1,10-phenanthroline, SCN!, S2O3!2, tartrate
Citrate, N,N-dihydroxyethylglycine, EDTA, F!, PO4!3, reducing agents (ascorbic acid), tartrate, tiron
Citrate, CN!, diethyldithiocarbamate, 2,3!dimercaptopropanol, dimethylglyoxime, ethylenediamine,
EDTA, F!, glycine, H2O2, NH3, NO2!, 1,10-phenanthroline, Na5P3O10, SCN!, S2O3!2 tartrate
Acetate, (reduction with) ascorbic acid + KI, citrate, N,N-dihydroxyethylglycine, EDTA, F!, formate,
NaOH + H2O2, oxidation to CrO4!2, Na5P3O10, sulfosalicylate, tartrate, triethylamine, tiron
Ascorbic acid + KI, citrate, CN!, diethyldithiocarbamate, 2,3-dimercaptopropanol, ethylenediamine,
EDTA, glycine, hexacyanocobalt(III)(3!), hydrazine, I!, NaH2PO2, NH2OH.HCl, NH3, NO!2, 1,10phenanthroline, S!2, SCN! + SO3!2, sulfosalicylate, tartrate, thioglycolic acid, thiosemicarbazide,
thiocarbohydrazide, thiourea
Acetylacetone, (reduction with) ascorbic acid, C2O4!2, citrate, CN! 2,3-dimercaptopropanol, EDTA, F!,
NH3, NH2OH.HCl, OH!, oxine 1,10-phenanthroline, 2,2'-bipyridyl, PO4!3, P2O7!4, S!2, SCN!, SnCl2,
S2O3!2, sulfamic acid, sulfosalicylate, tartrate, thioglycolic acid, thiourea, tiron, triethanolamine,
trithiocarbonate
Citrate, Cl!, EDTA, OH!, oxalate, sulfosalicylate, tartrate
F!, oxalate, tartrate
See Zr
Acetone, (reduction with) ascorbic acid, citrate, Cl!, CN!, 2,3-dimercaptopropan-1-ol, EDTA, formate, I!,
SCN!, SO3!2, tartrate, thiosemicarbazide, thiourea, triethanolamine
Cl!, EDTA, F!, SCN!, tartrate thiourea, triethanolamine
Citrate, CN!, SCN!, tartrate, thiourea

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Metal Masking Agent
La
Mg
Mn

Mo
Nb
Nd
NH4+
Ni
Np
Os
Pa
Pb
Pd
Pt
Pu
Rare
Earths
Re
Rh
Ru
Sb
Sc
Se
Sn
Ta
Te
Th
Ti
Tl
U
V

Citrate, EDTA, F!, oxalate, tartrate, tiron
Citrate, C2O4!2, cyclohexane-1,2-diaminetetraacetic acid, N,N-dihydroxyethylglycine, EDTA, F!, glycol,
hexametaphosphate, OH!, P2O7!4, triethanolamine
Citrate, CN!, C2O4!2, 2,3!dimercaptopropanol, EDTA, F!, Na5P3O10, oxidation to MnO4!, P2O7!4,
reduction to Mn+2 with NH2OH.HCl or hydrazine, sulfosalicylate, tartrate, triethanolamine, triphosphate,
tiron
Acetylacetone, ascorbic acid, citrate, C2O4!2, EDTA, F!, H2O2, hydrazine, mannitol, Na5P3O10,
NH2OH.HCl, oxidation to molybdate, SCN!, tartrate, tiron, triphosphate
Citrate, C2O4!2, F!, H2O2, OH!, tartrate
EDTA
HCHO
Citrate, CN!, N,N-dihydroxyethylglycine, dimethylglyoxime, EDTA, F!, glycine, malonate, Na5P3O10,
NH3 1,10-phenanthroline, SCN!, sulfosalicylate, thioglycolic acid, triethanolamine, tartrate
F!
CN!, SCN!, thiourea
H2O2
Acetate, (C6H5)4AsCl, citrate, 2,3-dimercaptopropanol, EDTA, I!, Na5P3O10, SO4!2, S2O3!2, tartrate, tiron,
tetraphenylarsonium chloride, triethanolamine, thioglycolic acid
Acetylacetone, citrate, CN!, EDTA, I!, NH3, NO2!, SCN!, S2O3!2, tartrate, triethanol-amine
Citrate, CN!, EDTA, I!, NH3, NO2!, SCN!, S2O3!2, tartrate, urea
Reduction to Pu+4 with sulfamic acid
C2O4!2, citrate, EDTA, F!, tartrate
Oxidation to perrhenate
Citrate, tartrate, thiourea
CN!, thiourea
Citrate, 2,3-dimercaptopropanol, EDTA, I!, OH!, oxalate, S!2, S2!2, S2O3!2, tartrate, triethanolamine
Cyclohexane-1,2-diaminetetraacetic acid, F!, tartrate
Citrate, F!, I!, reducing agents, S!2, SO3!2, tartrate
Citrate, C2O3!2, 2,3-dimercaptopropanol, EDTA, F!, I!, OH!, oxidation with bromine water, PO4!3,
tartrate, triethanolamine, thioglycolic acid
Citrate, F!, H2O2, OH!, oxalate, tartrate
Citrate, F!, I!, reducing agents, S!2, sulfite, tartrate
Acetate, acetylacetone, citrate, EDTA, F!, SO4!2, 4-sulfobenzenearsonic acid, sulfosalicylic acid, tartrate,
triethanolamine
Ascorbic acid, citrate, F!, gluconate, H2O2, mannitol, Na5P3O10, OH!, SO4!2, sulfosalicylic, acid, tartrate,
triethanolamine, tiron
Citrate, Cl!, CN!, EDTA, HCHO, hydrazine, NH2OH.HCl, oxalate, tartrate, triethanolamine
Citrate, (NH4)2CO3, C2O4!2, EDTA, F!, H2O2, hydrazine + triethanolamine, PO4!3, tartrate
(reduction with) Ascorbic acid, hydrazine, or NH2OH.HCl, CN!, EDTA, H2O2, mannitol, oxidation to
vanadate, triethanolamine, tiron

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Metal Masking Agent
W
Y
Zn
Zr

Citrate, F!, H2O2, hydrazine, Na5P3O10, NH2OH.HCl, oxalate, SCN!, tartrate, tiron, triphosphate, oxidation
to tungstate
Cyclohexane-1,2-diaminetetraacetic acid, F!
Citrate, CN!, N,N!dihydroxyethylglycine, 2,3-dimercaptopropanol, dithizone, EDTA, F!, glycerol, glycol,
hexacyanoferrate(II)(4!), Na5P3O10, NH3, OH!, SCN!, tartrate, triethanolamine
Arsenazo, carbonate, citrate, C2O!2, cyclohexane-1,2-diaminetetraacetic acid, EDTA, F!, H2O2, PO4!3,
P2O7-4, pyrogallol, quinalizarinesulfonic acid, salicylate, SO4!2 + H2O2, sulfosalicylate, tartrate,
triethanolamine

Sources: Perrin (1979) and Dean (1995)

TABLE 14.17 — Masking agents for anions and neutral molecules
Anion or
Neutral
Molecule

Masking Agent

Boric Acid
Br!
Br2
BrO3!
Chromate(VI)
Citrate
Cl!
Cl2
ClO3!
ClO4!
CN!
EDTA
F!
Fe(CN)3!3
Germanic Acid
I!
I2
IO3!
IO4!
MnO4!
MoO4!2
NO2!
Oxalate
Phosphate
S
S!2
Sulfate
Sulfite
SO6!2
Se and its anions
TeI!

F!, glycol, mannitol, tartrate, and other hydroxy acids
Hg+2
Phenol, sulfosalicylic acid
Reduction with AsO4!5, hydrazine, sulfite, or thiosulfate
Reduction with AsO4!5, ascorbic acid, hydrazine, hydroxylamine, sulfite, or thiosulfate
Ca+2
Hg+2, Sb+3
Sulfite
Thiosulfate
Hydrazine, sulfite
HCHO, Hg+2, transition-metal ions
Cu+2
Al+3, Be+2, boric acid, Fe+3, Th+4, Ti+4, Zr+4
AsO4!5, ascorbic acid, hydrazine, hydroxylamine, thiosulfate
Glucose, glycerol, mannitol
Hg+2
Thiosulfate
Hydrazine, sulfite, thiosulfate
AsO4!5, hydrazine, molybdate(VI), sulfite, thiosulfate
Reduction with AsO4!5, ascorbic acid, azide, hydrazine, hydroxylamine, oxalic acid, sulfite, or
thiosulfate
Citrate, F!, H2O2, oxalate, thiocyanate + Sn+2
Co+2, sulfamic acid, sulfanilic acid, urea
Molybdate(VI), permanganate, Al+3
Fe+3, tartrate
CN!, S2!, sulfite
Permanganate + sulfuric acid, sulfur
Cr+3 + heat
HCHO, Hg+2, permanganate + sulfuric acid
Ascorbic acid, hydroxylamine, thiosulfate
Diaminobenzidine, sulfide, sulfite

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Anion or
Neutral
Molecule

Masking Agent

Tungstate
Vanadate

Citrate, tartrate
Tartrate

Sources: Perrin (1979) and Dean (1995)

Demasking refers to any procedure that eliminates the effect of a masking agent already present
in solution. There are a variety of methods for demasking, including changing the pH of the
solution and physically removing, destroying, or displacing the masking agent. The stability of
most metal complexes depends on pH, so simply raising or lowering the pH is frequently
sufficient for demasking. Another approach to demasking involves the formation of new
complexes or compounds that are more stable than the masked species. For example, boric acid
commonly is used to demask the fluoride complexes of Sn4+ or Mo6+, and hydroxide is used to
demask the thiocyanate complexes of Fe3+. In addition, it might be possible to destroy the
masking agent in solution through a chemical reaction (i.e., through the oxidation of EDTA in
acidic solutions by permanganate or another strong oxidizing agent).
14.10.9

Review of Specific Radionuclides

The analytical separation and analysis of radionuclides involves several scientific disciplines.
The decay of one radionuclide to another is referred to as “radioactive equilibrium.” A series of
mathematical expressions (derived from the Bateman equations, Friedlander et al., 1981) identify
three separate cases of these equilibria (see Attachment 14A, “Radioactive Decay and
Equilibrium”).
14.10.9.1 Americium
Americium is a metal of the actinide series which is produced synthetically by neutron activation
of uranium or plutonium followed by beta decay.
Isotopes
Twenty isotopes of americium are known, 232Am through 248Am, including three metastable
states. All isotopes are radioactive. Americium-243 and 241Am, alpha emitters, are the longest
lived with half-lives of 7,380 years and 432.7 years, respectively. Americium-241 and 243Am also
undergo spontaneous fission. Americium-242m has a half-life of 141 years, and the half-lives of
the remaining isotopes are measured in hours, minutes, or seconds. Americium-241 is the most
common isotope of environmental concern.

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Occurrence
None of the isotopes of americium occur naturally. It is produced synthetically by neutron
bombardment of 238U or 239Pu followed by beta decay of the unstable intermediates. Americium241 is found in various plutonium wastes and can be extracted from reactor wastes. Some
industrial ionization sources also contain americium. Decay of 241Pu injected in the atmosphere
during weapons testing contributes to the presence of 241Am.
The silver metal is prepared by reduction of americium fluoride (AmF 3) or americium oxide
(AmO2) with active metals at high temperatures and is purified by fractional distillation, taking
advantage of its exceptionally high vapor pressure compared to other transuranium elements.
Kilogram quantities of 241Am are available, but only 10 to 100 g quantities of 243Am are prepared.
Soft gamma emission from 241Am is used to measure the thickness of metal sheets and metal
coatings, the degree of soil compaction, sediment concentration in streams, and to induce X-ray
fluorescence in chemical analysis. As an alpha emitter, it is mixed with beryllium to produce a
neutron source for oil-well logging and to measure water content in soils and industrial process
streams. The alpha source is also used to eliminate static electricity and as an ionization source in
smoke detectors.
Solubility of Compounds
Among the soluble salts are the nitrate, halides, sulfate, and chlorate of americium (Am+3). The
fluoride, hydroxide, and oxalate are insoluble. The phosphate and iodate are moderately soluble
in acid solution. Americium(VI) is precipitated with sodium acetate to produce the hydrate,
NaAmO2(C2H3O2)3@ xH2O.
Review of Properties
The study of the properties of americium is very difficult because of the intense alpha radiation
emitted by 241Am and 243Am, but some properties are known. Americium metal is very ductile
and malleable but highly reactive and unstable in air, forming the oxide. It is considered to be a
slightly more active metal than plutonium and is highly reactive combing directly with oxygen,
hydrogen, and halides to form the respective compounds, AmO2, AmH3, and AmX3. Alloys of
americium with platinum, palladium, and iridium have been prepared by hydrogen reduction of
americium oxide in the presence of the finely divided metals.
Unless the transuranium elements are associated with high-level gamma emission, the principal
toxicological problems associated with the radionuclides are the result of internal exposure after
inhalation or ingestion. When inhaled or ingested, they are about equally distributed between
bone tissue and the liver. At high doses transuranics lead to malignant tumors years later. In
addition, large quantities of 241Am could conceivably lead to criticality problems, producing
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external radiation hazards or neutron exposure from (α,n) reactions. Americium-241 is also a
gamma emitter.
Americium is generally thought to be adsorbed by many common minerals at pH values found in
the environment. Complexation of Am+3 by naturally occurring ligands, however, would be
expected to strongly reduce its adsorption.
Solution Chemistry
Americium can exist in solution in the +3, +4, (V), and (VI) oxidation states. Simple aqueous
ions of Am+3 and AmO2+2 (VI oxidation state) are stable in dilute acid, but Am+3 is the
predominant oxidation state. Free radicals produced by radiolysis of water by alpha particles
reduce the higher states spontaneously to Am+3. The +3 oxidation state exists as Am(OH)3 in
alkaline solution. Simple tetravalent americium is unstable in mineral acid solutions, disproportionating rapidly to produce Am+3 and AmO2+1 [Am(V)] in nitric and perchloric acid solutions. In
contrast, dissociation of Am(OH)4 or AmO2 [both Am+4] in sulfuric acid solutions produces
solutions containing Am+3 and AmO2+2. Stability is provided by complexation with fluoride ions
and oxygen-containing ligands such as carbonate and phosphate ions. The AmO2+1 ion also
disproportionates in acid solutions to yield Am+3 and AmO2+2, but the process for 241Am is so
slow that radiation-induced reduction dominates. Evidence exists for the presence of Am(VII) in
alkaline solutions from the oxidation of AmO2+2.
OXIDATION-REDUCTION BEHAVIOR. Although disproportionation reactions convert the +4 and
(V) oxidation states into the +3 and (VI) states, radiolysis eventually converts the higher
oxidation state into Am+3. Redox processes are used, however, to produce solutions of alternate
oxidation states and to equilibrate the forms of americium into a common state, usually +3, but
sometimes (VI).
The +4 state is reduced to Am+3 by iodide. In dilute, nonreducing solutions, peroxydisulfate
(S2O8!2) oxidize both the +3 and (V) states to the (VI) state. Ce+4 and ozone (O3) oxidize the (V)
state to (VI) in perchloric acid solution. Electrolytic oxidation of Am+3 to AmO2+2 occurs in
phosphoric, nitric, and perchloric acid solutions and solutions of sodium bicarbonate (Na2CO3).
The latter ion is reduced to Am+3 by iodide, hydrogen peroxide, and the nitrite ion (NO2!1).
COMPLEXATION. The +3 oxidation state forms complexes in the following order of strength (in
aqueous solution): F! > H2PO4! > SCN! > NO3! > Cl!. Both Am+3 and Am+4 form complexes
with organic chelants. These are stable in aqueous and organic solvents. Americium (+4) can be
easily reduced unless special oxidizing conditions are maintained. The AmO2+2 ion also forms
significant complex ions with nitrate, sulfate, and fluoride ions.
HYDROLYSIS. The actinide elements are known for their tendency to hydrolyze and, in many
cases, form insoluble polymers. In the predominant +3 oxidation state in solution, americium,
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with its large radius, has the least tendency of the +3 actinides to hydrolyze; yet, hydrolysis is
expected to occur with some polymerization. Hydrolysis that does occur is complicated and
depends on the nature of the cations present and may start at pH values as low as 0.5–1.0. In
contrast, the AmO2+2, like all actinyl ions, undergoes hydrolysis to an appreciable extent. The
tendency to form polymers of colloidal dimensions, however, appears to be small relative to
other actinide ions in the (VI) oxidation state. Precipitation occurs early on after relatively small
polymeric aggregates form in solution. The strong tendency to form insoluble precipitates after a
small amount of hydrolysis makes characterization of the water-soluble polymers a difficult
problem.
RADIOCOLLOIDS. At trace concentrations, a colloidal form of Am+2 can easily be prepared, so
steps should be taken to avoid its formation during analytical procedures. At high pH ranges,
colloids form from the Am(OH)3, and at lower pH ranges through adsorption of Am+3 onto
foreign particles. Their formation depends on storage time, pH, and ionic strength of the solution.
Dissolution of Samples
Americium is generally dissolved from irradiated reactor fuels, research compounds, and soil,
vegetation, and biological samples. Spent fuel elements may be difficult to dissolve but eventually yield to digestion with hydrofluoric acid, nitric acid, or sulfuric acid. Aqua regia is used if
platinum is present, and hydrochloric acid with an oxidizing agent such as sodium chlorate.
Perchloric acid, while a good solvent for uranium, reacts too vigorously. Sodium hydroxideperoxide is a good basic solvent. Research compounds, usually salts, yield to hot concentrated
nitric or sulfuric acid. Soil samples are digested with concentrated nitric acid, hydrofluoric acid,
or hydrochloric acid. Vegetation and biological samples are commonly wet ashed, and the
residue is treated with nitric acid.
Separation Methods
The separation of americium, particularly from other transuranics, is facilitated by the
exceptional stability of Am+3 compared to the trivalent ions of other actinides, which more
readily convert to higher oxidation states under conditions that americium remains trivalent.
PRECIPITATION AND COPRECIPITATION. Coprecipitation with lanthanum fluoride (LaF3) is
achieved after reduction of higher oxidation states to Am+3. Select oxidation of other transuranic
elements such as neptunium and plutonium to the +4 or VI oxidation states solubilizes these
radionuclides leaving americium in the insoluble form. Although coprecipitation with rare earths
as fluorides or hydroxides from a bicarbonate solution of americium(VI), is used to purify
americium, it is not as effective as ion-exchange procedures. Other coprecipitating agents for
Am+3 include thorium oxalate [Th(C2O4)2], calcium oxalate (CaC2O4), ferric hydroxide
[Fe(OH)3), and lanthanum potassium sulfate [LaK(SO4)2]. Americium (+4) is also coprecipitated
with these reagents as well as with zirconium phosphate [Zr3(PO4)4]. Americium(VI) is not
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coprecipitated with any of these reagents but with sodium uranyl acetate [NaUO2(C2H3O2)2].
SOLVENT EXTRACTION. Organic solvents and chelating agents are available for separating
americium from other radionuclides by selectively extracting either americium or the alternate
radionuclide from aqueous solutions into an organic phase. Tributylphosphate (TBP) in kerosene
or TTA in xylene removes most oxidation states of neptunium and plutonium from Am+3 in the
presence of dilute nitric acid. The addition of sodium nitrate (6 M) tends to reverse the trend
making americium more soluble in TBP than uranium, neptunium, or plutonium radionuclides.
Bis(2-ethylhexyl) phosphoric acid (HDEHP) in toluene is highly effective in extracting Am+3 and
is used in sample preparation for alpha spectroscopic analysis.
Plutonium in the +4 oxidation state can interfere with Am analysis. See Section 14.10.9.8 on
plutonium for a discussion of how to separate americium from plutonium.
ION EXCHANGE. Separation of americium can be achieved by cation-exchange chromatography.
Any of its oxidation states exchange with a cation resin in dilute acid solution, but the higher
oxidation states are not important in cation-exchange separations because they are unstable
toward reduction to the +3 state. Generally, Am+3 is the last tripositive ion among the actinides
eluted from a cation-exchange matrix, although the order may not be maintained under all
conditions. Many eluting agents are available for specific separations. Concentrated hydrochloric
acid, for example, has been used for separating actinides such as americium from the lanthanides.
Anion-exchange chromatography has been widely used for separating americium. Anionic
complexes of Am+3 form at high chloride concentrations, providing a chemical form that is easily
exchanged on an anion-exchange column. The column can be eluted using dilute hydrochloric
acid or a dilute hydrochloric acid/ammonium thiocyanate solution. Anion-exchange separations
of americium are also realized with columns prepared with concentrated nitric acid solutions.
The sequential separation of the actinides is accomplished readily using anion-exchange
chromatography. Americium, plutonium, neptunium, thorium, protactinium, curium, and
uranium can all be separated by the proper application of select acid or salt solutions to the
column.
ELECTRODEPOSITION. Americium can be electrodeposited for alpha spectrometry measurement
on a highly polished platinum cathode. The sample is dissolved in a dilute hydrochloric acid
solution that has been adjusted to a pH of about six with ammonium hydroxide solution using
methyl red indicator. The process runs for one hour at 1.2 amps.
Methods of Analysis
Americium-241 is detected and quantified by alpha or gamma spectrometry, or by gas
proportional counting (GPC). Trace quantities of 241Am are analyzed by GPC, after separation
from interfering radionuclides by solvent extraction, coprecipitation, or ion-exchange
chromatography. The isolated radionuclide is collected and mounted on a planchet or
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electroplated onto a platinum electrode for counting by alpha spectrometry. Americium-243 is
added to the analytical solution as a tracer to measure chemical yield. Americium-241 may be
determined directly (i.e., no radiochemical separation) in bulk soil samples by gamma
spectroscopy.
Compiled from: Ahrland, 1986; Baes and Mesmer, 1976; Choppin et al., 1995; Considine
and Considine, 1983; Cotton and Wilkinson, 1988; DOE, 1990 and 1997, 1995; 1997;
Ehmann and Vance, 1991; Greenwood and Earnshaw, 1984; Haissinsky and Adolff, 1965;
Horwitz et al., 1993, 1995; Katz et al., 1986; Lindsay, 1988; Metz and Waterbury, 1962;
NEA, 1982; SCA, 2001; Penneman, 1994; Penneman and Keenan, 1960; Schulz and
Penneman, 1986; Seaborg and Loveland, 1990.
14.10.9.2 Carbon
The chemistry of carbon compounds is too extensive to be summarized here. Fortunately, only
one isotope of carbon, 14C, is significant in analytical separation. This chapter will focus on the
two principal radioisotopes of carbon that are in use: 11C and 14C.
Isotopes
Carbon-11 has a half-life of 20 minutes. It is used for medical diagnoses and is prepared by
proton bombardment of a boron target in an accelerator. The 11C in the target then may be
incorporated as part of a tracer molecule that would be used for the diagnosis. This isotope is also
formed in nuclear reactors by the two reactions, 11B(p, n) 11C and 12C(n, 2n)11C.
The chemical environment in the reactor coolant system is highly reducing (overpressure of
hydrogen gas is used to minimize oxygen formation from radiolysis of water). Thus, the chemical
form of the carbon is most likely 11CH4. The radioisotope decays to 11B by positron emission. It
may be detected by liquid scintillation or gamma ray detection of the 511 keV annihilation peak.
Its short half-life obviates the need for its environmental analysis.
Carbon-14 is also formed as a result of activation in reactor coolant systems of fission reactors
from the following reaction: 17O(n,α)14C. As with 11C, the chemical form will most likely be
14
CH4.
Occurrence
Carbon-14 is a naturally occurring radionuclide with a half-life of 5,720 years. It is formed as a
result of 14N(n, p)14C. The nitrogen atoms in the upper atmosphere are bombarded with highenergy neutrons emitted from the sun. The carbon becomes incorporated as part of a CO2
molecule due to the presence of oxygen and many highly energetic particles and free radicals in
the upper atmosphere. Carbon dioxide freely exchanges with all carbon using organisms in the
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environment. The living organism rapidly reaches a state of equilibrium with the environment
because of the long half-life of the carbon. The rate of radioactive decay of naturally occurring
14
C is approximately 780 Bq (13 dpm) per gram of total carbon. However, once an organism dies,
it ceases to exchange that carbon with the environment. Thus, the activity per gram of carbon
would decrease with the characteristic half-life of 14C (as long as the material is undisturbed).
This is the basis for carbon dating of materials.
Solubility and Solution Chemistry
Organic compounds have a vast range of chemical and physical properties. Many of the 14C
containing materials one encounters will be insoluble in aqueous solution, but soluble in some
organic solvents. Carbon is basically tetravalent in all compounds, and forms covalent bonds.
Thus, when using separation techniques involving a carrier, such as CO3-2, it is necessary to
ensure not only that the sample is dissolved, but that sufficient oxidative power has been
employed to convert the analyte to the same chemical form. Carbon is also unique in that CO2 is
a common oxidation product of carbon and can easily escape from solution. The equilibria
CO2 + H2O 6 HCO3!1 + H+
HCO3!1 + H2O 6 CO3!2 + H+
demonstrate the significant effect that acid concentration can have on the loss of carbon, as CO2,
from solution. This must be taken into consideration whenever processing 14C samples.
Dissolution
Many applications involve 14C as tracers. As discussed later, no sample dissolution may be
needed and analysis by one of the two analytical techniques may proceed directly.
Dissolution of samples containing 14C where other isotopes are present involves the complete
destruction of the organic matter in the sample, and simultaneously not allowing the volatilization of the carbon. This is most commonly achieved by permanganate oxidation in a basic
solution. As seen in the equilibrium equations for carbon, in basic solution it is present as the
CO3!2 species, which is nonvolatile.
Samples also may be prepared by high temperature oxidation, in which the carbon is converted to
CO2. The exit gasses from the combustion process must be directed through a trap which will
remove carbon dioxide. These include such materials as molecular sieve, barium chloride
solutions or Ascarite® columns.

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Methods of Analysis
Carbon-14 decays only by β- emission. The Eβmax of this emission is 0.156 MeV. Although it is
detectable by gas proportional counting, the only two methods of analysis commonly used for
this isotope are liquid scintillation and mass spectroscopic analysis. The methods for liquid
scintillation analyses are described in Chapter 15, Quantification of Radionuclides, and Kessler
(1989).
14.10.9.3 Cesium
Cesium is the last member of the naturally occurring alkali metals in Group IA of the Periodic
Table, with an atomic number of 55. Its radiochemistry is simplified because the Group IA
metals form only +1 ions. Elemental cesium is a very soft, silver-white metallic solid in the pure
state with a melting point of only 28.5 EC. It tarnishes quickly to a golden-yellow color when
exposed to small amounts of air. With sufficient air, it ignites spontaneously. It is normally
stored under xylente or toluene to prevent contact with air.
Isotopes
Cesium isotopes of mass number 112 to 148 have been identified. Cesium-133 is the only stable
isotope. Cesium-134, 136Cs and 137Cs are the only isotopes of significance from an environmental
perspective. They are formed from the nuclear fission process. Their half-lives are 2.06 years,
13.2 days, and 30.17 years, respectively. Cesium-135 also is formed as a result of the fission
process. However, it is not a significant isotope, because it is a low-energy (0.21 MeV) beta-only
emitter with a long half-life (2.2×106 years).
Occurrence
Cesium is widely distributed in the Earth’s crust with other alkali metals. In granite and
sedimentary rocks the concentration is less than 7 ppm. In seawater it is about 0.002 ppm, but in
mineral springs the concentration may be greater than 9 mg/L. Cesium-137 is produced in
nuclear fission and occurs in atmospheric debris from weapons tests and accidents. It is a very
important component of radioactive fallout; and because of its moderately long half-life and high
solubility, it is a major source of long-lived external gamma radiation from fallout. It accounts
for 30 percent of the gamma activity of fission products stored for one year, 70 percent in two
years, and 100 percent after five years.
Cesium metal’s most recognized use is in the atomic clock that serves to define the second.
Cesium has been considered as a fuel in ion-propulsion engines for deep space travel and as a
heat-transfer medium for some applications. Cesium-137 has replaced 60Co in the treatment of
cancer and has been used in industrial radiography for the control of welds. Cesium-137 is also
used commercially as a sealed source in liquid scintillation spectrometers. The 661 keV gamma
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ray it emits is used to create an electron (Compton effect) distribution, which allows the degree
of sample quench to be determined.
Solubility of Compounds
Most cesium salts are very soluble in water and dilute acids. Among the salts of common anions,
the notable exceptions are cesium perchlorate and periodate (CsClO 4 and CsIO4). Several cesium
compounds of large anions are insoluble. Examples include the following: silicotungstate
[Cs8SiW12O42], permanganate (CsMnO4), chloroplatinate (Cs2PtCl6), tetraphenylborate
[CsB(C6H5)4], alum [CsAl(SO4)2], and cobaltnitrate complex [Cs3Co(NO3)6].
Review of Properties
Cesium is the most active and electropositive of all the metals. It forms compounds with most
inorganic and organic anions; it readily forms alums with all the trivalent cations that are found
in alums. The metal readily ionizes, and in ammonia solutions it is a powerful reducing agent.
When exposed to moist air, it tarnishes initially forming oxides and a nitride and then quickly
melts or bursts into flame. With water the reaction is violent. Cesium reacts vigorously with
halogens and oxygen, and it is exceptional among the alkali metals in that it can form stable
polyhalides such as CsI3. Reaction with oxygen forms a mixture of oxides: cesium oxide (Cs2O),
cesium peroxide (Cs2O2), and cesium superoxide (CsO2). The toxicity of cesium compounds is
generally not important unless combined with another toxic ion.
Cesium-137, introduced into the water environment as cations, is attached to soil particles and
can be removed by erosion and runoff. However, soil sediment particles act as sinks for 137Cs,
and the radionuclide is almost irreversibly bound to mica and clay minerals in freshwater
environments. It is unlikely that 137Cs will be removed from these sediments under typical
environmental conditions. Solutions of high ionic strength as occur in estuarine environments
might provide sufficient exchange character to cause cesium to become mobile in the ecosphere.
Solution Chemistry
The cesium ion exists in only the +1 oxidation state, and its solution chemistry is not complicated
by oxidation-reduction reactions. As a result, it undergoes complete, rapid exchange with carriers
in solution. The cesium ion is colorless in solution and is probably hydrated as a hexaaquo
complex.
COMPLEXATION. Cesium ions form very few complex ions in solution. The few that form are
primarily with nitrogen-donor ligands or beta-diketones. Anhydrous beta-diketones are insoluble
in water, but in the presence of additional coordinating agents, including water, they become
soluble in hydrocarbons. One solvent-extraction procedure from aqueous solutions is based on
chelation of cesium with TTA in hydrocarbon solvents. Cesium is sandwiched between crown
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ligands, associated with the oxygen atoms of the ether, in [Cs9(18-C-6)14]+9.
HYDROLYSIS. With the small charge and large radius of the cesium ion, hydrolysis reactions are
inconsequential.
ADSORPTION. When cesium is present in extremely low concentrations, even in the presence of 2
M acid, adsorption on the walls of glass and plastic containers leads to complications for the
radioanalyst. Half the activity of cesium radionuclides, for example, can be lost from acid
solutions stored for one month in these containers. Experiments indicate that addition of 1 µg
cesium carrier per milliliter of solution is sufficient to stabilize acid solutions for six months.
Dissolution of Samples
Radiochemists generally dissolve cesium samples from irradiated nuclear fuel, activated cesium
salts, natural water, organic material, agriculture material, and soils. Nuclear fuel samples are
generally dissolved in HCl, HNO3, HF, or a combination of these acids. Care should be taken to
ensure that the sample is representative if 137Cs has been used as a burn-up monitor. Precautions
should also be taken with these samples to prevent loss of cesium because of leaching or incomplete sample dissolution. Most cesium salts dissolve readily in water and acid solutions. In water
samples, the cesium might require concentration, preferably by ion exchange, or by precipitation
or coprecipitation if interfering ions are present. Organic materials are either decomposed by
HNO3 or dry ashed, and the cesium is extracted with hot water or hot acid solution. Extraction
and leaching procedure have been use to assess exchangeable or leachable cesium using
ammonium acetate solutions or acid solutions, but soils are generally completely solubilized in
HNO3, HCl, HF, H2SO4, or a mixture of these acids in order to account for all the cesium in a soil
sample.
Separation Methods
PRECIPITATION AND COPRECIPITATION. Cesium is separated and purified by several precipitation
and coprecipitation methods using salts of large anions. Gravimetric procedures rely on precipitation to collect cesium for weighing, and several radiochemical techniques isolate cesium radionuclides for counting by precipitation or coprecipitation. Cesium can be precipitated, or
coprecipitated in the presence of cesium carrier, by the chlorate, cobaltinitrate, platinate, and
tetraphenylborate ions. Other alkali metals interfere and should be removed before a pure
insoluble compound can be collected. Cesium can be isolated from other alkali metals by
precipitation as the silicotungstate. The precipitate can be dissolved in 6 M sodium hydroxide,
and cesium can be further processed by other separation procedures. The tetraphenylborate
procedure first removes other interfering ions by a carbonate and hydroxide precipitation in the
presence of iron, barium, lanthanum, and zirconium carriers. Cesium is subsequently precipitated
by the addition of sodium tetraphenylborate to the acidified supernatant. Alum also precipitates
cesium from water samples in the presence of macro quantities of the alkali metals. Trace
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quantities of cesium radionuclides are precipitated using stable cesium as a carrier.
ION EXCHANGE. The cesium cation is not retained by anion-exchange resins and does not form a
suitable anion for anion-exchange chromatography. The process is used, however, to separate
cesium from interfering ions that form anionic complexes. Cesium elutes first in these
procedures. Cesium is retained by cation-exchange resins. Because the cesium ion has the largest
ionic radius and has a +1 charge, it is less hydrated than most other cations. Therefore, cesium
has a small hydrated radius and can approach the cation exchange site to form a strong electrostatic association with the ion-exchange resin. Binding of alkali metal ion to cation exchange
resins follows the order: Cs+1 > Rb+1 > K+1 > Na+1 > Li+1. Cesium is generally the last alkali metal
ion to elute in cation-exchange procedures. In some procedures, the process is not quantitative
after extensive elution.
SOLVENT EXTRACTION. Cesium does not form many complex ions, and solvent extraction is not
a common procedure for its separation. One solvent-extraction procedure, however, is based on
chelation of cesium with TTA in a solvent of methyl nitrate/hydrocarbons. Cesium can also be
extracted from fission product solutions with sodium tetraphenylborate in amyl acetate. It can be
stripped from the organic phase by 3 M HCl.
Methods of Analysis
Macroscopic quantities of cesium have been determined by gravimetric procedures using one of
the precipitating agents described above. Spectrochemical procedures for macroscopic quantities
include flame photometry, emission spectroscopy, and X-ray emission.
Gamma ray spectrometry allows detection of 134Cs, 136Cs, and 137Cs down to very low levels. The
gamma ray measured for 137Cs (661 keV) actually is emitted from it progeny 137mBa. However,
because the half-life of the barium isotope is so short (2.5 min) it is quickly equilibrated with its
parent cesium isotope (i.e., secular equilibrium). Cesium-137 is used as part of a group of
nuclides in a mixed radioactivity source for calibration of gamma ray spectrometers. It is also
used in some liquid scintillation spectrophotometers to generate a Compton distribution to
determine the quench.
Compiled from: Choppin et al., 1995; Considine and Considine, 1983; Cotton and
Wilkinson, 1988; Emsley, 1989; EPA, 1973; EPA, 1973; EPA, 1980; Finston and Kinsley,
1961; Friedlander et al., 1981; Hampel, 1968; Hassinsky and Adolff, 1965; Kallmann, 1964;
Lindsay, 1988; Sittig, 1994.
14.10.9.4 Cobalt
Cobalt, atomic number 27, is a silvery-grey, brittle metal found in the first row of the transition
elements in the periodic table, between iron and nickel. Although it is in the same family of
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elements as rhodium and iridium, it resembles iron and nickel in its free and combined states.
Isotopes
Cobalt-59 is the only naturally occurring isotope of the element. The other twenty-two isotopes
and their metastable states, ranging from mass numbers 50 to 67, are radioactive. Isotopes with
mass numbers less than 59 decay by positron emission or electron capture. Isotopes with mass
numbers greater than 59 decay by beta and gamma emission. Except for 60Co, the most important
radionuclide, their half-lives range from milliseconds to days. The principal isotopes of cobalt
(with their half-lives) are 57Co (t½ . 272 d), 58Co (t½ . 71 d), and 60Co (t½ . 5.27 y). Isotopes 57
and 58 can be determined by X-ray as well as gamma spectrometry. Isotope 60 is easily
determined by gamma spectrometry.
Occurrence and Uses
The cobalt content of the crust of the Earth is about 30 ppm, but the element is widely distributed
in nature, found in soils, water, plants and animals, meteorites, stars, and lunar rocks. Over 200
cobalt minerals are known. Commercially, the most important are the arsenides, oxides, and
sulfides. Important commercial sources also include ores of iron, nickel, copper, silver, manganese, and zinc. Cobalt-60 is produced by neutron activation of stable 59Co. Cobalt-56 and 57Co are
prepared by bombardment of iron or nickel with protons or deuterons. Cobalt-58 (formed by
activation of nickel) is now the dominant isotope formed in nuclear power plants during a fuel
cycle, because most power plants have replaced their cobalt-bearing alloys, such as stellite.
Some of the metallic cobalt is isolated from its minerals, but much of the metal is produced
primarily as a byproduct of copper, nickel, or lead extraction. The processes are varied and
complicated because of the similar chemical nature of cobalt and the associated metals.
Since ancient times, cobalt ores has been used to produce the blue color in pottery, glass, and
ceramics. Cobalt compounds are similarly used as artist pigments, inks, cotton dyes, and to speed
the drying of paints and inks. They also serves as catalysts in the chemical industry and for
oxidation of carbon monoxide in catalytic converters. One of the major uses of cobalt is the
preparation of high-temperature or magnetic alloys. Jet engines and gas turbines are
manufactured from metals with a high content of cobalt (up to 65 percent) alloyed with nickel,
chromium, molybdenum, tungsten, and other metals.
Little use if made of pure cobalt except as a source of radioactivity from 60Co. The radionuclide
is used in cancer radiotherapy, as a high-energy gamma source for the radiography of metallic
objects and other solids, as a food irradiation source for sterilization, or as an injectable radionuclide for the measurement of flow rates in pipes. The half-life of 60Co (t½ . 5.2 y), and its
gamma emissions make it a principal contributor to potential dose effects in storage and transport
of radioactive waste.
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Solubility of Compounds
Most simple cobalt compounds contain Co+2, but Co+2 and Co+3 display varied solubilities in
water. To some extent, their solubilities depend on the oxidation state of the metal. For example,
all the halides of Co+2 are soluble but the only stable halide of Co+3, the fluoride, is insoluble. The
sulfates of both oxidation states are soluble in water. The acetate of Co+2 is soluble, but that of
Co+3 hydrolyses in water. The bromate, chlorate, and perchlorate of Co+2 are also soluble.
Insoluble compounds include all the oxides of both oxidation states, Co+2 sulfide, cyanide,
oxalate, chromate, and carbonate. The hydroxides are slightly soluble. Several thousand complex
compounds of cobalt are known. Almost all are Co+3 complexes and many are soluble in water.
Review of Properties
Metallic cobalt is less reactive than iron and is unreactive with water or oxygen in air unless
heated, although the finely divided metal is pyrophoric in air. On heating in air it forms the
oxides, Co+2 oxide (CoO) below 200 EC and above 900 EC and Co+2-Co+3 oxide (Co3O4) between
the temperature extremes. It reacts with common mineral acids and slowly with hydrofluoric and
phosphoric acids to form Co+2 salts and with sodium and ammonium hydroxides. On heating, it
reacts with halogens and other nonmetals such as boron, carbon, phosphorus, arsenic, antimony,
and sulfur.
Cobalt exists in all oxidation states from !1 to +4. The most common are the +2 and +3
oxidation states. The +1 state is found in a several complex compounds, primarily the nitrosyl
and carbonyl complexes and certain organic complexes. The +4 state exist in some fluoride
complexes. Co+2 is more stable in simple compounds and is not easily hydrolyzed. Few simple
compounds are known for the +3 state, but cobalt is unique in the numerous stable complex
compounds it forms.
The toxicity of cobalt is not comparable to metals such as mercury, cadmium, or lead. Inhalation
of fine metallic dust can cause irritation of the respiratory system, and cobalt salts can cause
benign dermatosis. Cobalt-60 is made available in various forms, in sealed aluminum or monel
cylinders for industrial applications, as wires or needles for medical treatment, and in various
solid and solution forms for industry and research. Extreme care is required in handling any of
these forms of cobalt because of the high-energy gamma radiation from the source.
Solution Chemistry
In aqueous solution and in the absence of complexing agents, Co+2 is the only stable oxidation
state, existing in water as the pink-red hexaaquo complex ion, Co(H2O)6+2. Simple cobalt ions in
the +3 oxidation state decompose water in an oxidization-reduction process that generates Co+2:
4 Co+3 + 2 H2O 6 4 Co+2 + O2 + 4 H+1
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Complexation of Co+3 decreases its oxidizing power and most complex ions of the +3 oxidation
state are stable in solution.
COMPLEXATION. Several thousand complexes of cobalt have been prepared and extensively
studied, including neutral structures and those containing complex cations or anions. The +2
oxidation state forms complexes with a coordination of four or six, and in aqueous solution,
[Co(H2O)6]+2 is in equilibrium with some [Co(H2O)4]+2. In alkaline solution Co+2 precipitates as
Co(OH)2, but the ion is amphoteric; and in concentrated hydroxide solutions, the precipitate
dissolves forming [Co(OH)4]!2. Many complexes of the form [Co(X)4]!1 exist with monodentate
anionic ligands such as Cl!1, Br!1, I!1, SCN!1, N3!1, and OH!1. Many aquo-halo complexes are
known; they are various shades of red and blue. The aquo complex, [Co(H2O)6]+2, is pink.
Chelate complexes are well-known and are used to extract cobalt from solutions of other ions.
Acetylacetone (acac) is used, for example, in a procedure to separate cobalt from nickel. Co+2 and
Ni+2 do not form chelates with the acac, Co+3 does, however, and can be easily extracted.
OXIDATION-REDUCTION BEHAVIOR. Most simple cobalt +3 compounds are unstable because the
+3 state is a strong oxidizing agent. It is very unstable in aqueous media, rapidly reducing to the
+2 state at room temperature. The aqueous ion of Co+2, [Co(H2O)6]+2, can be oxidized, however,
to the +3 state either by electrolysis or by ozone (O3) in cold perchloric acid (HClO4); solutions at
0 EC have a half-life of about one week. Compounds of the Co+3 complex ions are formed by
oxidizing the +2 ion in solution with oxygen or hydrogen peroxide (H2O2) in the presence of
ligands. The Co+3 hexamine complex forms according to:
4 CoX2 + 4 NH4X + 20 NH3 + O2 º 4 [Co(NH3)6]X3 + 2 H2O
HYDROLYSIS. The hydrolysis of the +2 oxidation state of cobalt is not significant in aqueous
media below pH 7. At pH 7, hydrolysis of 0.001 M solution of the cation begins and is
significant at a pH above 9. The hydrolysis of the +3 oxidation state is reminiscent of the
hydrolysis of Fe+3, but it is not as extensive. Hydrolysis of Co+3 is significant at pH 5. In contrast,
the hydrolysis of Fe+3 becomes significant at a pH of about 3.
Dissolution of Samples
Cobalt minerals, ores, metals, and alloys can be dissolved by treatment first with hydrochloric
acid, followed by nitric acid. The insoluble residue remaining after application of this process is
fused with potassium pyrosulfate and sodium carbonate. In extreme cases, sodium peroxide
fusion is used. Biological samples are dissolved by wet ashing, digesting with heating in a
sulfuric-perchloric-nitric acid mixture.

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Separation Methods
PRECIPITATION AND COPRECIPITATION. Cobalt can be precipitated by hydrogen sulfide (H2S),
ammonium sulfide (NH4S), basic acetate (C2H3O2!1/HO!1), barium carbonate (BaCO3), zinc
oxide (ZnO), potassium hydroxide and bromine (KOH/Br2), ether and hydrochloric acid
[(C2H5)2O and HCl], and cupferron. Cobalt sulfide (CoS) is coprecipitated with stannic sulfide
(SnS2) when low-solubility sulfides are precipitated in mineral acids. Care should be taken to
avoid coprecipitation of zinc sulfide (ZnS).
Cobalt can be separated from other metals by hydroxide precipitation using pH control to
selectively precipitate metals such as chromium, zinc, uranium, aluminum, tin, iron (+3),
zirconium, and titanium at low pH. Cobalt precipitates at pH 6.8, and magnesium, mercury,
manganese, and silver at a pH greater than 7. Cobalt is not be separated from metals such as iron,
aluminum, titanium, zirconium, thorium, copper, and nickel using ammonium hydroxide
(NH4OH) solutions (aqueous ammonia), because an appreciable amount of cobalt is retained by
the hydroxide precipitates of these metals produced using this precipitating agent. Various
precipitating agents can be used to remove interfering ions prior to precipitating cobalt: iron by
precipitating with sodium phosphate (Na3PO4) or iron, aluminum, titanium, and zirconium with
zinc oxide.
The separation of cobalt from interfering ions can be achieved by the quantitative precipitation of
cobalt with excess potassium nitrite (KNO2) to produce K3[Co(NO2)6] (caution: heating
K3[Co(NO2)6] after standing for some time makes it unstable). Ignition can be used to collect the
cobalt as its mixed oxide (Co3O4). Cobalt can also be precipitated with α-nitroso-β-napthol (1nitroso-2-napthol) to separate it from interfering metals. Nickel can interfere with this precipitation, but can be removed with dimethylglyoxime. Precipitation of Co+2 as mercury tetracyanatocobaltate (+2) {Hg[Co(SCN)4]} also is used, particularly for gravimetric analysis, and
precipitation with pyridine in thiocyanate solution is a quick gravimetric product,
[Co(C5H5N)4](SCN)2.
SOLVENT EXTRACTION. Various ions or chelates have been used in solvent extraction systems to
isolate cobalt from other metals. Separation has been achieved by extracting either cobalt itself
or, conversely, extracting contaminating ions into an organic solvent in the presence of hydrofluoric acid (HF), hydrochloric acid, and calcium chloride (HCl/CaCl2), hydrobromic acid (HBr),
hydroiodic acid (HI), or ammonium thiocyanate (NH4SCN). For example, Co+2 has been
separated from Ni+2 by extracting a hydrochloric acid solution containing calcium chloride with
2-octanol. The ion is not extracted by diethyl ether from hydrobromic acid solutions, but it is
extracted from ammonium thiocyanate solutions by oxygen-containing organic solvents in the
presence of Fe+3 by first masking the iron with citrate.
Several chelate compounds have been used to extract cobalt from aqueous solutions. Acetylacetone (acac) forms a chelate with Co+3, but not Co+2, that is soluble in chloroform at pH 6 to 9,
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permitting separation from several metals including nickel. Co+2 can be oxidized to Co+3 with
hydrogen peroxide (H2O2) prior to extraction. The chelating agent α-nitroso-β-napthol has also
been used in the separation of Co+3 by solvent extraction. Diphenylthiocarbazone (dithizone) has
been used at pH 8 to extract cobalt into carbon tetrachloride and chloroform after metals that
form dithizonates in acid solution (pH 3-4) have been removed. 8-quinolinol has been used in a
similar manner at pH up to 10. Masking agents added to the system impede the extraction of iron,
copper, and nickel.
ION-EXCHANGE CHROMATOGRAPHY. Anion-exchange resins have been used extensively to
separate cobalt from other metals. The chloro-metal complexes, prepared and added to columns
in molar hydrochloric acid solutions, are eluted at varying concentrations of hydrochloric acid.
Trace amounts of 59Fe, 60Co, and 65Zn and their respective carriers have been separated from
neutron-irradiated biological tissue ash with a chloride system. Cobalt-60 has been eluted carrierfree from similar samples and columns prepared with hydrobromic acid. Cobalt and contaminated metals in nitric-acid systems behave in a manner similar to hydrochloric-acid systems. Co+2cyanide and cyanate complexes have been used to separate cobalt from nickel. The basic form of
quaternary amine resins (the neutral amine form) has been used in the column chromatography of
cobalt. Both chloride- and nitrate-ion systems have resulted in the association of cobalt as a
complex containing chloride or nitrate ligands as well as the neutral (basic) nitrogen atom of the
amine resin. Resins incorporating chelates in their matrix system have been used to isolate
cobalt. 8-quinolinol resins are very effective in separating cobalt from copper.
ADSORPTION CHROMATOGRAPHY. Several inorganic adsorbents such as alumina, clays, and silica
are used to separate cobalt. Complex ions of cobaltamines separate on alumina as well as Co +2
complexes of tartaric acid and dioxane. A complex of nitroso-R-salts are adsorbed onto an
alumina column while other metals pass through the column. Cobalt is eluted with sulfuric acid.
Cobalt dithizonates adsorb on alumina from carbon tetrachloride solutions. Cobalt is eluted with
acetone. The separation of cobalt from iron and copper has been achieved on aluminum
hydroxide [Al(OH)3]. Clay materials—kaolinite, bentonite, and montmorillonite—separate Co+2
from Cu+2. Cu+2 adsorbs and Co+2 elutes with water. Silica gel and activated silica have both been
used as adsorbents in cobalt chromatography.
Organic adsorbents such as 8-hydroxyquinoline and dimethylglyoxime have been used in cobaltadsorption chromatographic systems. Powdered 8-hydroxyquinoline separates Co+2 from other
cations and anions, for example, and dimethylglyoxime separates cobalt from nickel. Cobaltcyano complexes adsorb on activated charcoal, and cobalt is eluted from the column while the
anionic complexes of metals such as iron, mercury, copper, and cadmium remain on the column.
Numerous paper chromatograph systems employing inorganic or chelating ligands in water or
organic solvents are available to separate cobalt from other metals. In one system, carrier-free
60
Co and 59Fe from an irradiated manganese target were separated with an acetone-hydrochloric
solvent.
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ELECTRODEPOSITION. Most electroanalytical methods for cobalt are preceded by isolating the
cobalt from interfering ions by precipitation or ion exchange. The electrolyte is usually an
ammonia solution that produces the hexamine complex of Co+2, Co(NH3)6+2 in solution.
Reducing agents such as hydrazine sulfate are added to prevent anodic deposits of cobalt and the
oxidation of the Co+2-amine ion. Cobalt and nickel can be separated electrolytically by using an
aqueous solution of pyridine with hydrazine to depolarize the platinum anode. The nickel is
deposited first, and the voltage is increased to deposit cobalt.
Methods of Analysis
Cobalt-57, 58Co, and 60Co maybe concentrated from solution by coprecipitation and determined
by gamma-ray spectrometry. Cobalt-60 is most commonly produced by the neutron activation of
59
Co, in a reactor or an accelerator. Cobalt-58 is most commonly produced from the following
reaction in nuclear reactors, 58Ni(n,p)58Co, due to the presence of nickel bearing alloys which
undergo corrosion and are transported through the reactor core. Cobalt-58 is the most significant
contributor to the gamma ray induced radiation fields in these facilities. Cobalt-57 can be
produced by either of the following, 58Ni(n,d)57Co [reactor] or 56Fe(d,n)57Co [accelerator], Cobalt57 and 60Co are frequently used as part of a mixed radionuclide source for calibration of gamma
ray spectrometers.
Compiled from: Baes and Mesmer, 1976; Bate and Leddicotte, 1961; Cotton and Wilkinson,
1988; Dale and Banks, 1962; EPA, 1973; Greenwood and Earnshaw, 1984; Haissinsky and
Adloff, 1965; Hillebrand et al., 1980; Larsen, 1965; Latimer, 1952; Lingane, 1966.
14.10.9.5 Iodine
Iodine is a nonmetal, the last naturally occurring member of the halogen series, with an atomic
number of 53. In the elemental form it is a diatomic molecule, I 2, but it commonly exists in one
of four nonzero oxidation states: !1 with metal ions or hydrogen; and +1, (V), and (VII) with
other nonmetals, often oxygen. Numerous inorganic and organic compounds of iodine exist,
exhibiting the multiple oxidation states and wide range of physical and chemical properties of the
element and its compounds. Existence of multiple oxidation states and the relative ease of
changing between the !1, 0, and (V) state allows readily available methods for separation and
purification of radionuclides of iodine in radiochemical procedures.
Isotopes
There are 42 known isotopes of iodine, including seven metastable states. The mass numbers
range from 108 to 142. The only stable isotope is naturally occurring 127I. The half-lives of the
radionuclides range from milliseconds to days with the single exception of long-lived 129I (t1/2 .
1.57×107 y). Iodine radionuclides with lower mass numbers decay primarily by electron capture.
The high mass numbers are, for the most part, beta emitters. The significant radionuclides are 123I
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(t1/2 . 13.2 h), 125I (t1/2 . 60.1 d, electron capture), 129I (β), and 131I (t1/2 .8 d, β).
Occurrence and Uses
Iodine is widely distributed, but never found in the elemental form. The average concentration in
the Earth’s crust is about 0.3 ppm. In seawater, iodine concentration, in the form of sodium or
potassium iodide, is low (about 50 ppb), but it is concentrated in certain seaweed, especially kelp.
It is also found in brackish waters from oil and salt wells. The sources are saltpeter and nitratebearing earth in the form of calcium iodate, well brine, and seaweed. Iodine is produced from
calcium iodate by extraction of the iodate from the source with water and reduction of the iodate
with sodium bisulfite to iodine. Iodine is precipitated by mixing with the original iodate liquor to
cause precipitation. Iodine can also be obtained from well brine, where the iodide ion is oxidized
with chlorine, and then the volatile iodine is blown out with a stream of air. Sodium or potassium
iodide in seaweed is calcined to an ash with sulfuric acid, which oxidizes the iodide to iodine.
Iodine from any of these processes can be purified by sublimation.
Isotopes of iodine of mass $ 128 may all be formed as a result of fission of uranium and
plutonium. Nuclear reactors and bomb tests are the most significant sources of these radioisotopes with the exception of 131I. That isotope is routinely produced for use in medical imaging
and diagnosis. The isotopes released from the other sources represent a short-term environmental
health hazard should there be an abnormal release from a reactor or testing of bombs.
This was the case in 1979 and 1986 when the reactor incidents at Three Mile Island and
Chernobyl caused releases of radioiodines. During the former event, a ban on milk distribution in
the downwind corridor was enforced as a purely preventative measure. In the latter case, significant releases of iodines and other isotopes caused more drastic, long term measures for food
quarantine.
Deposits on the surface of plants could provide a quick source of exposure if consumed directly
from fruits and vegetables or indirectly from cow’s milk. It would readily accumulate in the
thyroid gland, causing a short-term exposure of concern. It represent the greatest short-term
exposure after a nuclear detonation and has been released in power plant accidents. Iodine-129,
with of a half-life of more than 15 million years, represent a long-term environmental hazard. In
addition to its long half-life, the environmental forms of iodine in the environment are highly
soluble in groundwater and are poorly sorbed by soil components. It is not absorbed at all by
granite, and studies at a salt repository indicate that 129I would be only one of few radionuclides
that would reach the surface before it decayed. Therefore, research on the fate of 129I that might
be released suggests that the radionuclide would be highly disseminated in the ecosystem.
Iodine-131 is analyzed routinely in milk, soil and water. Iodine-129 is a low energy beta and
gamma emitter, which has a very long half-life (t½ . 1.47×107 y). The most significant concern
for this isotope is in radioactive waste, and its potential for migration due to the chemistry of
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iodine in the environment. Iodine-131 is produced for medical purposes by neutron reaction as
follows: 130Te(n,γ)131Te 6 beta decay 6 131I (t½ . 8 d).
The major use of iodine, iodine radionuclides, and iodine compounds is in medical diagnosis and
treatment. Iodine-123, 125I, and 131I are use for diagnostic imaging of the thyroid gland and the
kidneys. Iodine-131 is used to treat hyperthyroidism and thyroid cancer. Stable iodine in the form
of potassium iodide is added to commercial salt to prevent enlargement of the thyroid (goiter).
Iodine in the form of the hormone thyroxine is also used for thyroid and cardiac treatment and
hormone replacement therapy in iodine deficiency. Iodine radionuclides are used as a tracer in
the laboratory and industry to study chemistry mechanisms and processes and to study biological
activity and processes. Iodine is a bactericide and is used as an antiseptic and sterilization of
drinking water. It is used as a catalyst in chemical processes and as silver iodide in film
emulsions.
Solubility of Compounds
Molecular iodine is only very slightly soluble in water (0.33 g/L), but it is soluble in solutions of
iodide ion, forming I3!1. It is appreciably soluble in organic solvents. Carbon tetrachloride (CCl4)
or chloroform (CHCl3) are commonly used to extract iodine from aqueous solutions after
alternate forms of the element, typically I!1 and IO3!1, are converted to I 2. The solutions have a
violet color in organic solvents, and iodine dimerizes to some extent in these solutions:
2 I2 º I4
Numerous compounds of iodine are soluble in water. All metallic iodides are soluble in water
except those of silver, mercury, lead, cupurous ion, thallium, and palladium. Antimony, bismuth,
and tin iodides require a small amount of acid to keep them in solution. Most of the iodates and
periodates are insoluble. The iodates of sodium, potassium, rubidium, and the ammonium ion are
soluble in water. Those of cesium, cobaltous ion, magnesium, strontium, and barium are slightly
soluble in water but soluble in hot water. Most other metallic iodates are insoluble.
Review of Properties
Elemental iodine (I2) is a purple-black, lustrous solid at room temperature with a density of 4.9
g/cm3. The brittle crystals have a slightly metallic appearance. Iodine readily sublimes and stored
in a closed clear, colorless container, it produces a violet vapor with an irritating odor. Iodine has
a melting point of 114 EC and a boiling point of 184 EC.
The chemical reactivity of iodine is similar to the other halogens, but it is the least electronegative member of the family of elements and the least reactive. It readily reduces to iodide, and
is displaced from its iodides by the other halogens and many oxidizing agents. Iodine combines
directly with most elements to form a large number of ionic and covalent compounds. The
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exceptions are the noble gases, carbon, nitrogen, and some noble metals.
The inorganic compounds of iodine can be classified into three groups: (1) iodides, (2)
interhalogen, and (3) oxides. Iodine forms iodides that range from ionic compounds such as
potassium iodide (KI) to covalent compounds such as titanium tetraiodide (TiI 4) and phosphorus
triiodide (PI3), depending on the identity of the combining element. More electropositive (less
electronegative) metals (on the left side of the Periodic Table, such as alkali metals and alkaline
earths) form ionic compounds. Less electropositive metals and more electronegative nonmetals
tend to form covalent compounds. Interhalogen compounds include the binary halides, such as
iodine chloride (ICl), iodine trichloride (ICl3), and iodine pentafluoride (IF5), or contain
interhalogen cations and anions, such as ICl2+1, IF6+1, I+3, ClIBr!1, ICl4!1, and I6!2. Oxygen
compounds constitute the oxides, I2O5 and I4O9 (containing one I+3 cation and three IO3!1 anions),
for example; the oxyacids, such as hypoiodous acid (HIO) and iodic acid (HIO3); and compounds
containing oxyanions, iodates (IO3!1) and periodates (IO4!1) are the common ones.
Organoiodides include two categories: (1) iodides and (2) iodide derivatives with iodine in a
positive oxidation state because iodine is covalently bonded to another, more electronegative
element. Organoiodides contain a carbon iodide bond. They are relatively dense and volatile and
more reactive than the other organohalides. They include the iodoalkanes such as ethyl iodide
(C2H5I) and iodobenzene (C6H5I). Dimethyliodonium (+3) hexafluoroantimonate
[(CH3)2I+3SbF6-3], a powerful methylating agent, is an example of the second category.
The radionuclides of iodine are radiotoxic, primarily because of their concentration in the thyroid
gland. Toxicity of 129I, if released, is a concern because of its extremely long half-life. Iodine-131,
with a half-life of eight days, is a short-term concern. The whole-body effective biological halflives of 129I and 131I are 140 d and 7.6 d, respectively.
Solution Chemistry
OXIDATION-REDUCTION BEHAVIOR. Iodine can exist in multiple oxidation states in solution, but
the radiochemist can control the states by selection of appropriate oxidizing and reducing agents.
In acid and alkaline solutions, the common forms of iodine are: I!1, I2, and IO3!1. Hypoiodous
acid (HIO) and the hypoiodite ion (IO!1) can form in solution, but they rapidly disproportionate:
5 HIO º 2 I2 + IO3!1 + H+1 + 2 H2O
3 IO!1 º 2 I!1 + IO3!1
Iodine itself is not a powerful oxidizing agent, less than that of the other halogens (F2, Cl2, and
Br2), but its action is generally rapid. Several oxidizing and reducing agents are used to convert
iodine into desired oxidation states during radiochemical procedures. These agents are used to
promote radiochemical equilibrium between the analyte and the carrier or tracer or to produce a
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specific oxidation state before separation: I2 before extraction in an organic solvent or I!1 before
precipitation, as examples. Table 14.18 presents oxidizing and reducing agents commonly used
in radiochemical procedures:
Table 14.18 — Common radiochemical oxidizing
and reducing agents for iodine
Redox Process

Redox Reagent

I!1 6 I2

HNO2

!1

!1

(NaNO2 in acid)

I 6 IO3

MnO2 in acid

I2 6 I-

6 M HNO3
NaHSO3 and NaHSO4 (in acid)
Na2SO3 and Na2S2O3
Fe2(SO4)3 (in acid)
SO2 gas
NaHSO3 and (NH4)2SO3

I!1 6 IO4!1

KMnO4
50% CrO3 in 18N H2SO4

I!1 6 IO4!1

NaClO in base

!1

NH2OH·HCl

!1

IO3 6 I2

NH2OH·HCl
H2C2O4 in 18N H2SO4

IO4!1 6 I!1

NaHSO3 in acid

IO4 6 I2

Radiochemical exchange between I2 and I!1 in solution is complete within time of mixing and
before separation. In contrast, exchange between I2 and IO3!1 or IO4!1 in acid solution and
between IO3!1 and IO4!1 in acid or alkaline solution is slow. For radiochemical analysis of iodine,
experimental evidence indicates that the complete and rapid exchange of radioiodine with carrier
iodine can be accomplished by the addition of the latter as I!1 and subsequent oxidation to IO4!1
by NaClO in alkaline solution, addition of IO4!1 and reduction to I!1 with NaHSO3, or addition of
one followed by redox reactions first to one oxidation state and then back to the original state.
COMPLEXATION. As a nonmetal, iodine is generally not the central atom of a complex, but it can
act as a ligand to form complexes such as SiI6!2 and CoI6!3. An important characteristic of
molecular iodine is its ability to combine with the iodide ion to form polyiodide anions. The
brown triioide is the most stable:
I2 + I!1 º I3!1
The equilibrium constant for the reaction in aqueous solution at 25 EC is 725, so appreciable
concentrations of the anion can exist in solution, and the reaction is responsible for the solubility
of iodine in iodide solutions.
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HYDROLYSIS. Iodine hydrolyzes in water through a disproportionation reaction:
I2 + H2O º H+1 + I!1 + HIO
Because of the low solubility of iodine in water and the small equilibrium constant (k=2.0×10-13),
hydrolysis produces negligible amounts of the products (6.4×10!6 M) even when the solution is
saturated with iodine. Disproportionation of HIO produces a corresponding minute quantity of
IO3!1 (see the reaction above). In contrast, in alkaline solution, I2 produces I!1 and IO!1:
I2 + 2 OH!1 º I!1 + IO!1 + H2O
The equilibrium constant favors the products (K = 30), but the actual composition of the solution
is complicated by the disproportionation of IO!1 (illustrated above), giving I!1 and IO3!1. The
equilibrium constant for the reaction of IO!1 with hydroxide ion is very large (1020), and the rate
of the reaction is very fast at all temperatures. Therefore, the actual products obtained by
dissolving iodine in an alkaline solution are indeed I!1 and IO3!1, quantitatively, and IO!1 does not
exist in the solution.
Dissolution of Samples
Iodine compounds in rocks are often in the form of iodides that are soluble in either water or
dilute nitric acid when the finely divided ores are treated with one of these agents. Those that are
insoluble under these conditions are solubilized with alkali fusion with sodium carbonate or
potassium hydroxide, followed by extraction of the residue with water. Insoluble periodiates can
be decomposed by cautious ignition, converting them to soluble iodides.
Metals containing iodine compounds are dissolved in varying concentrations of nitric, sulfuric, or
hydrochloric acids. Dissolution can often be accomplished at room temperature or might require
moderation in an ice bath.
Organoiodides are decomposed with a sodium peroxide, calcium oxide, or potassium hydroxide
by burning in oxygen in a sealed bomb. Wet oxidation with mixtures of sulfuric and chromic
acids or with aqueous hydroxide is also used.
Separation Methods
PRECIPITATION. The availability of stable iodine as a carrier and the relative ease of producing
the iodide ion make precipitation a simple method of concentrating and recovering iodine
radionuclides. The two common precipitating agents are silver (Ag+1) and palladium (Pd+2)
cations, which form silver iodide (AgI) and palladium iodide (PdI2), respectively. Silver iodide
can be solubilized with a 30 percent solution of potassium iodide. Palladium precipitates iodide
in the presence of chloride and bromide, allowing the separation of iodide from these halides.
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The precipitating agent should be free of Pd+4, which will precipitate chloride. If Pd+2 iodide is
dried, precaution should be taken as the solid slowly looses iodine if heated at 100 EC. Iodate can
be precipitated as silver iodate, and periodate as lead periodate.
SOLVENT EXTRACTION. One solvent extraction method is commonly used to isolate iodine. After
preliminary oxidation-reduction steps to insure equilibrium of all iodine in solution, molecular
iodine (I2) is extracted from aqueous solutions by a nonpolar solvent, usually carbon tetrachloride
or chloroform. It is not uncommon to add trace quantities of the oxidizing or reducing agent to
the extraction solution to ensure and maintain all iodine in the molecular form. Hydroxylamine is
added, for example, if iodate is the immediate precursor of iodine before extraction.
ION-EXCHANGE CHROMATOGRAPHY. Both cation and anion exchange procedures are used to
separate iodine from contaminants. Cation-exchange chromatography has been used to remove
interfering cations. To remove 137Cs activity, an iodine sample in the iodide form is exchanged on
a cation resin and eluted with ammonium sulfite [(NH4)2SO3] to ensure maintenance of the iodide
form. Cesium cations remain on the resin. Bulk resin also is used, and iodide is washed free of
the resin also with sodium hypochlorite (NaClO) as the oxidizing agent. Anion resins provide for
the exchange of the iodide ion. The halides have been separated from each other on an anionexchange column prepared in the nitrate form by eluting with 1 M sodium nitrate. Iodide can also
be separated from contaminants by addition to an anion exchanger and elution as periodate with
sodium hypochlorite. The larger periodate anion is not as strongly attracted to the resin as the
iodide ion. Iodine-131 separation, collection, and analysis is performed by absorbing the
radionuclide on an anion-exchange resin and gamma counting it on the sealed column after
eluting the contaminants.
DISTILLATION. Molecular iodine is a relatively volatile substance. Compared to many
contaminating substances, particularly metal ions in solution, its boiling point of 184 EC is very
low, and the volatility of iodine provides a method for its separation from other substances. After
appropriate oxidation-reductions steps to convert all forms of iodine into the molecular form,
iodine is distilled from aqueous solution into sodium hydroxide and collected by another
separation process, typically solvent extraction. In hydroxide solution, molecular iodine is
converted to a mixture of iodide and hypoiodite ions and then into iodide and periodate ions, and
suitable treatment is required to convert all forms into a single species for additional procedures.
Methods of Analysis
Macroquantities of iodine can be determined gravimetrically by precipitation as silver iodide,
palladium iodide, or cuprous iodide. The last two substances are often used to determine the
chemical yield in radiochemical analyses. Microquantities of 129I and 131I are coprecipitated with
palladium iodide or cuprous iodide using stable iodide as a carrier and counted for quantification.
Iodine-129 usually is beta-counted in a liquid-scintillation system, but it also can be determined
by gamma-ray spectrometry. Iodine-129 can undergo neutron activation and then be measured by
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gamma-ray spectrometry from the 130I (t½ . 12.4 h) produced by the neutron-capture reaction.
The method uses conventional iodine valence adjustments and solvent extraction to isolate the
Iodine fraction. Chemically separated 129I is isolated on an anion exchange resin before being
loaded for irradiation. A lower limit of detection (0.03 ng) can be achieved with a neutron flux of
5×1014 n/cm2·s for 100 seconds. Iodine-129 also can be determined directly by mass spectrometry. The measurement limit by this technique is approximately 2 femtograms. Special counting
techniques, such as beta-gamma coincidence, have also been applied to the analysis of 129I.
Iodine-131 is determined by gamma-ray emission. Mass spectrometry has been used for
measurement of 125I and 129I.
Compiled from: Adams, 1995; APHA, 1998; Armstrong et al., 1961; Bailar et al., 1984; Bate
and Stokely, 1982; Choppin et al., 1995; Considine and Considine, 1983; Cotton and
Wilkinson, 1988; DOE, 1990 and 1997, 1997; EPA, 1973; EPA, 1980; Ehmann and Vance,
1991; Greenwood and Earnshaw, 1984; Haissinsky and Adloff, 1965; Kleinberg and Cowan,
1960; Latimer, 1952; Lindsay, 1988; McCurdy et al., 1980; Strebin et al., 1988.
14.10.9.6 Neptunium
Neptunium, atomic number 93, is a metal and a member of the actinide series. The relatively
short half-lives of the neptunium isotopes obviate naturally occurring neptunium from being
detected in environmental samples (except in some rare instances). Thus, all detected isotopes
are produced artificially, principally by neutron bombardment of uranium. Neptunium has six
possible oxidation states: +2, +3, +4, (V), (VI), and (VII). The most stable ionic form of
neptunium is the NpO2+1 ion. The ionic states of neptunium are similar to that of manganese,
however the chemistry is most closely associated with uranium and plutonium.
Isotopes
There are 17 isotopes of neptunium, which include three metastable states. The mass range of
neptunium isotopes is from 226 to 242. All isotopes are radioactive, and the longest-lived
isotope, 237Np, has a half-life of 2.1×106 years and decays by alpha emission (principal decay
mode) or spontaneous fission (very low probability of occurrence). The most common mode of
decay for the other neptunium isotopes is by β-particle emission or electron capture.
Neptunium is formed in nuclear reactors from two separate neutron-capture reactions with
uranium. Thus the largest quantity of neptunium isotopes are associated with spent nuclear fuel.
In fuel reprocessing, the focus is on the recovery of uranium and plutonium isotopes. Thus the
neptunium isotopes are part of the waste stream from that process.
The short-lived 239Np can be used as a tracer when separated from its parent 243Am. With the halflife of the americium at 7,370 years, and that of the neptunium is only 2.3 days, tracer quantities
can be successfully removed every 6–10 days from an americium source.
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Occurrences and Uses
Neptunium was the first of the actinides to be produced synthetically (in 1940). Neptunium-239
(t½ . 25 min) resulted from neutron bombardment of natural uranium.
Neptunium-237 is formed as a result of successive neutron capture on a 235U nucleus to form
237
U. This uranium isotope has a reasonably short half-life (6.75 d). After a 235U target has been
irradiated with neutrons, most of the 237U activity will have decayed to 237Np after about 30 days
(no radiochemical equilibrium; see Attachment 14A, “Radioactive Decay and Equilibrium”). At
that time, the 237Np may be “milked” from the source.
Neptunium-237 (t½ . 2.1×106 y), is irradiated with neutrons to form 238Np, which decays to 238Pu.
Plutonium-238 is used in space vehicles as a power source because of its superior energy
characteristics. Neptunium-237 can be used in neutron detection equipment because it has a
significant (n,γ) capture cross-section. The 238Np produced has a half-life of 2.1 days with easily
determinable beta or gamma emissions.
Solubility of Compounds
Neptunium solubility is strongly dependent upon oxidation state. The +3 and +4 states form very
insoluble fluorides, while the (V) and (VI) states are soluble. This property is an effective means
of separation of neptunium from uranium. Neptunium (+4) may be carried on zirconium
phosphate precipitate, indicating its insolubility as a phosphate only in that oxidation state.
Neptunium forms two oxides, NpO2 and Np3O8, both of which are soluble in concentrated
hydrochloric, perchloric and nitric acids. The most soluble of the neptunium compounds are
Np(SO4)2, Np(C2O4)2, Np(NO3)5, Np(IO3)4, and (NH4)2Np2O7. Neptunium (+3) compounds are
easily oxidized to Np+4 when exposed to air.
Review of Properties
Neptunium is a silvery, white metal, which is rapidly oxidized in air to the NpO2 compound.
NpF3 is formed by the action of hydrogen and HF on NpO2. NpF4 is formed by the action of
oxygen and HF on NpF3. These reactions, and similar ones for the other halides take place at
~500 EC. All the halides are volatile above 450 EC, with the hexafluoride boiling at 55 EC. All
the halides undergo hydrolysis in water to form the oxo-complex or ions.
Neptunium is found in the environment at very low concentrations due to the short half-lives of
its isotopes and the few reactions through which 237Np, its long-lived isotope, can be formed. The
principal nuclear reactions are identified here:
238

U(n, 2n)237U 6

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Np + β!

237

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235

U(n,γ) 6

236

U(n,γ) 6

237

U 6

Np + β!

237

Solution Chemistry
Neptunium most closely resembles uranium in its solution chemistry, although it has many
differences that allow it to be easily separated. The +4 and (V) oxidation states are the two most
commonly encountered in chemical and environmental analysis of neptunium.
COMPLEXATION. Neptunium forms complexes with fluorides, oxalates, phosphates, sulfates, and
acetates in the +4 oxidation state at the macro level. However, for chemical separation of
neptunium in concentrations found in environmental samples, the sulfate or the fluoride of the +4
oxidation state can be co-precipitated with BaSO4 or LaF3, respectively.
Neptunium (+4) also forms strong complexes in HCl and HNO 3 with the chloride and nitrate
anions. These complexes appear to have similar complexation constants and charge densities as
those of U(VI) and Pu(VI) in the same media. Neptunium(V) forms weak complexes with
oxalate ions. Complexation in basic media with potassium phosphotungstate or lithium
hydroxide has been shown to be a useful method for oxidation-reduction potential measurements
as the individual oxidation states are stabilized significantly.
OXIDATION-REDUCTION. The most stable oxidation state of neptunium in aqueous solution is
(V). Oxidation in basic solution to (VI) can be achieved with MnO4!, or BrO3!. Like manganese,
neptunium can form the (VII) state. This can be achieved in basic solution with nitrous oxide,
persulfate, or ozone.
Solutions of Np(V) can undergo disproportionation to yield the (VI) and +4 oxidation states. This
reaction has a small equilibrium constant. However, in sulfuric acid media this may be
accelerated a thousand fold, because sulfates complex with the Np+4 ion, driving the disproportionation reaction towards completion.
Dissolution of Samples
The dissolution of samples containing neptunium must be rigorous in ensuring complete
dissolution, because no stable isotopes of neptunium exist to act as carriers. High temperature
furnace oxidation of soil, vegetable, and fecal samples will ensure that the neptunium will be in
the (VI) oxidation state. The resultant ash can be dissolved using lithium metaborate or
perchloric acid. At that point it may be selectively reduced to either the (V) or +4 oxidation state,
depending upon the other analytes from which it must be separated.
Separation Methods
PRECIPITATION AND COPRECIPITATION. The only samples that will have a significant amount of
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neptunium will be high-level wastes (HLW) resulting from spent fuel. Thus, for other sample
analyses, the methods of precipitation of neptunium usually involve the use of a co-precipitant. In
this respect, neptunium acts just like uranium. The +4 oxidation state is the one that will coprecipitate with LaF3. If Np(V) or (VI) are formed, they will not precipitate with fluoride but stay
in solution. This is analogous to the chemistry of the U+4 and U(VI) ions in solution.
Neptunium, like the other actinides, will flocculate with a general precipitating reagent such as
iron hydroxide or titanium hydroxide.
SOLVENT EXTRACTION. Neptunium can be extracted into organic solvents such as methyl
isobutyl ketone (MIBK), TBP, xylene and dibutoxytetraethylene glycol. The +4, (V), and (VI)
oxidation states are extracted using these solvents under a variety of conditions. In all cases, care
must be taken to eliminate or mask any fluorides, oxalates, or sulfates that are present, because
they will have a significant effect on the extraction efficiency. The extraction process is aided by
complex-forming compounds such as TTA, TIOA, trioctylphosphine oxide (TOPO), or
tributylamine (TBA). Several different methods have been developed that use combinations of
these chelates as well. In these instances a synergistic effect has been noted.
ION-EXCHANGE CHROMATOGRAPHY. The four principal neptunium oxidation states are soluble in
dilute to concentrated HCl, HClO4, HNO3, and H2SO4. Although neptunium forms complexes
with these ions in solution the exchange constant for a cation exchange resin is much greater, and
the Np ions are readily removed for the aqueous system. The elution pattern of the oxidation
states is, as with the other transuranics, lowest to highest ionic charge density. Thus the most
strongly retained is the +4:
NpO2+ < NpO22+ < Np3+ < Np4+.
Neptunium can be separated effectively from uranium and plutonium using an anion exchange
method. The plutonium and neptunium are reduced to the +4 state with uranium as (VI) in HCl.
The uranium elutes, while the neptunium and plutonium are retained. The plutonium may then be
reduced to the +3 state using iodide or hydrazine, and will be eluted off the resin in the HCl
solution.
More recently, resin loaded with liquid extractants has been used very successfully to separate
the actinides. Neptunium can be separated selectively from plutonium and uranium using a
TEVA® column, after the neptunium has been reduced to the +4 state using ferrous sulfamate.
This process has been shown to be successful for water, urine, soil, and fecal samples.
Methods of Analysis
Neptunium-237 is the radioisotope most commonly used as a tracer for neptunium recovery. The
principal means of detection of this isotope is alpha spectrometry following a NdF3 or LaF3
coprecipitation step. The 4.78 MeV alpha peak is easily resolved from other alpha emitters
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(notably plutonium) whose chemistry is analogous to that of neptunium. The 239Np radioisotope
could also be used as a tracer. It could be isolated from the parent 243Am source, whose
characteristic gamma-ray of 106 keV is used for quantitation. The other neptunium isotopes are
most easily determined after separation and appropriate sample mounting using gas flow
proportional counting.
Compiled from: Horwitz et al., 1995; Morss and Fuger, 1992; Sill and Bohrer, 2000.
14.10.9.7 Nickel
Isotopes
Twenty-four isotopes of nickel exist from mass number 51 to 74. It has five stable isotopes, and
the most significant of its radioisotopes are 63Ni (t½ . 100 y) and 59Ni (t½ . 7.6 × 104 y). All other
isotopes have half-lives of 5 days or less.
Occurrence
Nickel is found in nature as one of two principal ores, pentlandite or pyrrhotite. It is also a
significant constituent of meteorites. It is a silvery white metal used in the production of Invar,
Hastalloy, Monel, Inconel and stainless steels. Its other principal use is in coins. Corrosion
resistant alloys containing nickel are used in the fabrication of reactor components. During the
life cycle of the reactor, the nickel is converted to the two long-lived radionuclides through the
following reactions: 58Ni(n,γ)59Ni and 62Ni(n,γ)63Ni.
The Code of Federal Regulations (Title 10, Part 61) identifies these isotopes as having specific
limits “in activated metal,”because the material must be physically sampled and dissolved in
order to assess the level of contamination of these isotopes in the metal.
Nickel-63 is a key component in the electron-capture detector of gas chromatographic systems.
This technique is used particularly for organic compounds containing chlorine and phosphorus.
Nickel-63 decays by emission of a low-energy beta (Eβmax = 0.066 MeV), which establishes a
baseline current in the detector system. When a compound containing phosphorus or chlorine
passes the source, these elements can “capture an electron.” The response to this event is an
electrical current less than the baseline current, which is converted into a response used to
quantify the amount of material.
Solubility of Compounds
The soluble salts of nickel are chlorides, fluorides, sulfates, nitrates, perchlorates, and iodides.
Nickel sulfide is very insoluble and will dissolve initially from solutions at low pH. However,
upon exposure to air, such solutions will form the very insoluble compound Ni(OH)S. Nickel
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hydroxide is also insoluble (Ksp = 2 × 10!16) and forms a very gelatinous precipitate, which can
scavenge other radionuclides. Thus, avoiding the formation of this compound is very important.
Solutions of neutral pH, where nickel is suspected of being a component, should be treated with
ammonia to maintain the solubility of this metal ion.
Review of Properties
Nickel metal is highly resistant to air or water oxidation. It exists in the +2 oxidation state under
normal conditions. It can be oxidized to the +3 oxidation state, to NiO(OH), by treatment of Ni+2
with aqueous bromine in potassium hydroxide. It can exist as a +4 ion in compounds such as
NiO2 (used in NiCad batteries), by oxidation with strong oxidants such as peroxydisulfate. In the
+4 oxidation state nickel is a very strong oxidant and will react with water in aqueous solutions.
Nickel metal has been used in the radiochemistry laboratory as an electrode for the galvanic
plating of polonium from hydrochloric acid solutions (see Section 14.10.9.17). In these instances,
the polonium is being removed as interference in the alpha analysis of uranium or plutonium.
Solution Chemistry
Acid solutions of macroscopic quantities of nickel are emerald green. This is due to the
formation of the hexaaquonickel complex, which is very stable.
OXIDATION. Nickel metal will readily dissolve in most mineral acids. The exception is in
concentrated nitric acid, where the metal forms a passive oxide layer resistant to normal
oxidation. Under normal laboratory conditions it will only form the +2 ion.
An usual property of nickel metal is that it forms a volatile carbonyl complex (boiling point
50 EC) when treated with carbon monoxide gas at low temperatures. This carbonyl compound
decomposes to nickel metal at 200 EC. Thus, for samples with a high organic content that may be
placed in a furnace for combustion, a high flow of air or oxygen should be assured if nickel is
going to be analyzed for in the residue.
COMPLEXATION. Nickel forms strong complexes with nitrogen containing compounds such as
ammonia, ethylene diamine, EDTA, and diethylenetriamine. The complex with ammonia forms a
deep blue color distinct from the green color of the normal aqueous ion. The nickel ammonia
complex has a large formation constant and is very stable in the pH range 7–10. This particular
property of nickel is used to separate it from other metals and transuranics that may precipitate in
ammonaical solution at this pH.
Nickel forms a weak complex with chloride ion as the tetrachloronicollate (+2) anion. This forms
the basis of its separation from other first row transition elements iron and cobalt. The complex,
Ni+2 + 4Cl! 6 NiCl4!2, is only stable in solutions greater than 10 M in HCl (see ion exchange
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section). Nickel forms complexes with the chelating agent diphenylthiocarbazone, which can be
extracted into organic solvents to form the basis of a separation form other transition metals.
Dissolution of Samples
Samples containing nickel radionuclides are most likely to be corrosion products, pure metals
that have been irradiated, or environmental water or soil samples. Dissolution of nickel and its
compounds from these matrices can be achieved using any combination of concentrated mineral
acids.
Separation Methods
PRECIPITATION. The classical method of nickel determination by gravimetric analysis is through
precipitation with dimethylglyoxime (DMG). This material is very specific to nickel and forms a
crystalline precipitate that is easily dried and weighed. The precipitation is carried out at pH 2-3,
in the absence of other macroscopic metal contaminants. Aluminum, iron, and chromium can
interfere but can be sequestered at pH 7–10 in ammoniacal solution with added citrate or tartrate.
The Ni-DMG precipitate may be dried, weighed, and the mass used as the determination for yield
of added nickel carrier.
SOLVENT EXTRACTION. Among the many solvent extraction methods for nickel, the following
compounds are notably efficient: Cupferron, acetylacetone, TTA, dibenzoylmethane, and
8-hydroxyquinoline. The extractions almost uniformly are most effective at pH 5–10. Unfortunately, in each of these separation techniques, the most effective solvents are chloroform,
benzene, or carbon tetrachloride, all of which have been phased out as analytical aids in
separation analysis.
ION EXCHANGE. Nickel can be separated from other transition metals on an anion exchange
column by dissolution of the sample in 12 M HCl. After the sample is loaded onto the column,
lowering the HCl concentration to 10 M will elute the nickel.
Nickel also can be separated from cobalt in oxalate media using a cation exchange resin. The
cobalt forms an anionic complex with the oxalate while the nickel does not. The cobalt passes
through the resin and the nickel is retained.
Methods of Analysis
The 59Ni and 63Ni isotopes do not emit gamma radiation. Liquid scintillation or proportional
counting after radiochemical separation can determine both isotopes. Nickel-59, as a very thin
test source, also can be determined using a low energy gamma/X-ray detector. It decays by
electron capture, and yields a characteristic X-ray of 6.93 keV. In a 63Ni analysis, if 59Ni is
present in the test source, a correction for the liquid scintillation yield of the 59Ni will be
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necessary. Chemical yield is determined by using a stable carrier and gravimetric analysis or
spectrophotometric techniques.
Compiled from: Cotton and Wilkinson, 1966; Freiser, 1983; Kraus and Nelson, 1958;
Minczewski et al., 1982.
14.10.9.8 Plutonium
Plutonium, with an atomic number of 94, is an actinide and the second element in the transuranic
series. Essentially all plutonium is an artifact, most produced by neutron bombardment of 238U
followed by two sequential beta emissions, but trace quantities of plutonium compounds can be
found in the natural environment. Plutonium radiochemistry is complicated by the five possible
oxidation states that can exist; four can be present in solution at one time.
Isotopes
Plutonium has 18 isotopes with mass numbers ranging from 232 to 247, and all isotopes are
radioactive. Some have a long half-life: the isotope of greatest importance, 239Pu, has a half-life
of 24,110 years, but 242Pu and 244Pu have a half-lives of 376,000 and 76,000,000 years, respectively. Plutonium-238, 240Pu, and 241Pu have a half-lives of 87.74, 6,537, and 14.4 years, respectively. Four of these isotopes decay by alpha emission accompanied by weak gamma rays: 238Pu,
239
Pu, 240Pu, and 242Pu. In contrast, 241Pu decays by beta emission with weak gamma rays, but its
progeny is 241Am, an intense gamma emitter. Plutonium-239 and 241Pu are fissile materials—they
can be split by both fast and slow neutrons. Plutonium-240, and 242Pu are fissionable but have
very small neutron fission cross-sections. Plutonium-240 partly decays by spontaneous fission,
although a small amount of spontaneous fission occurs in most plutonium isotopes.
Occurrence and Uses
There are minute quantities of plutonium compounds in the natural environment as the result of
thermal neutron capture and subsequent beta decay of naturally occurring 238U. All plutonium of
concern is an artifact, the result of neutron bombardment of uranium in a nuclear reactor.
Virtually all nuclear power-plants of all sizes and the waste from the plants contain plutonium
because 238U is the main component of fuel used in nuclear reactors. It is also associated with the
nuclear weapons industry and its waste. Virtually all the plutonium in environmental samples is
found in air samples as the result of atmospheric weapons testing. Plutonium in plant and crop
samples is essentially caused by surface absorption.
Plutonium is produced in nuclear reactors from 238U that absorbs neutrons emitted by the fission
of 235U, which is a naturally occurring uranium isotope found with 238U. Uranium-239 is formed
and emits a beta particle to form 239Np that decays by beta emission to form 239Pu. Once started,
the process is spontaneous until the uranium fuel rods become a specific uranium-plutonium
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mixture. The rods are dissolved in acid, and plutonium is separated primarily by solvent
extraction, finally producing a concentrated plutonium solution. Pure plutonium metal can be
prepared by precipitating plutonium peroxide or oxalate, igniting the precipitate to PuO2,
converting the oxide to PuF3, and reducing Pu+3 to the metal in an ignited mixture containing
metallic calcium.
Large quantities of 239Pu have been used as the fissile agent in nuclear weapons and as a reactor
fuel when mixed with uranium. It is also used to produce radioactive isotopes for research,
including the study of breeder reactors, and 238Pu is used as a heat source to power instruments
for space exploration and implanted heart pacemakers.
Solubility of Compounds
General solubility characteristics include the insolubility of the hydroxides, fluorides, iodates,
phosphates, carbonates, and oxalates of Pu+3 and Pu+4. Some of these can be dissolved in acid
solution, however. The corresponding compounds of PuO2+1 and PuO2+2 are soluble, with the
exception of the hydroxides. The binary compounds represented by the carbides, silicides,
sulfides, and selenides are of particular interest because of their refractory nature. One of the
complicating factors of plutonium chemistry is the formation of a polymeric material by hydrolysis in dilute acid or neutral solutions. The polymeric material can be a complicating factor in
radiochemical procedures and be quite unyielding in attempts to destroy it.
Review of Properties
Plutonium metal has some unique physical properties: a large piece is warm to the touch because
of the energy produced by alpha decay, and it exists in six allotropic forms below its melting
point at atmospheric pressure. Each form has unusual thermal expansion characteristics that
prevents the use of unalloyed plutonium metal as a reactor fuel. The delta phase, however, can be
stabilized by the addition of aluminum or gallium and be used in reactors. Chemically, plutonium
can exist in five oxidation states: +3, +4, (V), (VI), and (VII). The first four states can be
observed in solution, and solid compounds of all five states have been prepared. The metal is a
silver-grey solid that tarnishes in air to form a yellow oxide coating. It is chemically reactive
combining directly with the halogens, carbon, nitrogen, and silicon.
Plutonium is a very toxic substance. Outside the body, however, it does not present a significant
radiological hazard, because it emits only alpha, low-energy beta, gamma, or neutron radiation.
Ingested plutonium is not readily absorbed into the body, but passes through the digestive tract
and expelled before it can cause significant harm. Inhaled plutonium presents a significant
danger. Particularly, inhalation of particles smaller than one micron would be a serious threat due
to the alpha-emitting radionuclide being in direct contact with lung tissue. Plutonium would also
be very dangerous if it were to enter the blood stream through an open wound, because it would
concentrate in the liver and bones, leading to damage to the bone marrow and subsequent related
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problems. For these reasons, plutonium is handled in gloveboxes with associated precautions
taken to protect the worker from direct contact with the material. When working with plutonium
in any form, precautions should also be taken to prevent the accumulation of quantities of
fissionable plutonium that would achieve a critical mass, particularly in solution where it is more
likely to become critical than solid plutonium.
Most of the plutonium in the environment is the result of weapons testing. More than 99 percent
of the plutonium from these activities was released during atmospheric tests, but a small portion
was also released during ground tests. An even smaller quantity is released by nuclear fuel
reprocessing plants, some in the ocean, and by nuclear waste repositories. Part of the atmospheric
plutonium, originally part of the weapons, settled to the Earth as an insoluble oxide, locating in
the bottom sediments of lakes, rivers, and oceans or becoming incorporated in sub-surface soils.
The majority of environmental plutonium isotopes are the result of atmospheric nuclear bomb
tests. If the bomb material is made from uranium, the oxide is enriched to high percentages of
235
U, the fissile isotope. The 238U isotope does not fission, but absorbs 1–2 neutrons during the
explosion forming isotopes of 239U and 240U. These isotopes beta decay within hours to their
neptunium progeny, which in turn decay to 239Pu and 240Pu. Bombs made from plutonium would
yield higher fractions of 240/241/242Pu.
Plutonium formed as a result of atmospheric tests is most likely to be in the form of a fine
particulate oxide. If as in the case of a low altitude or underground test, there is a soil component,
the plutonium will be fused with siliceous minerals. The behavior of the soluble form of
plutonium would be similar to that released from fuel reprocessing plants and from nuclear waste
sites. Like the insoluble oxide, most of the soluble form is found in sediments and soils, but a
small percentage is associated with suspended particles in water. Both the soluble form of
plutonium and the form suspended on particulate matter are responsible for plutonium transportation in the environment. Plutonium in soil is found where the humic acid content is high. In nonhumic, carbonate-rich soils, plutonium migrates downward. Migration in the former soil is slow
(#0.1 cm/y) and in the latter it is relatively fast (1–10 cm/y). In subsurface oxic soil, plutonium is
relatively mobile, transported primarily by colloids. In wet anoxic soils, most of the plutonium is
quickly immobilized, although a small fraction remains mobile. The average time plutonium
remains in water is proportional to the amount of suspended material. For this reason, more than
90 percent of plutonium is removed from coastal water, while the residence time in mid-ocean
water where particulate matter is less is much longer.
Solution Chemistry
The equilibration problems of plutonium are among the most complex encountered in radiochemistry. Of the five oxidation states that plutonium may have, the first four are present in
solution as Pu+3, Pu+4, PuO2+1, PuO2+2. They coexist in dilute acid solution, and sometimes all
four are present in substantial quantities. Problems of disproportionation and auto-oxidation in
freshly prepared solutions also complicate the chemistry of plutonium. The (VII) state can form
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in alkaline solutions, and it has been suggested that the ion in solution is PuO5!3. Plutonium ions
tend to hydrolyze and form complex ions in solution. The +4 ion can form long chain polymers
that do not exhibit the usual chemical behavior of the +4 oxidation state. Finally, the different
oxidation states exhibit radically different chemical behavior. As a result of these effects, it is
possible to mix a plutonium sample with plutonium tracer, subject the mixture to a relatively
severe chemical treatment using hot acids or similar reagents, and still selectively recover
portions of either the tracer or the sample. This characteristic explains the challenge in achieving
reproducible radiochemical results for plutonium.
OXIDATION-REDUCTION BEHAVIOR. Numerous redox agents are available to oxidize and reduce
any of the five states of plutonium to alternate oxidation states. Table 14.19 provides a
convenient method of preparation of each state and illustrates the use of redox reagents in
plutonium chemistry.
Table 14.19 — Redox agents in plutonium chemistry
Oxidation State Form

Method of Preparation

+3

Pu+3

Dissolve Pu metal in HCl and reduce Pu+4 with NH2OH, N2H4,
SO2, or by cathodic reduction

+4

Pu+4

Oxidize Pu+3 with hot HNO3; treat Pu+3 or PuO2+2 with NO2!1

+4

PuO2@nH2O
(polymer)

Heat Pu+4 in very dilute acid; peptize Pu(OH)4

V

PuO2+1

Reduce PuO2+2 with stoichiometric amount of I!1 or ascorbic acid;
electrolytic reduction of PuO2+2

VI

PuO2+2

Oxidize Pu+4 with hot dilute HNO3 or AgO; ozonize Pu+4 in cold
dilute HNO3 with Ce+3 or Ag+1 catalyst

VII

PuO5!3

Oxidize PuO2+2 in alkali with O3, S2O8!2 or radiation

Unlike uranium, the +3 oxidation state is stable enough in solution to be useful in separation
chemistry. Disproportionation reactions convert Pu+4 to Pu+3 and PuO2+2 releasing H+1. The
presence of acid in the solution or complexing agents represses the process. Similarly, PuO2+1
disproportionates producing the same products but with the consumption of H +1. For this reason,
PuO2+1 is not predominant in acid solutions. These disproportionation reactions can be involved
in redox reactions by other reagents. Instead of direct oxidation or reduction, the disproportionation reaction can occur first, followed by direct oxidation or reduction of one of the products.
It is possible to prepare stable aqueous solutions in which appreciable concentrations of the first
four oxidation states exist simultaneously: the +3, +4, (V), and (VI) states. The relative
proportions of the different oxidation states depend on the acid, the acid concentration, the
method of preparation of the solution, and the initial concentrations of each of the oxidation
states. These relative concentrations will change over time and ultimately establish an
equilibrium specific to the solution. In 0.5 M HCl at 25 EC, for example, the equilibrium
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percentages of the four oxidation states prepared from initially pure Pu+4 are Pu+3 (27.2%), Pu+4
(58.4%), Pu(V) (0.7%), and Pu(VI) (13.6%). Freshly prepared plutonium samples are frequently
in the +4 state, while an appreciable amount of the +3 and +6 oxidation states will be present in
long-standing tracer solutions.
A convenient solution to this plutonium equilibration problem takes the form of a two-step
process:
• Boil the combined sample and tracer with a concentrated inorganic acid (e.g., HNO3) to
destroy any +4 polymers that might have formed, and
• Cool and dilute the solution; then rapidly (to avoid reforming polymers) treat the solution
with excess iodide ion (solution turns brown or black) to momentarily reduce all of the
plutonium to the +3 oxidation state.
The solution will immediately start to disproportionate in the acid medium, but the plutonium
will have achieved a true equilibrium starting at a certain time from one state in the solution.
Alpha particles emitted by 239Pu can decompose solutions of the radionuclide by radiolysis. The
radiolysis products then oxidize or reduce the plutonium, depending on the nature of the solution
and the oxidation state of the element. The nature of the anion present greatly influences the rate
of the redox process. For the radiochemist it is important to recognize that for old plutonium
solutions, particularly those in low acidity, the oxidation labeled states are not reliable.
HYDROLYSIS AND POLYMERIZATION. Hydrolysis is most pronounced for relatively small and
highly charged ions such as Pu+4, but plutonium ions in any oxidation state are more easily
hydrolyzed than their larger neptunium and uranium analogues.
Trivalent plutonium tends to hydrolyze more than neptunium or uranium, but the study of its
hydrolysis characteristics has been hindered by precipitation, formation of Pu+4, and unknown
polymerization. In strongly alkaline solutions, Pu(OH)3 precipitates; the solubility product
constant is estimated to be 2×10!20.
Plutonium (+4) exists as a hydrated ion in solutions that are more acidic than 0.3 M H+1. Below
0.3 M, it undergoes much more extensive hydrolysis than any other plutonium species, or at
lower acidities (0.1 M) if the plutonium concentration is lower. Thus, the start of hydrolysis
depends on the acid/plutonium ratio as well as the temperature and presence of other ions. On
hydrolysis, only Pu(OH)+3 is important in the initial phases, but it tends to undergo irreversible
polymerization, forming polymers with molecular weights as high as 1010 and chemical
properties much different from the free ion. Presence of the polymer can be detected by its bright
green color. When Pu+4 hydroxide [Pu(OH)4] is dissolved in dilute acid, the polymer also forms.
Similarly, if a solution of Pu+4 in moderately concentrated acid is poured slowly into boiling
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water, extensive polymerization occurs. The colloidal character of the polymer is manifested by
its strong adsorption onto glass, silica, or small bits of paper or dirt. The chemical characteristics
of the polymer, with regard to precipitation, ion-exchange, and solvent extraction, is markedly
different than the chemistry of the common +4 oxidation state of plutonium. Care should be
taken in the laboratory to avoid the formation of these polymers. For instance, these polymers can
be formed by overheating solutions during evaporation. Moreover, diluting an acidic plutonium
solution with water can cause polymerization because of localized areas of low acidity, even
when the final concentration of the solution is too high for polymerization. Therefore, plutonium
solutions should always be diluted with acid rather than water. Polymeric plutonium can also be
formed if insufficient acid is used when dissolving Pu+4 hydroxide.
Immediately after formation, these polymers are easy to decompose by acidification with
practically any concentrated inorganic acid or by oxidation. Because depolymerization is slow at
room temperature and moderate acid concentrations, solutions should be made at least 6 M and
boiled to destroy the polymers. The polymer is rapidly destroyed under these conditions. Adding
strong complexing agents such as fluoride, sulfate, or other strong complexing agents can
increase the rate of depolymerization. However, if the polymers are allowed to “age,” they can be
very difficult to destroy.
The PuO2+1 ion has only a slight tendency to hydrolyze, beginning at pH 8, but study of the extent
of the process is inhibited by the rapid disproportionation of hydrolyzed plutonium(V).
Hydrolysis of PuO2+2 is far more extensive than expected for a large +2 ion. Hydrolysis begins at
pH of about 2.7 to 3.3, giving an orange color to the solution that yields to bright yellow by pH 5.
Between pH 5 and 7, dimerizatons seem to occur, and by pH 13 several forms of plutonium
hydroxide have been precipitated with solubility products of approximately 2.5×10!25.
COMPLEXATION. Plutonium ions tend to form complex ions in the following order:
Pu+4 > Pu+3 . PuO2+2 > PuO2+1
Divalent anions tend to form stronger complexes, and the order for simple anions with Pu+4 is:
carbonate > oxalate > sulfate > fluoride > nitrate >
chloride > bromide > iodide > perchlorate
Complexation is preferably through oxygen and fluorine rather than nitrogen, phosphorus, or
sulfur. Plutonium also forms complexes with ligands such as phosphate, acetate, and TBP.
Strong chelate complexes form with EDTA, tartrate, citrate, TTA, acetylacetone (acac), and
cupferron. Pu+4 forms a strong complex with fluoride (PuF+3) that is used to solubilize plutonium
oxides and keep it in the aqueous phase during extraction of other elements with organic
solvents. The complex with nitrate, Pu(NO3)6!2, allows the recovery of plutonium from nuclear
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fuels. Carbonate and acetate complexes prevent precipitation of plutonium from solution even at
relatively high pH.
Dissolution of Samples
Metallic plutonium dissolves in halogen acids such as hydrochloric acid, but not in nitric or
concentrated sulfuric acids. The metal dissolves in hydrofluoric nitric acid mixtures. Plutonium
oxide dissolves with great difficulty in usual acids when ignited. Boiling with concentrated nitric
acid containing low concentrations of hydrofluoric acid or with concentrated phosphoric acid is
used. Fusion methods have also been used to dissolve the oxide as well as other compounds of
plutonium. Plutonium in biological samples is readily soluble, in the case of metabolized
plutonium in excreted samples, or highly refractory, in the case of fallout samples. Most
procedures for fallout or environmental samples involve treatment with hydrofluoric acid or
fusion treatment with a base.
Separation Methods
Extensive work has been done on methods to separate plutonium from other elements. Both
laboratory and industrial procedures have received considerable treatment. The methods
described below represents only a brief approach to separation of plutonium, but they indicate the
nature of the chemistry employed.
PRECIPITATION AND COPRECIPITATION. Macro quantities of plutonium are readily precipitated
from aqueous solution, and the methods are the basis of separating plutonium from other
radionuclides in some procedures. Contamination of other metals can be a problem, however;
zirconium and ruthenium give the most trouble. Plutonium is precipitated primarily as the
hydroxide, fluoride, peroxide, or oxalate. Both Pu+3 and Pu+4 are precipitated from acid solution
by potassium or ammonium hydroxide as hydrated hydroxides or hydrous oxides. On
redissolving in acid, Pu+4 tends to form the polymer, and high concentration of acid is needed to
prevent its formation. Pu+4 peroxide is formed on the addition of hydrogen peroxide to Pu+3, Pu+4,
Pu(V), and Pu(VI) because of the oxidizing nature of hydrogen peroxide. The procedure has been
used to prepare highly pure plutonium compounds from americium and uranium.
Coprecipitation of plutonium can be very specific with the control of its oxidation states and
selection of coprecipitating reagents. Lanthanum fluoride, a classical procedure for coprecipitation of plutonium, will bring down Pu+3 and Pu+4 but not Pu(VI). Only elements with similar
redox and coprecipitation behavior interfere. Separation from other elements as well as
concentration from large volumes with lanthanum fluoride is also important because not many
elements form acid-soluble lanthanum fluoride coprecipitates. Bismuth phosphate (BiPO4) is also
used to coprecipitate Pu+3 and Pu+4. In contrast to lanthanum fluoride and bismuth phosphate,
zirconium phosphate [Zr3 (PO4)4] and an organic coprecipitate, zirconium phenylarsenate
[Zr(C6H5)AsO4], will coprecipitate Pu+4 exclusively.
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SOLVENT EXTRACTION. A wide variety of organic extractants have been developed to separate
plutonium from other radionuclides and metals by selectively extracting them from aqueous
media. The extractants, among others, include organophosphorus compounds such as phosphates
(organoesters of phosphoric acid), amines and their quaternary salts, alcohols, ketones, ethers,
and amides. Chelating agents such as TTA and cupferron have also been used. Numerous studies
have been performed on the behavior of these systems. It has been found that the performance of
an extracting system is primarily related to the organic solvent in which the extractant is
dissolved and the concentration of the extractant in the solvent, the nature of the aqueous
medium (the acid present and its concentration [pH] and the presence of salting agents), the
temperature of the system, and the presence and nature of oxidizing agents. One common system,
used extensively in the laboratory and in industrial process to extract plutonium from fission
products, illustrates the use of solvent extraction to separate plutonium from uranium and other
metals. The PUREX process (plutonium uranium reduction extraction) is used in most fuel
reprocessing plants to separate the radionuclides. It employs TBP, tri-n-butyl phosphate
[(C4H9O)3PO], in a hydrocarbon solvent, as the extractant. The uranium fuel is dissolved in nitric
acid as Pu+3, and plutonium is oxidized to Pu+4 and uranium to U(VI) by oxidizing agents.
Plutonium and uranium are extracted into a 30 percent TBP solution, and the organic phase is
scrubbed with nitric acid solution to remove impurities. The plutonium is removed by backextracting it as Pu+3 with a nitric acid solution containing a reducing agent.
Solvent extraction chromatography, which uses an inert polymeric material as the support for
adsorbed organic chelating agents, has provided an efficient, easy technique for rapidly
separating plutonium and other transuranic elements. A process using CMPO in TBP and fixed
on an inert polymeric resin matrix has been used to isolate Pu+4. Aliquat-336® also has been used
successfully. All plutonium in the analyte is adjusted to Pu+4, and the column is loaded from 2 M
nitric acid. Plutonium is eluted with 4 M hydrochloric acid and 0.1 M hydroquinone or 0.1 M
ammonium hydrogen oxalate (NH4HC2O4). Environmental samples contain Fe3+ that may
interfere with this process and subsequently interfere with the analysis for plutonium. Ascorbic
acid can be used to reduce Fe+3 to Fe+2, which also reduces Pu+4 to Pu+3. Alternatively, nitrite may
be added after the ascorbic acid, which will not oxidize the iron but will convert the Pu+3 to Pu+4.
This process is an example of selective oxidation-reduction of plutonium and iron, and is used in
many different separation schemes for plutonium, including separation from americium.
ION-EXCHANGE CHROMATOGRAPHY. Ion-exchange chromatography has been used extensively
for the radiochemical separation of plutonium. All cationic plutonium species in noncomplexing
acid solutions readily exchanges onto cation resins at low acid concentrations and desorb at high
acid concentrations. Plutonium in all its oxidation states form neutral or anionic complexes with
various anions, providing an alternate means for eluting the element. Various cation-exchange
resins have been used with hydrochloric, nitric, perchloric, and sulfuric acids for separation of
plutonium from metals including other actinides. The most common uses of plutonium cationexchange chromatography is concentrating a dilute solution or separating plutonium from nonexchangable impurities, such as organic or redox agents.
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Anion-exchange chromatography is one of the primary methods for the separation of plutonium
from other metals and the separation of the plutonium oxidation states. On a strong anionexchange resin, for example, exchange of the higher oxidation states (+4, V, and VI) occurs at
hydrochloric acid concentrations above 6 M, while elution occurs at 2 M acid. Plutonium (+3)
does not absorb on the column, and Pu(VI) absorbs from 2 to 3 M hydrochloric acid solution.
Plutonium can be separated from other actinides and most other elements by exchanging the
plutonium cations—Pu+4 and Pu(VI)—onto a strong-anion resin from 6 M hydrochloric acid, and
subsequently eluting the plutonium by reducing it to Pu+3. Plutonium (+4) may be separated
effectively on anion exchange resin in 7-8 M nitric acid as the [Pu(NO3)6 ]!2 complex. Uranium
will elute slowly in this media, and sufficient volume must be processed in order to avoid cross
contamination of uranium with plutonium when the plutonium is subsequently eluted. Elution is
achieved at a lower acid concentration, or by reduction to Pu+3.
ELECTRODEPOSITION. Separation methods based on electrodeposition are not common, but one
method for the alpha analysis of plutonium is in use. Plutonium is electrodeposited on a stainless
steel disc from an ammonium sulfate solution at 1.2 amps for one hour. The separation is used
after isolating the radionuclide by extraction chromatography. This technique allows the
plutonium isotopes to be resolved by alpha spectroscopy.
Methods of Analysis
Once isolated, purified, and in solution, 238Pu, 239Pu, 240Pu, and 241Pu are collected for analysis
either by electrodepositon on a platinum or nickel disc or by microprecipitation with lanthanum
or neodymium fluoride. Mass spectrometry also can be used for longer-lived isotopes of
plutonium. Radionuclides of 238Pu, 239Pu, and 240Pu are determined by alpha spectrometry or gas
flow proportional counting. Plutonium-241 measured by gas proportional counting. Plutonium236 and 242Pu are used as tracers for measuring chemical yield.
When analyzing most samples containing 238Pu or 239Pu, the analyst can use either 236Pu or 242Pu
as a tracer. However, 242Pu should be avoided as a tracer when analyzing samples that inherently
contain 242Pu, such as waste generated by commercial nuclear reactors. When analyzing samples
that have higher (> 1 Bq) activity levels of 238Pu or 239Pu, most laboratories will use 236Pu as a
tracer because its higher-energy alpha-energy peaks (5.768 and 5.721 MeV) are well separated
from the lower energy peaks of 238Pu (highest alpha energy of 5.499 MeV) or 239Pu. Thus, the
isolated peaks of the 236Pu tracer can be quantified easily,1 and any minimum amount of 236Pu
peak tailing into the lower energy peaks of 238Pu or 239Pu (containing appreciably more counts)
will not significantly affect their quantification. However, when analyzing samples containing
very low concentrations of 238Pu or 239Pu (most environmental samples), 242Pu can be used as a

1

It should be noted that any contribution from a tracer into the peak(s) of an analyte of interest must be quantified
properly, and the affected analyte peak result corrected, to avoid a biased result or Type I error (false positive).

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tracer because its highest peak energy of 4.90 MeV is about 0.2 MeV lower than the lowest peak
energy of 238Pu or 239Pu. For such low activity samples, the 242Pu activity added to the sample
aliquant being processed should be more than the expected 238Pu or 239Pu test source activity.
Therefore, any tailing of the 239Pu alpha peaks into the 242Pu peaks would be minimized.
Compiled from: Baes and Mesmer, 1976; Choppin et al., 1995; Coleman, 1965; Cotton and
Wilkinson, 1988; DOE 1990 and 1997; EPA 1973 and 1980; Maxwell and Fauth, 2000; Metz
and Waterbury, 1962; Seaborg and Loveland, 1990; Weigel et al., 1986.
14.10.9.9 Radium
Radium, with an atomic number of 88, is the heaviest (last) member of the family of alkaline
earth metals, which, in addition, includes beryllium, magnesium, calcium, strontium, and barium.
Radium is the most alkaline and reactive of the series, and exists exclusively as +2 cations in
compounds and solution. All isotopes are radioactive, and essentially all analyses are made by
radioactive measurements or by mass spectrometry.
Isotopes
There are 25 isotopes of radium, from 205Ra to 234Ra. The most important with respect to the
environmental contamination are members of the 238U and 232Th naturally occurring decay series:
226
Ra and 228Ra, respectively. Radium-226 (t½ . 1,602 y) is the most abundant isotopic form. A
member of the 238U series, it is produced by alpha emission from 230Th. Radium-226 emits an
alpha particle and, in turn, produces 222Rn, an inert gas that is also an alpha emitter. Radium-226
generates radon at the rate of 0.1 µL per day per gram of radium, and its radioactivity decreases
at the rate of about 1 percent every 25 years. Radium-228 (t ½ . 5.77 y) is produced in the 232Th
decay series by emission of an alpha particle from 232Th itself.
Occurrence
In nature, radium is primarily associated with uranium and thorium, particularly in the uranium
ores—carnotite and pitchblende, where 226Ra is in radioactive equilibrium with 238U and its other
progeny. The widespread dispersal of uranium in rocks and minerals results in a considerable
distribution of radium isotopes throughout nature. Generally found in trace amounts in most
materials, the radium/uranium ratio is about 1 mg radium per 3 kg uranium (1 part radium in
3×106 parts uranium). This leads to a terrestrial abundance of approximately 10!6 ppm: 10!12 g/g
in rocks and minerals. Building materials, such as bricks and concrete blocks for example, that
contain mineral products also contain radium. With leaching from soil, the concentration is about
10!13 g/L in river and streams, and uptake in biological systems produces concentrations of 10-14
g/g in plants and 10!15 g/g in animals.
Uranium ores have been processed with hot mineral acids or boiling alkali carbonate to remove
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radium and uranium. Extracted radium was usually coprecipitated with barium sulfate, converted
to carbonate or sulfide, and solubilized with hydrochloric acid. Separation from barium was
usually accomplished by fractional crystallization of the chlorides, bromides, or hydroxides,
because barium salts are usually slightly more soluble. The free metal has been prepared by
electrolysis of radium chloride solutions, using a mercury cathode. The resulting amalgam is
thermally decomposed in a hydrogen atmosphere to produce the pure metal. The waste streams
from these industrial operations contain radium, primarily as a coprecipitate of barium sulfate.
Because many other natural ores also contain uranium and radium, processing can result in
uranium and its equilibrium progeny appearing in a product or byproduct. Apatite, a phosphate
ore, is used to produce phosphoric acid, and the gypsum byproduct contains all the radium
originally present in the ore.
Radium-226 extracted from ores has historically been used in diverse ways as a source of
radioactivity. It has been mixed with a scintillator to produce luminous paint, and at one time, the
most common use for its salts was radiation therapy. As a source of gamma radiation, radium
activity was enhanced by sealing a radium salt in a capsule that prevented escape of the gaseous
progeny, 222Rn, and allowing the radon to decay into its successive progeny. Two progeny are
214
Pb and 214Bi, the principal emitters of gamma radiation in the source. For the most part, radium
has been replaced in medical technology by other sources of radioactivity, but numerous capsules
containing the dry, concentrated substances still exist.
Radium salts are used in various instruments for inspecting structures such as metal castings by
gamma-ray radiography, to measure the thickness of catalyst beds in petroleum cracking units,
and to continuously measure and control the thickness of metals in rolling mills. Radium is also
used for the preparation of standard sources of radiation, as a source of actinium and protactinium, and as a source of ionizing radiation in static charge eliminators. In combination with
beryllium, it is a neutron source for research, in the analysis of materials by neutron activation,
and radio-logging of oil wells.
Radium in the environment is the result of natural equilibration and anthropological activity,
such as mining and processing operations. Radium is retained by many rock and soil minerals,
particularly clay minerals, and migrates only very slowly in through these materials. The decay
progeny of 226Ra, gaseous 222Rn, is an important environmental pollutant and represents the most
significant hazard from naturally occurring radium. Concentration of the alpha-emitting gas in
some occupied structures contributes to the incidence of lung cancer in humans. During the
decay of 226Ra, the recoil of the parent nucleus after it emits an alpha particle, now 222Rn, causes
an increased fraction of radon to escape from its host mineral, a larger fraction than can be
explained by intramineral migration or diffusion.
In groundwater, radium likely encounters dissolved sulfate and/or carbonate anions, which could
precipitate radium sulfate or radium carbonate. Although both salts are relatively insoluble, a
sulfate concentration of 0.0001 M would still allow an equilibrium concentration of about 0.1
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ppm Ra+2 to exist in solution. Thus, the insolubilities of either of these salts are not likely to
prevent contamination of the environment.
Radium also contaminates the environment because of past disposal practices of some processing, milling, and reclamation operations. Radium process tailings have been discovered in land
areas as seams or pockets of insoluble radium compounds, such as barium radium sulfate, or
unprocessed radium (uranium) ore, such as carnotite. Release of solid or liquid process streams
and subsequent mixing with local soil has resulted in intimate contamination of soil particles,
primarily as Ra+2 absorbed onto clay-sized fractions. This form of absorbed radium is tightly
bound to soil but can be extracted partially by hot concentrated acid solutions.
Solubility of Compounds
The solubility of radium compounds can usually be inferred from the solubility of the corresponding barium compound and the trend in the solubilities of the corresponding alkaline earth
compounds. The common water-soluble radium salts are the chloride, bromide, nitrate, and
hydroxide. The fluoride, carbonate, phosphate, biphosphate (hydrogen phosphate), and oxalate
are only slightly soluble. Radium sulfate is the least soluble radium compound known, insoluble
in water and dilute acids, but it is soluble in concentrated sulfuric acid, forming a complex ion
with sulfate anions, Ra(SO4)2!2.
Radium compounds are essentially insoluble in organic solvents. In most separation procedures
based on extraction, other elements, not radium, are extracted into the organic phase. Exceptions
are known (see “Separation Methods,” below), and crown ethers have been developed recently
that selectively remove radium from an aqueous environment.
Review of Properties
Radium is toxic exclusively because of its radioactive emissions: gamma radiation of the element
itself and beta particles emitted by some of its decay progeny. It concentrates in bones replacing
calcium and causing anemia and cancerous growths. Its immediate progeny, gaseous radon, is an
alpha emitter that is a health threat when inhaled.
Metallic radium is brilliant white and reacts rapidly with air, forming a white oxide and black
nitride. It is an active metal that reacts with cold water to produce radium hydroxide, hydrogen,
and other products. The radium ion in solution is colorless. Its compounds also are colorless
when freshly prepared but darken and decompose on standing because of the intense alpha
radiation. The original color returns when the compound is recrystallized. Alpha emissions also
cause all radium compounds to emit a blue glow in air when sufficient quantities are available.
Radium compounds also are about 1.5 EC higher in temperature than their surroundings because
of the heat released when alpha particles loose energy on absorbance by the compound. Glass
containers turn purple or brown in contact with radium compounds and eventually the glass
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crystallizes and becomes crazed.
Like all alkaline earths, radium contains two valence electrons (7s2) and forms only +2 ions in its
compounds and in solution. The ionic radius of radium in crystalline materials is 152 pm (0.152
nm or 1.52 D), the largest crystalline radius of the alkaline earth cations (Ra+2 > Ba+2 > Sr+2 >
Ca+2 >Mg+2 > Be+2). In contrast, the hydrated ion radius in solution is the smallest of the alkaline
earth cations, 398 pm (Be+2 > Mg+2 > Ca+2 > Sr+2 > Ba+2 > Ra+2). With the smallest charge-tocrystal-radius ratio among the alkaline earths of 1.32 (+2/1.52), the smallest hydrated radius of
radium is expected, because the ratio represents the least attractive potential for water molecules
in solution.
Solution Chemistry
Existing exclusively in the +2 oxidation state, the chemistry of radium is uncomplicated by
oxidation-reduction reactions that could produce alternate states in solution. It is made even less
complicated by its weak tendency to form complex ions or hydrolyze in solution. These
properties are a reflection of the small charge-to-crystal-radius ratio of 1.32, described above. In
general, radiochemical equilibrium is established with carriers by stirring, followed by either
standing or digesting in the cold for several minutes. Adsorption of trace amounts of radium on
surfaces, however, is an important consideration in its radiochemistry.
COMPLEXATION. Radium, like other alkaline-earth cations, forms few complexes in acid
solution. Under alkaline conditions, however, several one-to-one chelates are formed with
organic ligands: EDTA, diethylene triamine pentaacetic acid (DTPA), ethyleneglycol bis(2aminoethylether)-tetraacetate (EGTA), nitrilotriacetate (NTA or NTTA), and citrate. The most
stable complex ion forms with DTPA. The tendency to form complexes decreases as their
crystalline size increases and their charge-crystal-radius ratio decreases. Because crystalline sizes
of the cations are in the order: Ra+2 > Ba+2 > Sr+2 > Ca+2, radium has the least tendency to form
complex ions, and few significant complexes of radium with inorganic anions are known. One
notable exception is observed in concentrated sulfuric acid, which dissolves highly insoluble
radium sulfate (RaSO4) by forming Ra(SO4)2!2.
Complex-ion chemistry is not used in most radium radiochemical procedures. Complexing
agents are primarily employed as elution agents in cation exchange, in separations from barium
ions by fractional precipitation, and in titration procedures. Alkaline citrate solutions have been
used to prevent precipitation of radium in the presence of lead and barium carriers until complete
isotopic exchange has been accomplished.
HYDROLYSIS. Similar to their behavior complex-ion formation, alkaline earths show less and less
tendency to hydrolyze with increasing size of the ions, and the tendency decreases with
increasing ionic strength of the solution. Therefore, hydrolysis of radium is an insignificant factor
in their solution chemistry.
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ADSORPTION. The adsorption of trace amounts of radium on surfaces is an important consideration in its radiochemistry. Although not as significant with radium as with some ions with higher
charges, serious losses from solution can occur under certain conditions. Adsorption on glass is a
particular problem, and adsorption on polyethylene has been reported. Adsorption gradually
increases with increasing pH and depends strongly on the nature of the surface. In the extreme,
up to 50 percent radium has been observed to adsorb onto glass from neutral solution in 20 days,
and 30 percent from 0.13 M hydrochloric acid (HCl). Fortunately, adsorbed radium can be
removed from glass with strong acid.
The presence of insoluble impurities, such as traces of dust or silica, increases adsorption, but
adsorption is negligible from very pure solutions at low pH values. Tracer radium solutions,
therefore, should be free from insoluble impurities, and radium should be completely in solution
before analysis. The solutions should also be maintained in at least 1 M mineral acid or contain
chelating agents. Addition of barium ion as a carrier for radium will probably decrease the
amount of radium adsorption. Radium residues from solubilization of samples that contain silica
or lead or barium sulfates and those that result in two or more separate solutions should be
avoided, because the radium might divide unequally between the fractions. Destruction of silica
with HF, reduction of sulfates to sulfides with zinc dust, and subsequent dissolution of the
residue with nitric acid are procedures used to avoid this problem.
Dissolution of Samples
Soil, mineral, ore samples, and other inorganic solids are dissolved by conventional treatment
with mineral acids and by fusion with sodium carbonate (Na2CO3). Hydrofluoric acid (HF) or
potassium fluoride (KF) is used to remove silica. Up to 95 percent radium removal has been
leached from some samples with hot nitric acid (HNO3), but such simple treatment will not
completely dissolve all the radium in soil, rock, and mineral samples. Biological samples are wet
ashed first with mineral acids or decomposed by heating to remove organic material. The residue
is taken up in mineral acids or treated to remove silica. Any dissolution method that results in
two or more separate fractions should be avoided, because the adsorption characteristics of trace
quantities of radium may cause it to divide between the fractions.
Barium sulfate (BaSO4), often used to coprecipitate radium from solution, can be dissolved
directly into alkaline EDTA solutions. Radium can be repeatedly reprecipitated and dissolved by
alternate acidification with acetic acid and dissolution with the EDTA solution.
Solutions resulting from dissolution of solid samples should be made at least 1 M with mineral
acid before storage to prevent radium from absorbing onto the surface of glass containers.
Separation Methods
COPRECIPITATION. Radium is almost always present in solution in trace amounts, and even the
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most insoluble radium compound, radium sulfate, can not be used to separate and isolate radium
from solution by direct precipitation. Therefore, the cation is commonly removed from solution
in virtually quantitative amounts by coprecipitation. Because radium forms the same types of
insoluble compounds as barium: sulfates (SO4!2), chromates (CrO4!2), carbonates (CO3!2),
phosphates (PO4!3), oxalates (C2O4!2), and sulfites (SO3!2), it coprecipitates with all insoluble
barium compounds, and to a lesser extent with most insoluble strontium and lead compounds.
Barium sulfate and barium chromate are most frequently used to carry radium during coprecipitation. Other compounds that are good carriers for radium include: ferric hydroxide when
precipitated at moderately high pH with sodium hydroxide (NaOH) or ammonium hydroxide
(NH4OH), barium chloride (BaCl2) when precipitated from a cold mixed solvent of water and
alcohol saturated with hydrochloric acid, barium iodate [Ba(IO3)2], and various insoluble
phosphates, fluorides, and oxalates (e.g., thorium phosphate [Th3(PO4)4], lanthanum fluoride
(LaF3), and thorium oxalate [Th(C2O4)2]. Lead sulfate (PbSO4) can be used if a carrier-free
radium preparation is required, because quantitative lead-radium separations are possible while
quantitative barium-radium separations are very difficult.
ION EXCHANGE. Radium has been separated from other metals on both cation- and anionexchange resins. Barium and other alkaline earths are separated on cation-exchange columns
under acidic conditions. In hydrochloric acid solutions (3 M), the affinity of the cation for the
exchange site is dominated by ion-dipole interactions between the water molecules of the
hydrated ion and the resin. Ions of smaller hydrated radius (smaller charge-to-crystal-radius ratio)
tend to displace ions of larger hydrated radius. The affinity series is Ra+2 > Ba+2 > Sr+2 > Ca+2,
and radium elutes last. Increasing the acid concentration to 12 M effectively reverses the order of
affinity, because the strong acid tends to dehydrate the ion, and ion-resin affinity is dominated
more by ionic interactions, increasing in the order of increasing crystal radius: Ca+2 > Sr+2 > Ba+2
> Ra+2, and calcium elutes last. Radium has also been separated from tri- and tetravalent ions
because these ions have a much stronger affinity for the cation-exchange resin. Radium with its
+2 charge is only partially absorbed, while trivalent actinium and tetravalent thorium, for
example, will be completely absorbed. Tracer quantities of radium also has been separated from
alkaline earths by eluting a cation-exchange column with chelating agents such as lactate, citrate,
and EDTA; radium typically elutes last, because it forms weaker interactions with the ligands.
Anion-exchange resins have been used to separate radium from other metal ions in solutions of
chelating agents that form anionic complexes with the cations. The affinity for the columns
decreases in the order Ca > Sr > Ba > Ra, reflecting the ability of the metal ions to form stable
complex anions with the chelating agents. The difficult separation of barium from radium has
been accomplished by this procedure. Radium is also separated from metals such as uranium,
polonium, bismuth, lead, and protactinium that form polychloro complex anions. Because radium
does not form a chlorocomplex, it does not absorb on the anion exchanger (carrying a positive
charge), and remains quantitatively in the effluent solution.
Ion-exchange methods are not easily adapted for the separation of macro-scale quantities of
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radium, because the intense radiation degrades the synthetic resin and insoluble radium
compounds usually form in the ion-exchange column.
SOLVENT EXTRACTION. Radium compounds have very low solubilities in organic solvents. In
most extraction procedures, other organic-soluble complexes of elements, not radium, are
extracted into the nonaqueous phase, leaving radium in the water. Radium is separated from
actinium, thorium, polonium, lead, bismuth, and thallium, for example, by extracting these
elements as TTA complexes. Radium does not form the complex except at very high pH, and is
not extracted. One notable exception to this generality is the extraction of radium tetraphenylborate by nitrobenzene from an alkaline solution. The presence of EDTA inhibits formation of
the tetraphenylborate, however, and radium is not extracted in the presence of EDTA either.
More recent developments have employed crown ethers to selectively extract radium as a
complex ion from water samples for analysis. Radium-selective extraction membranes have also
been used to isolate radium from solutions.
Methods of Analysis
Radium is detected and quantified by counting either alpha or gamma emissions of the radionuclide or its progeny. Gamma-ray spectroscopy can be used on macro 226Ra samples (approximately
50 g or more) without pretreatment unless 235U, even in very small quantities, is present to interfere with the measured peak. The most sensitive method for the analysis of 226Ra is de-emanation
of 222Rn from the radium source, complete removal, followed by alpha counting the 222Rn and its
progeny. The procedure is lengthy and expensive, however. The radium in a liquid sample is
placed in a sealed tube for a specified time to allow the ingrowth of 222Rn. The radon is collected
in a scintillation cell and stored for several hours to allow for ingrowth of successive progeny
products. The alpha radiation is then counted in the scintillation cell called a Lucas cell. The
primary alpha emissions are from 222Rn, 218Po, and 214Po. Complete retention of radon can also be
accomplished by sealing the radium sample hermetically in a container and gamma-counting.
Radium-228 can also be determined directly by gamma spectroscopy, using the gamma-rays of
its progeny, 228Ac, without concern for interference. Alower detection limit is obtained if the
228
Ac is measured by beta counting. In the beta-counting procedure, 228Ra is separated, time is
allowed for actinium ingrowth, the 228Ac is removed by solvent extraction, ion-exchange, or
coprecipitation, and then measured by beta counting.
Radium-224 can be determined by chemically isolating the 212Pb, which is in equilibrium with
the 224Ra. After an appropriate ingrowth period, 212Pb is determined by alpha-, beta-, or gammacounting its progeny, 212Bi and 212Po.
Compiled from: Baes and Mesmer, 1976; Choppin et al., 1995; Considine and Considine,
1983; DOE, 1990 and 1997, 1997; EPA, 1984; Friedlander et al., 1981; Green and Earnshaw,
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1984; Hassinsky and Asloff, 1965; Kirby and Salutsky, 1964; Lindsay, 1988; Salutsky, 1997;
Sedlet, 1966; Shoesmith, 1964; Sunderman and Townley, 1960; Turekian and Bolter, 1966;
Vdovenko and Dubasov, 1975.
14.10.9.10 Strontium
Strontium, atomic number 38, is the fourth member of the alkaline-earth metals, which includes
beryllium, magnesium, calcium, strontium, barium, and radium. Like radium, it exists
exclusively in the +2 oxidation state in both compounds and in solution, making its chemistry
simpler than many of the radionuclides reviewed in this section.
Isotopes
Strontium exists in 29 isotopic forms, including three metastable states, ranging in mass number
from 77 to 102. Natural strontium is a mixture of four stable isotopes: 84Sr, 86Sr, 87Sr, and 88Sr.
The lower mass number isotopes decay by electron capture, and the isotopes with higher mass
numbers are primarily beta emitters. The half-lives of most isotopes are short, measured in
milliseconds, seconds, minutes, hours, or days. The exception is 90Sr, a beta emitter with a halflife of 29.1 years.
Occurrence and Uses
Strontium is found in nature in two main ores, celestite (SrSO4) and strontianite (SrCO3), widely
distributed in small concentrations. Small amounts are found associated with calcium and barium
minerals. The Earth’s crust contains 0.042 percent strontium, ranking twenty-first among the
elements occurring in rock and making it as abundant as chlorine and sulfur. The element ranks
eleventh in abundance in sea water, about 8–10 ppm. The only naturally occurring radioactive
isotopes of strontium are the result of spontaneous fission of uranium in rocks. Other nuclear
reactions and fallout from nuclear weapons test are additional sources of fission products.
Strontium-90 is a fission product of 235U, along with 89Sr, and short-lived isotopes, 91Sr to 102Sr.
Strontium-85 can be produced by irradiation of 85Rb with accelerated protons or deuterons.
The beta emission of 90Sr and its progeny, 90Y (t½ . 64 h), has found applications in industry,
medicine, and research. The radionuclides are in equilibrium in about 25 days. The radiation of
90
Y is more penetrating than that of strontium. It is used with zinc sulfide in some luminescent
paints. Implants of 90Sr provide radiation therapy for the treatment of the pituitary gland and
breast and nerve tissue. The radiation from strontium has been used in thickness gauges, level
measurements, automatic control processes, diffusion studies of seawater, and a source of
electrical power. Because 90Sr is one of the long-lived and most energetic beta emitters, it might
prove to be a good source of power in space vehicles, remote weather stations, navigational
buoys, and similar long-life, remote devices. Both 89Sr and 90Sr have been used in physical
chemistry experiments and in biology as tags and tracers. Ratios of 88Sr to 87Sr ratios are used in
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geological dating, because 87Sr is formed by decay of long-lived 87Rb.
Solubility of Compounds
Several simple salts of strontium are soluble in water. Among these are the acetate, chloride,
bromide, iodide, nitrate, nitrite, permanganate, sulfide, chlorate, bromate, and perchlorate.
Strontium hydroxide is slightly soluble and is precipitated only from concentrated solutions.
Review of Properties
Strontium is a low-density (2.54 g/cm 3) silver-white metal. It is as soft as lead and is malleable
and ductile. Three allotropic forms exit with transition temperatures of 235 and 540 EC. Freshly
cut strontium is silver in appearance, but it rapidly turns a yellowish color on formation of the
oxide in air. It is stored under mineral oil to prevent oxidation.
Strontium isotopes are some of the principal constituents of radioactive fallout following
detonation of nuclear weapons, and they are released in insignificant amounts during normal
operations of reactors and fuel reprocessing operations. Their toxicity is higher, however, than
that of other fission products, and 90Sr represent a particular hazard because of its long half-life,
energetic beta emission, tendency to contaminate food, especially milk, and high retention in
bone structure. Strontium in bone is difficult to eliminate and has a biological half-life of
approximately eleven years (4,000 d).
Strontium occurring in groundwater is primarily in the form of divalent strontium ions. Its
solubility under oxidizing and reducing conditions is approximately 0.001 M (0.15 g/L or 150
g/m3).
Solution Chemistry
Strontium exists exclusively in the +2 oxidation state in solution, so the chemistry of strontium is
uncomplicated by oxidation-reduction reactions that could produce alternate states in solution.
COMPLEXATION. Strontium has little tendency to form complexes. Of the few complexing agents
for strontium, the significant agents in radiochemistry to date are EDTA, oxalate, citrate,
ammoniatriacetate, methylanine-N,N-diacetate, 8-quinolinol, and an insoluble chelate with
picrolonate. The most stable complex ion forms with EDTA. Coordination compounds of
strontium are not common. These chelating agents are used primarily in ion-exchange
procedures. Amine chelates of strontium are unstable, and the β-diketones and alcohol chelates
are poorly characterized. In contrast, cyclic crown ethers and cryptates form stronger chelates
with strontium than with calcium, the stronger chelating metal with EDTA and more traditional
chelating agents. Cryptates are a macrocyclic chelate of the type, N[(CH2CH2O)2CH2CH2]3N, an
octadentate ligand containing six oxygen atoms and two nitrogen atoms as ligand bonding sites
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that encapsulates the cation. It might find use in the extraction chemistry of strontium.
HYDROLYSIS. The tendency of the alkaline-earth cations to hydrolyze decreases as their atomic
number increases. The tendency is greater than that of the corresponding alkali metals, but
hydrolysis of potassium, for example, is insignificant. An indication of the tendency of a cation
to hydrolyze is the solubility of their hydroxides, and the solubility of the alkaline earths
increases with increasing atomic number. Strontium hydroxide is slightly soluble in water (8 g/L
at 20 EC). In comparison, the hydroxide of beryllium, the first element in the alkaline earth
series, has a solubility of approximately 3×10!4 g/L.
Dissolution of Samples
Dissolution of samples for the analysis of strontium is generally simple. Water is used to dissolve
soluble compounds: acetate, bromide, chloride, iodide, chlorate, perchlorate, nitrate, nitrite, and
permangenate. Hydrochloric or nitric acid dissolves the fluoride, carbonate, oxalate, chromate,
phosphate, sulfate, and oxide. Strontium in limestone, cement, soil, bone, and other biological
material can be dissolved from some samples in hot hydrochloric acid. Insoluble silica, if present,
can be filtered or centrifuged. In some cases, soil can be leached to remove strontium. As much
as 99.5 percent of the strontium in some crushed soil samples has been leached with 1 M nitric
acid by three extractions. Soil samples have also been suspended overnight in ammonium acetate
at pH 7 to leach strontium. If leaching is not successful, soil samples can be dissolved by alkali
fusion of the ground powder with potassium hydroxide, nitrate, or carbonate. Strontium is taken
up from the residue in nitric acid. Biological materials such as plant material or dairy products
are solubilized by ashing at 600 EC and taking up milk residue in hot, concentrated hydrochloric
acid and plant residue in aqua regia. Wet ashing can be used by treating the sample with nitric
acid followed by an equal-volume mixture of nitric and perchloric acids. Human and animal bone
samples are ashed at 900 EC and the residue dissolved in concentrated hydrochloric acid.
Separation Methods
PRECIPITATION AND COPRECIPITATION. The common insoluble salts of strontium are the fluoride,
carbonate, oxalate, chromate, and sulfate. Most are suitable for radiochemical procedures, and
strontium separation have the advantage of stable forms of strontium that can be used as a carrier
and are readily available. Precipitation of strontium nitrate in 80 percent nitric acid has been used
to separate stable strontium carrier and 90Sr from its progeny, 90Y, and other soluble nitrates
(calcium, for example). The solubility of strontium chloride in concentrated hydrochloric
solution has been used to separate strontium from barium—barium chloride is insoluble in the
acid. Barium and radium (as coprecipitant) have been removed from strontium by precipitating
barium as the chromate at a carefully controlled pH of 5.5. Strontium chromate will not
precipitate unless the pH is raised. Strontium can also be separated from yttrium by precipitation
of the much less soluble yttrium hydroxide by raising an acid solution of the cations to a pH of
about 8 with ammonium hydroxide. Strontium hydroxide is slightly soluble and will not
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precipitate without high concentrations of hydroxide or strontium or both. Carrier-free strontium
is coprecipitated with ferric hydroxide, and lead sulfate is also used.
SOLVENT EXTRACTION. The application of organic solvents for separation of strontium from
other metals has not been extensive. TTA has been used to extract carrier-free strontium at a pH
greater than 10. At pH 5, 90Y is extracted with TTA from strontium, which remains in aqueous
solution. 8-hydroxyquinolinol in chloroform has also been used to extract strontium. The few
procedures that have been available are mainly used to separate the alkaline earths from each
other. A 1:1 mixture of ethyl alcohol and diethyl ether with di-2-ethylhexyl phosphoric acid
extracts calcium from strontium.
In recent years, extraction procedures have been developed based on the complexation of
strontium cations with crown ethers in 1-octanol. Strontium can be extracted with these mixture
from 1 M to 7 M nitric acid solutions. The most advantageous application of strontium extraction
procedures has been found in extraction chromatography. An extraction resin consisting of
4,4'(5')-bis(t-butylcyclohexano)-18-crown-6 (DtBuCH18C6) in 1-octanol on an inert polymeric
matrix is highly selective for strontium nitrate and will separate the cation from many other
metals including calcium, barium, and yttrium. This column is used to separate strontium from
potassium, cerium, plutonium, and neptunium (K+1, Ce+4, Pu+4, Np+4, respectively). The column
is prepared and loaded from 8 M nitric acid. The ions listed above are eluted with 3 M nitric acid
containing oxalic acid. Strontium is eluted with 0.05 M nitric acid.
ION-EXCHANGE CHROMATOGRAPHY. Ion-exchange chromatography is used to separate trace
quantities of strontium, but separation of macro quantities is very time consuming. Strontium is
absorbed on cation-exchange resins, and elution is often based on the formation of a stable
complex. Carrier-free strontium is separated from fission products, including barium, on a
cation-exchange resin and eluted with citrate. In a similar process, strontium was also separated
from other alkaline earths, magnesium, calcium, barium, and radium, eluting with ammonium
lactate at pH 7 and 78 EC. Good separations were also obtained with hydrochloric solutions and
ammonium citrate. Strontium-90 and 90Y are separated on a cation-exchange column, eluting
yttrium with ammonium citrate at pH 3.8 and strontium at pH 6.0. Strontium and calcium have
also been separated in EDTA solutions at pH 5.3. Strontium is retained on the column, and
calcium elutes as the calcium-EDTA complex. Strontium elutes with 3 M hydrochloric acid.
Strontium does not form many anionic complextes, Thus, not many procedures use anionexchange chromatography for separation of strontium. Strontium-90 has been separated from 90Y
on an anion-exchange resin pretreated with hydroxide. Strontium is eluted from the column with
water, and yttrium is eluted with 1 M hydrochloric acid. The alkaline earths have been separated
by anion-exchange column pretreated with dilute ammonium citrate, loading the column with the
chloride form of the metals, and eluting with ammonium citrate at pH 7.5.

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Methods of Analysis
Macroquantities of strontium are determined by gravimetric methods and atomic absorption
spectrometry, and emission spectrometry. Strontium is precipitated as strontium carbonate or
sulfate in gravimetric procedures. For atomic absorption analysis, the separated sample is ashed,
and the product is dissolved in hydrochloric acid. Lanthanum is added to the solution to
precipitate interfering anions, phosphate, sulfate, or aluminate, that would occur in the flame.
Strontium-89 and 90Sr are determined by analysis of their beta emissions. With a short half-life of
50.5 d, 89Sr is only found in fresh fission products. Strontium-90 is a beta emitter with a half-life
of 27.7 y. Its progeny is 90Y, which emits beta particles with a half-life of 64.0 h, producing
stable 90Zr. Neither 90Sr nor 90Y is a gamma emitter. Strontium-90 is determined directly from its
beta emission, before 90Y grows in, by beta counting immediately (three to four hours) after it is
collected by precipitation. The chemical yield can be determined gravimetrically by the addition
of stable strontium, after the separation of calcium. Alternatively, 90Sr can be measured from the
beta emission of 90Y while it reaches secular equilibrium (two to three weeks). The 90Y is
separated by solvent extraction and evaporated to dryness or by precipitation, then beta counted.
The chemical yield of the yttrium procedure can be determined by adding stable yttrium and
determining the yttrium gravimetrically. Strontium-89 has a half-life of 50.5 d and is only present
in fresh fission material. If it is present with 90Sr, it can be determined by the difference in
activity of combined 89Sr and 90Sr (combined or total strontium) and the activity of 90Sr. Total
strontium is measured by beta counting immediately after it is collected by precipitation, and 90Sr
is measured by isolating 90Y after ingrowth. Strontium-85 can be used as a tracer for determining
the chemical yield of 90Sr (determined by isolating 90Y), but its beta emission interferes with beta
counting of total strontium and must be accounted for in the final activity.
An alternative method for determining 89Sr and 90Sr in the presence of each other is based on the
equations for decay of strontium radionuclides and ingrowth of 90Y. Combined strontium is
collected and immediately counted to determine the total strontium. During ingrowth, the
mixture is recounted, and the data from the counts are used to determine the amount of 89Sr and
90
Sr in the original (fresh) mixture.
Cerenkov radiation counting techniques also may be used for 89/90Sr analysis. When beta particle
energies exceed the speed of light in the medium in which the beta particles are emitted, the
excess energy is emitted in the energy range of 350-600 nm. In water, the energy to be exceeded
is 0.263 MeV. As a practical matter, however, Cerenkov radiation counting is not very useful for
beta energies less than 1 MeV beta maximum (Eβmax) typically found in environmental
laboratories. NCRP (1985) cites a 3 percent detection efficiency for a 204Tl Eβmax of 0.764 MeV,
with corresponding average beta energy of 0.240 MeV. Only at a 143Pr of 0.932 MeV does the
detection efficiency go to 6.2 percent—a detection efficiency of marginal usefulness as a figure
of merit.
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The three isotopes that are involved with this analysis are 89Sr (Eβmax = 1.5 MeV), 90Sr (Eβmax = 0.5
MeV), and 90Y (Eβmax = 2.3 MeV). The analysis requires chemical separation of the strontium
from the sample matrix by conventional techniques. Cerenkov counting relies on the beta
energies (the 90Sr beta does not contribute significantly). For example, strontium may be
separated chemically as an oxalate precipitate (after yttrium has been removed by precipitation),
dissolved in nitric acid, and counted immediately (yielding the counts for 89Sr). After about 10
days, the sample would be recounted, yielding a total for 89Sr + 90Y. The value for the 90Y is then
determined by applying spectral interference factors for spectral overlap and appropriate
background subtraction techniques. Alternatively, 90Y can be separated from the strontium
solution after a period of ingrowth and Cerenkov-counted to determine the 90Sr concentration.
Compiled from: Baes and Mesmer, 1976; Banavali et al., 1995; Choppin et al., 1995;
Considine and Considine, 1983; CRC, 1998-99; DOE, 1990 and 1997, 1997; EPA, 1973;
EPA, 1980; Greenwood and Earnshaw, 1984; Hassinsky and Adloff, 1965; NCRP, 1985;
Riley, 1995; Rucker, 1991; Sunderman and Townley, 1960; Turekian and Bolter, 1966.
14.10.9.11 Sulfur and Phosphorus
The radiochemistry of sulfur and phosphorus is somewhat different than most other radioisotopes. These two elements are nonmetallic and, like carbon, can be found in many different types
of compounds. These two elements are used most extensively as tracers by incorporation into
organic molecules, generally as covalent-bonded atoms. Thus, they do not react as sulfur or
phosphorus, but as the molecule of which they are a part. They may be present as inorganic
species, which have their own peculiar chemistry.
Isotopes
Sulfur has 17 isotopes, four of which are stable. Only two of the 13 radioisotopes have
significant radiochemical analytical applications. These are 35S (t½ .87.2 d) and 37S (t½ . 5 min).
Sulfur-35 decays only by beta emission with no gamma emission. Sulfur-37 decays by beta
emission with a 3.1 MeV delayed gamma emission.
Phosphorus also has 17 isotopes, only one of which is stable. Its two principal radioisotopes, 32P
(t½ . 14.3 d) and 33P (t½ . 25.3 d), both decay only by beta emission, with no gamma emission.
Occurrence
None of the radioisotopes of sulfur occurs naturally. They are produced by neutron activation of
stable parent isotopes or by accelerator bombardment techniques. Both 32P and 33P are formed
naturally in the upper atmosphere. The steady-state concentration of these radionuclides in
rainwater is about 0.05 Bq/L. They are also produced artificially by accelerator bombardment.
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Solubility and Solution Chemistry
The most stable forms of the two elements in aqueous solutions are sulfate and phosphate.
However, the relatively long half-lives of the radioisotopes of S and P allow them to be
incorporated easily into organic or biomolecules. In these instances, the chemical identity of the
radioisotope is sacrificed for the chemical property of the molecule. For example, 35S may be
incorporated into these species, but each will have a distinct chemical property:
SO4!2 , S!2 , CH3-S-CH2 -CH2 -C(H)(NH2)(COOH) [methionine]
H-S-CH2-C(H)(NH2)(COOH) [cystine]
If a solution of methionine had added to it methionine labeled with 35S, the radioisotopecontaining molecules would be indistinguishable chemically from the other methionine
molecules. However, if the methionine solution was equilibrated with a solution of 35S!2, no 35S
would be found in the methionine molecules, because methionine does not dissociate to give S !2.
Similarly, for phosphorus the radioisotope could be incorporated into the following species:
PO43-, (C8H17)3PO [tri-n-octylphosphine oxide]
H2PO4-{C9H14N5O3} [adenosine-5-phosphoric acid].
Here, the tri-n-octylphosphine oxide is soluble in organic solvents but not in water, while the
other two are readily water-soluble. For the two water-soluble molecules, under conditions of
neutral pH, no exchange of radiophosphorus would be expected between them. However under
certain conditions where the organic molecule could be hydrolyzed, exchange could occur.
Incorporation of the radioisotope into an organic molecule would occur by first forming the
radioisotope by nuclear bombardment, then reacting the activated material with the appropriate
reagents to form the molecule of interest. Attempting to form the radioisotope by activation of
the organic molecule would lead to the destruction of the organic molecule, and the radioisotope
would be part of other (potentially) unknown species. The chemical purity of the final product
would be verified through an independent means such as infrared, nuclear magnetic resonance, or
mass spectrometry. The specific activity of the new molecule then can be calculated by
measuring the activity due to the radioisotope.
OXIDATION-REDUCTION. For each of these elements, the most stable ionic form in aqueous
solution is as the SO4!2 or the PO43- ions (dependent upon pH). Sample oxidation for sulfur
should be performed with care to avoid loss as SO2 or as H2S. This can occur in nitric acid when
sulfides or organic sulfur compounds are present. Oxidation in basic solution using hydrogen
peroxide or permanganate can avoid such losses. Phosphorus does not suffer from this
disadvantage of acid oxidation. Generally, when present as phosphate or sulfate, reduction to
other species will not occur unless powerful reducing agents have been added to the solution.
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COMPLEXATION. Neither sulfate nor phosphate are strong complexing agents. This is due to their
negative charge being spread out among many atoms, yielding low charge density. Most
complexing ions are strongly nucleophilic.
Dissolution of Samples
The radioisotopes of phosphorus and sulfur generally are incorporated into in vivo or in vitro
studies of plant or animal tissues. The cost common methods of sample preparation for these
studies usually are maceration/suspension, tissue solubilization, and total oxidation. The method
of maceration is a reduction of the “size” of the sample. The material is suspended in a minimal
amount of fluid, and then a physical means such as a blender, mortar and pestle, or stirring rod is
used to suspend the material in the solvent. The chemical nature of the molecule containing the
radioisotope is unchanged.
Tissue solubilization is the addition of a chemical solvent such as toluene, which dissolves the
tissue in its entirety putting the sample into an organic solvent matrix. The chemical nature of the
molecule containing the radioisotope is unchanged.
Total oxidation is performed most frequently using either peroxide or nitric acid, which removes
all of the organic material as carbon dioxide, and the elements are in solution as phosphate or
sulfate. Care should be taken in this form of sample preparation for sulfur, because it can be
volatilized as SO2 or SO3 vapor.
The molecules of interest having biochemical activity may change chemically during the course
of such studies. Thus, one should consider what the potential decomposition products are, and
how they should be separated from the organic/biomolecules of interest, before preparing the
sample. If an environmental sample were to be analyzed for these radioisotopes, the sample
preparation would need to be total-sample-oxidation, because the type of organic material would
likely be unknown.
Separation Methods
Because many different organic forms exist for these elements, it would be difficult to identify all
of the different separation techniques used to separate them from specific mixtures of other
organic compounds. Generally, the techniques that are used are HPLC, GC, and electrophoresis.
In many instances, separation of the molecules containing the radioisotopes is not necessary,
because the sulfur or phosphorus is the only radioisotope present, having been used as a tracer in
following the reaction progress or products.
PRECIPITATION. Sulfur may be analyzed by sample oxidation followed by barium precipitation.
This takes place at about pH 2 in HCl solution. As with other separation techniques, sample
processing should ensure the elimination of other cations (such as radium or strontium), which
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could be present in environmental samples.
Phosphate is a strong Bronsted-Lowry base. Precipitation of phosphate salts would be carried out
best in basic media. However, most metal salts also form insoluble hydroxides, so this form of
separation is not used frequently. However, if other metal ions are removed, phosphate can be
completely precipitated using calcium ion in basic solution.
ION EXCHANGE. Both phosphate and sulfate may be exchanged easily on anion exchange media.
However, if the anion resin were in the hydroxide form, the exchange would release hydroxide
and potentially cause precipitation of metal ions either on the ion exchange resin or in the eluent.
Thus, converting the anion resin to the nitrate or chloride form prior to separation would permit
the free flow of eluent without precipitation. Such separation will occur on weak base anion
exchangers (such as those used in ion chromatography) or strong base ion exchangers.
Methods of Analysis
All of the radioisotopes of interest of phosphorus and sulfur are beta emitters. The most effective
method of analysis for these isotopes is liquid scintillation. For the analysis of organic/
biomolecules, the scintillation cocktail usually may be added directly to the analyte after one of
the methods of nonoxidative sample preparation described above. In some instances, these
analytes may contain double-labeled compounds. Other radioisotopes, such as 14C or 3H, also
may be incorporated into the molecule. These can also be analyzed directly by liquid scintillation
because of the significant differences in the beta particle energies. Samples of unknown origin
would require oxidation and separation prior to analysis.
14.10.9.12 Technetium
Technetium, atomic number 43, has no stable isotopes. Natural technetium is known to exist but
only in negligibly small quantities resulting from the spontaneous fission of natural uranium.
Technetium is chemically very similar to rhenium, but significant differences exist that cause
them to behave quite differently under certain conditions.
Isotopes
Thirty-one radioisotopes of technetium are known with mass numbers ranging from 86 to 113.
The half-lives range from seconds to millions of years. The lower mass number isotopes decay by
primarily by electron capture and the higher mass number isotopes by beta emission. The
significant isotopes (with half-lives/decay modes) are 95mTc (61 d/electron capture and isomeric
transition), 99mTc (6.01 h/isomeric transition by low-energy γ), and 99Tc (2.13×105 y/β to stable
99
Ru). Other long-lived isotopes are 97Tc (2.6×106/electron-capture) and 98Tc (4.2×106 y/β
emission).
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Occurrence and Uses
The first synthesis of technetium was through the production of 99Mo by bombardment of 98Mo
with neutrons and subsequent beta decay to 99Tc. Technetium is also a major constituent of
nuclear reactor fission products and has been found in very small quantities in pitchblende from
the spontaneous fission of naturally occurring uranium.
Technetium makes up about 6 percent of uranium fission products in nuclear power plant fuels. It
is recovered from these fuels by solvent extraction and ion-exchange after storage of the fuels for
several years to allow the highly radioactive, short-lived products to decay. Technetium is
recovered as ammonium pertechnetate (NH4TcO4) after its solutions are acidified with
hydrochloric acid, precipitated with sulfide, and the sulfide (Tc2S7) is reacted with hydrogen
peroxide. Rhenium and molybdenum are also removed by extraction with organic solvents. The
metal is obtained by reduction of ammonium pertechnetate with hydrogen at 600 EC.
Potassium pertechnetates (KTcO4) have been used in water (55 ppm) as corrosion inhibitors for
mild carbon steel in aerated distilled water, but currently there is no significant uses of elemental
technetium or its compounds, although technetium and some of its alloys are superconductors.
The corrosion protection is limited to closed systems to prevent release of the radioactive isotope.
Technetium-95m, with a half-life of only 61 days, has been used in tracer work. Technetium-99m
is used in medical diagnosis as a radioactive tracer. As a complex, the amount of 99mTc required
for gamma scanning is very small, so it is referred to as noninvasive scanning. It is used for
cardiovascular and brain studies and the diagnosis of liver, spleen, and thyroid disorders. There
are more than 20 99mTc compounds available commercially for diagnostic purposes. With iodine
isotopes, they are the most frequently used radionuclides for diagnostics. Technetium-99m also
has been used to determine the deadtime of counting detectors.
Solubility of Compounds
The nature of the compounds has not been thoroughly delineated, but ammonium pertechnetate
is soluble in water, and technetium heptoxide forms soluble pertechnetic acid (HTcO4) when
water is added.
Review of Properties
Technetium is a silver-grey metal that resembles platinum in appearance. It tarnishes slowly in
moist air to give the oxyacid, pertechnetic acid (HTcO4). It has a density of 11.5 g/cm3. The metal
reacts with oxygen at elevated temperatures to produce the volatile oxide, technetium heptoxide.
Technetium dissolves in warm bromine water, nitric acid, aqua regia, and concentrated sulfuric
acid, but it is insoluble in hydrochloric and hydrofluoric acids. Technetium forms the chlorides
(TcCl4 and TcCl6) and fluorides (TcF5 and TcF6) by direct combination of the metal with the
respective halogen. The specific halide is obtained by selecting the proper temperature and
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pressure for its formation.
The behavior of technetium in groundwater is highly dependent on its oxidation state. Under
oxidizing conditions, pertechnetate is the predominant species. It is very soluble and only slightly
absorbed to mineral components. For those reasons, it has a relatively high dissemination
potential in natural systems. Under reducing conditions, technetium precipitates as technetium
dioxide (TcO2), which is very insoluble. With the production of 99Tc in fission fuels and
considering its long half-life, the soluble form of the radionuclide is an environmental concern
wherever the fuel is reprocessed or stored. As a consequence, 99Tc would be expected to be one
of the principal contributors to a radioactive release to the environment, even from repositories
with barriers that could retain the radionuclide up to 10,000 years. Studies of a salt repository
indicate that 99Tc is one of the few radionuclides that might reach the surface before it decays.
Solution Chemistry
All oxidation states between !1 and +7 can be expected for technetium, but the important ones in
solution are +4 and +7. The +4 state exist primarily as the slightly soluble oxide, TcO2. It is
soluble only in the presence of complexing ligands; TcCl6!2, for example, is stable in solutions
with a chloride concentration greater than 1 M. The most important species in solution is the
pertechnetate ion [TcO4!1 as Tc(VII)], which is readily soluble and easily formed from lower
oxidation states with oxidizing agents such as nitric acid and hydrogen peroxide. There is no
evidence of polymeric forms in solution as a result of hydrolysis of the metal ion.
OXIDATION-REDUCTION BEHAVIOR. Most radioanalytical procedures for technetium are
performed on the pertechnetate ion, TcO4!1. The ion can be reduced by hydrochloric acid, the
thiocyanate ion (SCN!1), organic impurities, anion-exchange resins, and some organic solvents.
The product of reduction can be TcO 2 [Tc+4], although a multiplicity of other products are
expected in complexing media. Even though the +7 oxidation state is easy to reduce, the
reduction process is sometimes slow. Unless precautions are taken to maintain the appropriate
oxidation state, however, erratic results will be obtained during the radioanalytical procedure.
Several examples illustrate the precaution. Dissolution should always be performed under
strongly oxidizing conditions to ensure conversion of all states to the +7 oxidation state because
complications because of slow exchange with carrier and other reagents are less likely to occur if
this state is maintained. Technetium is extracted with various solvents in several radioanalytical
procedures, but the method can be very inefficient because of reduction of the pertechnetate ion
by some organic solvents. The presence of an oxidizing agent such as hydrogen peroxide will
prevent the unwanted reduction. In contrast, TcO4!1 is easily lost on evaporation of acid solutions
unless a reducing agent is present or evaporation is conducted at a relatively low temperature.
COMPLEXATION. Technetium forms complex ions in solution with several simple inorganic
ligands such as fluoride and chloride. The +4 oxidation state is represented by the TcX6!2 ion
where X = F, Cl, Br, and I. It is formed from TcO4!1 by reduction to the +4 state with iodide in
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HX. TcF6!2 is found in HF solutions during decomposition of samples, before further oxidation.
Complex ions formed between organic ligands and technetium in the (V) oxidation state are
known with the general formula, TcO3XLL, where X is a halide and L is an organic ligand. the
ligands typically bond through an oxygen or nitrogen atom. Other organic complexes of the (V)
state have the general formulas: TcOX2L2, TcOX4!1, and TcOX5!2.
Dissolution of Samples
Dissolution of samples containing technetium requires two precautions: it is essential that acid
solutions be heated only under reflux conditions to avoid losses by volatilization, and dissolution
should be done only with strongly oxidizing conditions to ensure conversion of all lower
oxidation states to Tc(VII). In addition, problems with slow carrier exchange are less likely for
the (VII) oxidation state. Molybdenum targets are dissolved in nitric acid or aqua regia, but the
excess acid interferes with many subsequent analytical steps. Dissolution in concentrated sulfuric
acid followed by oxidation with hydrogen peroxide after neutralization avoids these problems of
excess acid. Other technetium samples can be dissolved by fusion with sodium peroxide/sodium
hydroxide (Na2O2/NaOH) fluxes.
Separation Methods
PRECIPITATION AND COPRECIPITATION. The various oxidation states of technetium are
precipitated in different forms with different reagents. Technetium(VII) is primarily present in
solution as the pertechnetate anion, and macro quantities are precipitated with large cations such
as thallium (Tl+1), silver (Ag+1), cesium (Cs+1), and tetraphenylarsonium [(C6H5)4As+1]. the
latter ion is the most efficient if ice-bath conditions are used. Pertechnetate is coprecipitated
without interference from molybdenum with these cations and perrhenate (ReO4!1), perchlorate
(ClO4!1), periodate (IO4!1), and tetrafluoroborate (BF4!1). The salt consisting of tetraphenylarsenium and the perrhenate froms a coprecipitate fastest, in several seconds. Technetium(VII) can
be precipitated from solution as the heptasulfide (Tc2S7) by the addition of hydrogen sulfide (or
hydrogen sulfide generating compounds such as thioacetamide and sodium thiosulfate) from 4 M
sulfuric acid. Because many other transition metals often associated with technetium also from
insoluble compounds with sulfide, the method is primarily used to concentrate technetium.
Technetium (+4) is carried by ferric hydroxide. The method can be use to separate technetium
from rhenium. The precipitate is solubilized and oxidized with concentrated nitric acid, and iron
is removed by precipitation with aqueous ammonia. Technetium is coprecipitated as the hexachlorotechnetate (+4) (TcCl6-2) with thallium, and rhenium as the α,α’-dipyridylhexachlororhenate (+4).
Technetium(VI) (probably as TcO4!2) is carried quantitatively by molybdenum 8-hydroxyquinolate and by silver or lead molybdate. Tc +3 is carried quantitatively by iron or zinc hydroxide and
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the sulfide, hydroxide, and 8-hydroxyquinolate of molybdenum.
SOLVENT EXTRACTION. Technetium, primarily in the Tc(VII) state (pertechnetate) can be isolated
by extraction with organic solvents, but the principal disadvantage of all extraction systems is the
inevitable introduction of organic material that might reduce the pertechnetate anion and cause
difficulties in subsequent analytical steps. The pertechnetate ion is extracted with pyridine from a
4 M sodium hydroxide solution, but perrhenate and permanganate ions are also extracted. The
anion also extracts into chloroform in the presence of the tetraphenylarsonium ion as tetraphenylarsonium pertechnetate. Extraction is more favorable from neutral or basic sulfate solutions than
chloride solutions. Perrhenate and perchlorate are also extracted but molybdenum does not
interfere. Small amounts of hydrogen peroxide in the extraction mixture prevent reduction of
pertechnetate. Technetium is back-extracted into 0.2 M perchloric acid or 12 M sulfuric acid.
Other organic solvents are have also been used to extract pertechnetate from acid solutions,
including alcohols, ketones, and tributyl phosphate. Ketones and cyclic amines are more effective
for extraction from basic solutions. Tertiary amines and quaternary ammonium salts are more
effective extracting agents than alcohols, ketones, and tributyl phosphate. Back extraction is
accomplished several ways, depending on the extraction system. A change in pH, displacement
by another anion such as perchlorate, nitrate, or bisulfate, or addition of a nonpolar solvent to an
extraction system consisting of an oxygen-containing solvent.
A recent extraction method has been used successfully for extraction chromatography and
extractive filtration. A column material consisting of trioctyl and tridecyl methyl ammonium
chlorides impregnated in an inert apolar polymeric matrix is used to separate 99Tc by loading the
radionuclide as the pertechnetate ion from a 0.1 M nitric acid solution. It is stripped off the
column most readily with 12 M nitric acid. Alternatively, the extraction material is used in a
filter disc, and the samples containing 99Tc are filtered from water at pH 2 and rinsed with 0.01
M nitric acid. Technetium is collected on the disc.
Lower oxidation states of technetium are possible. The thiocyanate complexes of technetium(V)
are soluble in alcohols, ethers, ketones, and trioctylphosphine oxide or trioctylamine
hydrochloride in cyclohexane or 1,2-dichloroethane. Technetium (+4), as TcCl6!2, extracts into
chloroform in the presence of high concentrations of tetraphenylarsonium ion. Pertechnetate and
perrhenate are both extracted from alkaline solution by hexone (methyl isobutyl ketone), but
reduction of technetium to the +4 state with hydrazine or hydroxylamine results in the extraction
of perrhenate only.
ION-EXCHANGE CHROMATOGRAPHY. Ion-exchange chromatography is primarily performed with
technetium as the pertechnetate anion. Technetium does not exchange on cation resins, so
technetium is rapidly separated from other cations on these columns. In contrast, it is strongly
absorbed on strong anion exchangers and is eluted with anions that have a greater affinity for the
resin. Technetium and molybdenum are separated using ammonium thiocyanate as the eluent. A
good separation of pertechnetate and molybdate has been achieved on an anion-exchange resin in
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the phosphate form where the molybdate is preferentially absorbed. Good separation of
pertechnetate and perrhenate are obtained with perchlorate as the eluent.
VOLATILIZATION. The volatility of technetium heptoxide allows the co-distillation of technetium
with acids. Co-distillation from perchloric acid gives good yields, but only a partial separation
from rhenium is achieved. Molybdenum is also carried unless complexed by phosphoric acid.
Separation from rhenium can be achieved from sulfuric acid, but yields of technetium are can be
very poor because of its reduction by trace impurities in the acid. Much more reproducible results
can be obtained in the presence of an oxidizing agent, but ruthenium tetroxide (RuO4) also
distills under these conditions. It can be removed, however, by precipitation as ruthenium dioxide
RuO2. In distillation from sulfuric acid-water mixtures, technetium distills in the low-boiling
point aqueous fraction, probably as pertechnetic acid. Technetium and rhenium are separated
from sulfuric-hydrochloric acid mixtures; pertechnetate is reduced to nonvolatile Tc+4 and
remains in the acid solution. Technetium heptoxide can be separated from molybdenum trioxide
by fractional sublimation at temperatures $ 300 EC.
ELECTRODEPOSITION. Technetium can be electrodeposited as its dioxide (TcO2) from 2 M
sodium hydroxide. The metal is partially separated from molybdenum and rhenium, but
deposition only occurs from low technetium concentrations. Carrier-free 95Tc and 96Tc have been
electrolyzed on a platinum electrode from dilute sulfuric acid. Optimum electroplating of
technetium has been achieved at pH 5.5 in the presence of very dilute fluoride ion. Yields were
better with a copper electrode instead of platinum—about 90 percent was collected in two hours.
Yields of 98–99 percent were achieved for platinum electrodes at pH 2-5 when the plating time
of up to 20 hours was used. In 2 M sulfuric acid containing traces of fluoride, metallic
technetium instead of the dioxide is deposited on the electrode.
Methods of Analysis
Technetium-99 is analyzed by ICP-MS, gas proportional counting, or liquid scintillation from its
beta emission. No gamma rays are emitted by this radionuclide. For ICP-MS analysis, technetium
is stripped from an extraction chromatography resin and measured by the spectral system. The
results should be corrected for interference by 99Ru, if present. For beta analysis, technetium can
be electrodeposited on a platinum disc and beta counted. Alternatively, it is collected by
extraction-chromatography techniques. The resin from a column or the disc from a filtration
system is placed in a liquid scintillation vial and counted. Technetium-99m (t1/2=6.0 h), measured
by gamma-ray spectrometry, can be used as a tracer for measuring the chemical yield of 99Tc
procedures. Conversion electron ejection from the tracer should then be subtracted from the total
beta count when measuring 99Tc. Alternatively, samples are counted immediately after isolation
and concentration of technetium to determine the chemical recovery, then the 99mTc is allowed to
decay before analysis of the 99Tc. A widely used medical application is the technetium generator.
Molybdenum-98 is neutron-irradiated and chemically oxidized to 99MoO4!2. This solution is ionexchanged onto an acid-washed alumina column. After about 1.25 days, the activity of 99mTc has
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grown-in to its maximum concentration. The 99Tc is eluted with a 0.9% solution of NaCl, while
the 99Mo remains on the column. The column may have its 99mTc removed after another 1.25
days, but at a slightly smaller concentration. The 99mTc thus separated is carrier free. This process
historically was referred to as “milking,” and the alumina column was called the “cow.”
Neutron activation analysis methods for technetium have been employed since 1972. A method
was developed and applied for the analysis of 99Tc in mixed fission products. The method
employs chemical separation of 99Tc from most fission products by a cyclohexanone extraction
from a basic carbonate solution. Technetium-99 is stripped into water by addition of CCl4 to the
cyclohexanone phase and then isolated on an anion exchange column. Neutron irradiation of the
isolated 99Tc was made in the pneumatic facility at a high flux beam reactor (e.g., at a flux of
5×1014 n·cm2/sec for approximately 11 seconds. Thus, after irradiation 99Tc is converted to 100Tc,
which, because of its 15.8 second half-life, requires an automatic process to measure its 540 and
591 keV gamma lines.
Compiled from: Anders, 1960; Bate, 1979; CRC, 1998-99; Choppin et al., 1995; Cobble,
1964; Considine and Considine, 1983; Coomber, 1975; Cotton and Wilkinson, 1988; DOE,
1990 and 1997, 1997; Ehmann and Vance, 1991; Foti et al., 1972a, 1972b; Fried, 1995;
Greenwood and Earnshaw, 1984; Hassinsky and Adloff, 1965; Kleinberg et al., 1960;
Lindsay, 1988; SCA, 2001; Wahl and Bonner, 1951.
14.10.9.13 Thorium
Thorium, with an atomic number of 90, is the second member in the series of actinide elements.
It is one of only three of the actinides—thorium, protactinium, and uranium—that occur in nature
in quantities sufficient for practical extraction. In solution, in all minerals, and in virtually all
compounds, thorium exists in the +4 oxidation state; it is the only actinide exclusively in the +4
state in solution.
Isotopes
There are 24 isotopes of thorium ranging inclusively from 213Th to 236Th; all are radioactive.
Thorium-232, the parent nuclide in the natural decay series, represents virtually 100 percent of
the thorium isotopes in nature, but there are a trace amounts of 227Th, 228Th, 230Th, 231Th, and 234Th
(progeny of 232Th and 235/238U). The remaining isotopes are anthropogenic. The most important
environmental contaminants are 232Th and 230Th (a member of the 238U decay series). They have
half-lives of 1.41×1010 years and 75,400 years, respectively.
Occurrence and Uses
Thorium is widely but sparsely dispersed in the Earth’s crust. At an average concentration of
approximately 10 ppm, it is over three times as abundant as uranium. In the ocean and rivers,
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however, its concentration is about one-thousandth that of uranium (about 10!8 g/L) because its
compounds are much less soluble under environmental conditions. There are six minerals whose
essential element is thorium; thorite (uranothorite) and thorianite are common examples. Several
lanthanum and zirconium minerals are also thorium-bearing minerals; examples include
monazite sand and uraninite. In each mineral, thorium is present as its oxide, thorium dioxide
(ThO2). Monazite sand is the most common commercial mineral, but thorite is also a source of
thorium.
Thorium is extracted from its minerals with hot sulfuric acid or hot concentrated alkali,
converted into thorium nitrate [Th(NO3)4] (its chief commercial compound), extracted with
organic solvents (commonly kerosene containing tributylphosphate), stripped from the organic
phase by alkali solutions, and crystallized as thorium nitrate or precipitated with oxalate. The
metal can be produced by electodeposition from the chloride or fluoride dissolved in fused alkali
halides or by thermoreduction of thorium compounds by calcium (1,000–1,200 EC). Thorium can
also be produced as a by-product in the production of other valuable metals such as nickel,
uranium, and zirconium, in addition to the lanthanides. Unextracted minerals or partially
extracted mill tailings represent some forms of thorium contaminants found in the environment.
Very insoluble forms of thorium hydroxide [Th(OH)4] are other common species found.
Metallic thorium has been used as an alloy in the magnesium industry and as a deoxidant for
molybdenum, iron, and other metals. Because of its high density, chemical reactivity, poor
mechanical properties, and relatively high cost, it is not used as a structural material. Thorium
dioxide is a highly refractory material with the highest melting point among the oxides,
3,390 EC. It has been used in the production of gas mantles, to prevent crystallization of tungsten
in filaments, as furnace linings, in nickel alloys to improve corrosion resistance, and as a catalyst
in the conversion of methanol to formaldehyde. Thorium-232 is a fuel in breeder reactors. The
radionuclide absorbs slow neutrons, and with the consecutive emission of two beta particles, it
decays to 233U, a fissionable isotope of uranium with a half-life of 159,000 years.
Solubility of Compounds
Thorium exists in solution as a highly charged ion and undergoes extensive interaction with
water and with many anions. Few of the compounds are water soluble; soluble thorium
compounds include the nitrate [Th(NO3)4], sulfate [Th(SO4)2], chloride (ThCl4), and perchlorate
[Th(ClO4)4]. Many compounds are insoluble in water and are used in the precipitation of thorium
from solution, including the hydroxide [Th(OH)4], fluoride (ThF4), iodate [Th(IO3)4], oxalate
[Th(C2O4)2], phosphate [Th3(PO4)4], sulfite [Th(SO3)2], dichromate [Th(Cr2O7)2], potassium
hexafluorothorionate [K2ThF6], thorium ferrocyonide (+2) [ThFe(CN)6], and thorium peroxide
sulfate [Th(OO)2SO4].
The thorium ion forms many complex ions, chelates, and solvated species that are soluble in
organic solvents. This property is the basis of many procedures for the separation and purification
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of thorium (see below). For example, certain ions, such as nitrate and sulfate, form large
unsolvated complex ions with thorium that are soluble in organic solvents. Chelates of 1,3diketones, such as acetylacetone (acac) and TTA, form neutral molecular chelates with the
thorium ion that are soluble. In addition, many neutral organic compounds have strong solvating
properties for thorium, bonding to the thorium ion in much the same way water solvates the ion
at low pH. TBP, diethyl ether, methyl ethyl ketone, mesityl oxide, and monoalkyl and dialkyl
phosphates are examples of such compounds.
Review of Properties
Thorium is the first member of the actinide series of elements that includes actinium (Ac),
uranium, and the transuranium elements. Thorium is a bright, silver-white metal with a density
above 11 g/cm3. It tarnishes in air, forming a dark gray oxide coating. The massive metal is
stable, but in finely divided form and as a thin ribbon it is pyrophoric and forms thorium oxide
(ThO2). Thorium metal dissolves in hydrochloric acid, is made passive by nitric acid, but is not
affected by alkali. It is attacked by hot water and steam to form the oxide coating and hydrogen,
but its reactions with water are complicated by the presence of oxygen. Thorium has four valence
electrons (6d27s2). Under laboratory conditions, chlorides, bromides, and iodides of the bi- and
trivalent state have been prepared. In aqueous solution and in most compounds, including all
those found in nature, thorium exists only in the +4 oxidation state; its compounds are colorless
in solution unless the anion provides a color. Thorium forms many inorganic compounds in acid
solution.
Solution Chemistry
Because the only oxidation state of thorium in solution is the +4 state, its chemistry is not
complicated by oxidation-reductions reactions that might produce alternate species in solution.
With the +4 charge and corresponding charge-to-radius ratio of 4.0, however, thorium forms very
stable complex ions with halides, oxygen-containing ligands, and chelating agents. Although
Th+4 is large (0.99 D; 0.099 nm; 99 pm) relative to other +4 ions (Ti, Zr, Hf, Ce) and therefore
more resistant to hydrolysis, as a highly charged ion, it hydrolyzes extensively in aqueous
solutions above pH 3 and tends to behave more like a colloid than a true solution. The
concentration of Th+4 is negligible under those conditions. Below pH 3, however, the
uncomplexed ion is stable as the hydrated ion, Th(H2O)8 or 9+4.
COMPLEXATION. Thorium has a strong tendency to form complex ions in solution. The presence
of HF forms very stable complex ions, for example, with one, two, or three ligands:
Th+4 + HF 6 ThF+3 + H+1
ThF+3 + HF 6 ThF2+2 + H+1
ThF2+2 + HF 6 ThF3+1 + H+1
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These complex ions represent the predominant species in solutions containing HF. Stable
complex ions also form with oxygen-containing ligands such as nitrate, chlorate, sulfate,
bisulfate, iodate, carbonate, phosphate, most carboxylate anions, and chelate anions. Some
chelating agents such as salicylate, acetylacetonate (acac), TTA, and cupferron form complexes
that are more soluble in organic solvents, This property is the basis of several radiochemical
isolation methods for thorium. Through the formation of soluble complex ions, chelating agents
found in some industrial wastewater or natural water samples will interfere to varying degrees
with the isolation of thorium by ferric hydroxide [Fe(OH)3] coprecipitation. Alternative isolation
methods should be used, such as coprecipitation from an acidic solution with an alternative
reagent. Protonation of the anionic form of chelates with acid renders them useless as chelating
agents. Other complexing agents also interfere with precipitation by the formation of soluble
ions. Thorium, for example, does not precipitate with oxalate in the presence of carbonate ions.
A procedure for separating thorium from rare-earth ions takes advantage of the formation of a
soluble thorium-EDTA complex that inhibits thorium precipitation when the rare-earth ions are
precipitated with phosphate. The presence of high concentrations of other complexing agents
such as phosphate, chloride, and other anions found in some samples takes thorium into a
completely exchangeable form when it is solubilized in high-concentration nitric acid.
HYDROLYSIS. Beginning at pH 3, thorium ions undergo extensive hydrolysis to form monomeric
and polymeric complexes in solution, leaving little Th+4 in a saturated solution at pH 3
(approximately 5×10!6 M). Tracer solutions containing 234Th can be added at pH 2 to allow
equilibration because it is not likely to occur if part of the thorium is hydrolyzed and bound in
polymeric forms.
The hydrolysis process is complex, depending on the pH of the solution and its ionic strength.
Several species have been proposed: three are polynuclear species, Th2(OH)2+6, Th4(OH)8+8, and
Th6(OH)15+9; and two are monomeric species, Th(OH)+3 and Th(OH)2+2. The monomeric species
are of minor importance except in extremely dilute solutions, but they become more important as
the temperature increases. The presence of chloride and nitrate ion diminishes hydrolysis,
because the formation of corresponding complex ions markedly suppresses the process. Hydrolysis increases with increasing hydroxide concentration (pH), and eventually polymerization of the
species begins. At a pH of about 5, irreversible hydrolysis produces an amorphous precipitate of
thorium hydroxide, a polymer that might contain more than 100 thorium atoms. Just before
precipitation, polymerization slows and equilibration might take weeks or months to obtain.
Routine fuming of a sample containing organic material with nitric acid is recommended after
addition of tracer, but before separation of thorium as a hydroxide precipitate because there is
evidence for lack of exchange between added tracer and isotope already in solution. Complexing
with organic substances in the initial solution or existence of thorium in solution as some
polymeric ion have been suggested as the cause.
ADSORPTION. The insoluble hydroxide that forms in solution above pH 3 has a tendency to
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coagulate with hydrated oxides such as ferric oxide. The high charge of the Th+4cation, high
charge-to-radius ratio, and tendency to hydrolyze all contribute to the ability of thorium to adsorb
on surfaces by ion-exchange mechanisms or chemical adsorption mechanisms. These adsorption
properties greatly affect the interaction of thorium with ion-exchange resins and environmental
media such as soil.
Dissolution of Samples
Thorium samples are ignited first to remove organic materials. Most compounds will decompose
when sintered with sodium peroxide (Na2O2), and most thorium minerals will yield to alternate
sodium peroxide sintering and potassium pyrosulfate (K2S2O7) fusion. It is often necessary to
recover thorium from hydrolysis products produced by these processes. The hydrolysis products
are treated with hydrofluoric acid, and thorium is recovered as the insoluble fluoride. Rock
samples are often dissolved in hydrofluoric acid containing either nitric acid or perchloric acid.
Monazite is dissolved by prolonged sintering or with fuming perchloric or sulfuric acid. Thorium
alloys are dissolved in two steps, first with aqua regia (nitric and hydrochloric acid mixture)
followed by fusion with potassium pyrosulfate. Thorium targets are dissolved in concentrated
nitric acid containing hydrofluoric acid, mantles in nitric or sulfuric acid, and tungsten filaments
with aqua regia or perchloric acid.
Separation Methods
PRECIPITATION AND COPRECIPITATION. Precipitation and coprecipitation are used to separate and
collect thorium from aqueous solutions either for further treatment in an analytical scheme or for
preparation of a sample for counting. Formation of insoluble salts is used to precipitate thorium
from solution; examples include the hydroxide, peroxide, fluoride, iodate, oxalate, and
phosphate, among others. Tracer quantities of thorium are commonly coprecipitated with
lanthanum fluoride (LaF3), neodymium fluoride (NdF3), and cerium fluoride (CeF3) in separation
schemes and to prepare samples for alpha counting. Tracer quantities are also carried with
calcium oxalate [Ca(C2O4)], ferric hydroxide [Fe(OH)3], zirconium iodate [Zr(IO3)4], zirconium
phosphate [Zr3(PO4)4], and barium sulfate (BaSO4).
ION EXCHANGE. The highly charged thorium cation is strongly adsorbed onto cation exchangers
and is more difficult to elute than most other ions. Its strong adsorption property makes it
possible to remove trace quantities of thorium from a large volume of solution onto small
amounts of ion-exchange resin. Washing the resin with mineral acids of various concentrations
separates thorium from less strongly bound cations that elute from the resin. For example, Th+4
remains bonded at all hydrochloric concentrations, allowing other cations to be eluted at different
concentrations of acid. Thorium is eluted by complexing agents such as citrate, lactate, fluoride,
carbonate, sulfate, or oxalate that reduce the net charge of the absorbing species, causing reversal
of the adsorption process.
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Anion exchangers are useful for separating thorium, but the contrasting behavior of thorium with
the resin depends on whether hydrochloric or nitric acid is used as an eluent. In hydrochloric
acid, several metal ions, unlike thorium, form negative complexes that can be readily removed
from a thorium solution by adsorption onto the anionic exchanger. Thorium forms positively
charged chlorocation complexes or neutral thorium chloride (ThCl4) in the acid and is not
exchanged onto the resin at any hydrochloric acid concentration. In contrast, thorium forms
anionic complexes in nitric acid solution that adsorb onto the exchanger over a wide range of
nitric acid concentrations, reaching a maximum affinity near 7 M nitric acid. Behavior in nitric
acid solution is the basis for a number of important radiochemical separations of thorium from
rare earths, uranium, and other elements.
ELECTRODEPOSITION. Thorium separated from other actinides by chemical methods can be
electrodeposited for alpha counting from a dilute solution of ammonium sulfate adjusted to a pH
of 2. The hydrous oxide of thorium is deposited in one hour on a highly polished platinum or
stainless-steel disc serving as the cathode of an electrolytic cell. The anode is a platinum-iridium
alloy.
SOLVENT EXTRACTION. Many complexes and some compounds of thorium can be extracted from
aqueous solutions into a variety of organic solvents. The TTA (α-theonyltrifluoroacetone)
complex of metals is widely used in radiochemistry for the separation of ions. Thorium can be
separated from most alkali metal, alkaline earth, and rare earth metals after the complex is
quantitatively extracted into benzene above pH 1. Backwashing the organic solution with dilute
acid leaves the more soluble ions in benzene.
Extraction of nitrates and chlorides of thorium into organic solvents from the respective acid
solutions is widely used for isolation and purification of the element. One of the most common
processes is the extraction of thorium nitrate from a nitric acid solution with TBP. TBP is usually
diluted with an inert solvent such as ether or xylene/toluene to reduce the viscosity of the
mixture. Dilution reduces the extraction effectiveness of the mixture, but the solubility of many
contaminating ions is greatly reduced, increasing the effectiveness of the separation when the
thorium is recovered by backwashing.
Long-chain amine salts have been very effective in carrying thorium in laboratory and industrial
extraction process using xylene/toluene. Complex sulfate anions of thorium are formed in
sulfuric acid that act as the counter ion to the protonated quaternary amine cation. They
accompany the organic salt into the organic phase.
In recent years, solvent extraction chromatography procedures have been developed to separate
thorium. These procedures use extraction chromatography resins that consist of extractant
materials such as Aliquat-336® (tricaprylylmethylammonium chloride or methyltricaprylylammonium chloride), CMPO in TBP, or DPPP (dipentylpentylphosphonate), also called DAAP
(diamylamylphosphonate), or absorbed onto an inert polymeric material. They are used in a
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column, rather than in the traditional batch mode, and provide a rapid efficient method of
separating the radionuclide with the elimination of large volumes of organic waste.
Methods of Analysis
Chemical procedures are used for the analysis of macroscopic quantities of thorium in solution
after it has been separated by precipitation, ion exchange, extraction, and/or extraction chromatography from interfering ions. Gravimetric determination generally follows precipitation as the
oxalate that is calcined to the oxide (ThO2). Numerous volumetric analyses employ EDTA as the
titrant. In the most common spectrometric method of analysis, thorin, a complex organoarsenic
acid forms a colored complex with thorium that is measured in the visible spectrum.
Trace quantities of thorium are measured by alpha spectrometry after chemical separation from
interfering radionuclides. Thorium-227, 228Th, 230Th, and 232Th are determined by the
measurement of their respective spectral peaks (energies), using 234Th as a tracer to determine the
chemical yield of the procedure. The activity of the tracer is determined by beta counting in a
proportional counter. Thorium-234 also emits gamma radiation that can be detected by gamma
spectrometry; however, the peak can not be measured accurately because of interfering peaks of
other gamma-emitting radionuclides. Thorium-229 is sometimes used as a tracer to determine the
chemical yield of the alpha spectrometric procedure, but it produces considerable recoil that
might contaminate the detector.
Compiled from: Ahrland, 1986; Baes and Mesmer, 1976; Cotton, 1991; Cotton and
Wilkinson, 1988; DOE, 1990 and 1997, 1997; EPA, 1980 and 1984; Greenwood, 1984;
Grimaldi, 1961; Hassinsky and Adloff, 1965; Hyde, 1960; Katzin, 1986; Lindsey, 1988.
14.10.9.14 Tritium
Unlike the elements reviewed in this section, tritium is the only radionuclide of the element
hydrogen. It contains two neutrons and is represented by the symbols 3H, 3T, or simply, T. The
atom contains only one valence electron so its common oxidation state, besides zero, is +1,
although it can exist in the !1 state as a metal hydride.
Occurrence and Uses
Tritium is found wherever hydrogen is found, with and without the other isotopes of the element
(hydrogen and deuterium)—as molecular hydrogen (HT, DT, T2), water (HOT, DTO, T2O), and
inorganic and organic compounds, hydrides and hydrocarbons, respectively, for example. About
99 percent of the radionuclide in nature from any source is in the form of HOT. Natural processes
account for approximately one T atom per 1018 hydrogen atoms. The source of some natural
tritium is ejection from the sun, but the primary source is from bombardment of 14N with cosmic
neutrons in the upper atmosphere:
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14
7

N + 01n → 3 H +126 C

Most tritium from this source appears as HOT.
Tritium is produced in laboratory and industrial processes by nuclear reactions such as:
2
1

D + 21D → 31T +11 H

For large-scale production of tritium, 6Li alloyed with magnesium or aluminum is the target of
neutrons:
6
3

Li + 01n → 31T +42 He

The radionuclide is retained in the alloy until released by acid dissolution of the target. Large
quantities are handled as HT or HOT. HOT is formed from HT when it is exposed to oxygen or
water vapor. A convenient way to store tritium is as the hydride of uranium (UT3). It is formed by
reacting the gas with finely divided uranium and is released by heating the compound above
400 EC.
Tritium is also produced in nuclear reactors that contain water or heavy water from the neutron
bombardment of boron, lithium, and deuterium:
10

B(n, T) 2 4He
B(n, T) 9Be
6
Li (n, T) 4He
2
H (n,γ) T
11

and from the fission process as a ternary fission fragment. Significant uses for tritium are in
fission bombs to boost their yield, in thermonuclear weapons (the hydrogen bomb), in luminescent signs, and in night-vision military applications. Tritium bombarded with high-energy
deuterons undergoes fusion to form helium and releases neutrons:
3
1

H + 21H → 42 He +01n

A tremendous amount of energy is released during the nuclear reaction, much more than the
energy of the bombarding particle. Fusion research on controlled thermonuclear reactions should
lead to an energy source for electrical generation.
Tritium absorbed on metals are a source of neutrons when bombarded with deuterons. Mixed
with zinc sulfide, it produces radioluminescence that is used in luminescent paint and on watch
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dials. Gaseous tritium in the presence of zinc sulfide produces a small, permanent light source
found in rifle sights and exit signs. Tritium is also a good tracer because it does not emit gamma
radiation. Hydrological studies with HOT are used to trace geological water and the movement of
glaciers. It is also used as a tracer for hydrogen in chemical studies and biological research. In
medicine, it is used for diagnosis and radiotreatment.
Review of Properties
Tritium (t½ . 12.3 y) decays by emission of a low-energy beta particle to form 3He, and no
gamma radiation is released. The range of the beta particle is low, 6 mm in air and 0.005 mm in
water or soft tissue.
The physical and chemical properties of tritium are somewhat different than hydrogen or
deuterium because of their mass differences (isotope effects). Tritium is approximately 1.5 times
as heavy as deuterium and three times heavier than hydrogen, and the isotope effect can be large
for mass differences of these magnitudes. In its simple molecular form, tritium exists primarily as
T2 or DT. The oxide form is HOT, DTO, or T2O, with higher molecular weights than water
(H2O). Thus molecules of tritiated water are heavier, and any process such as evaporation or
distillation that produces a phase transition results in isotopic fractionation and enrichment of
tritium in water. In a mixture of the oxides, various mixed isotopic water species are generally
also present because of exchange reactions: in any mixture of H2O, D2O, and T2O, HOT and
DTO are found.
Tritium can be introduced into organic compounds by exposing T2 to the compound for a few
days or weeks, irradiation of the compound and a lithium salt with neutrons (recoil labeling), or it
can be selectively introduced into a molecule by chemical synthesis using a molecular tritium
source such as HOT. Beta radiation causes exchange reactions between hydrogen atoms in the
compound and tritium and migration of the isotope within the molecule. Phenol (C6H5OH), for
example, labeled with tritium on the oxygen atom (C6H5OT) will become C6H4TOH and
C6H4TOT. When tritium samples are stored in containers made from organic polymers such as
polyethylene, the container will adsorb tritium, resulting in a decrease in the concentration of
tritium in the sample. Eventually, the tritium atoms will migrate to the outer surface of the
container, and tritium will be lost to the environment. Catalytic exchange also occurs in tritiated
solutions or solutions containing T2 gas. Exchange is very rapid with organic compounds when
H+1 or OH!1 ions or if a hydrogen-transfer agent such as Pt or Pd is present.
Tritium as HT or HOT will absorb on most metallic surfaces. Penetration at room temperature is
very slow, and the radionuclide remains close to the surface. In the form of HOT, it can be
removed with water, or by hydrogen gas in the form of HT. Heating aids the removal. When
tritium is absorbed at elevated temperatures, it penetrates deeper into the surface. Adsorption
under these conditions will result in enough penetration to cause structural damage to the metal,
especially if the process continues for extended periods. Hydrogenous material such as rubber
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and plastics will also absorb tritium. It will penetrate into the material, and hydrogenous
materials are readily contaminated deep into the material, and it is impossible to completely
remove the tritium. Highly contaminated metal or plastic surfaces can release some of the loosely
bound tritium immediately after exposure in a process called outgassing.
Pure T2O can be prepared by oxidation of tritium gas with hot copper oxide (Cu+2) or direct
combination of the gas with oxygen in the presence of an electrical spark. It is never used for
chemical or biological processes because one milliliter contains 2,650 curies. The liquid is selfluminescent, undergoes rapid self-radiolysis, and considerable radiation damage is done to
dissolved species. For the same reason, very few compounds of pure tritium have ever been
prepared or studied.
Tritium is not a hazard outside the body. Gamma radiation is not released by its decay. The beta
emission is low in energy compared to most beta emitters and readily stopped by the outer layer
of skin. Only ingested tritium can be a hazard. Exposure to tritium is primarily in the form of HT
gas or HOT water vapor, although T2 and T2O may be present. Only about 0.005 percent of the
activity of inhaled HT gas is incorporated into lung tissue, and most is exhaled. In addition,
tritiated water can be absorbed through the skin or wounds unless protective equipment is used.
Tritium is found in tissue wherever hydrogen is found. The biological half-life is about ten days,
but the value varies significantly, depending on exertion rates and fluid intake.
Environmental tritium is formed in the gaseous and aqueous forms, but over 99 percent of tritium
from all sources is found in the environment after exchange with hydrogen in water in the form
of HOT. It is widely distributed in the surface waters of the Earth and makes a minor contribution
to the activity of ocean water. It can also be found in laboratories and industrial sites in the form
of metal hydrides, tritiated pump oil, and tritiated gases such as methane and ammonia.
Tritium found in environmental samples may be either exchangeable in acid media (labile) or
organically bound. In the latter case, combustion of the material is necessary to release the tritium
into an exchangeable form. This is performed usually by adding an oxidizing agent, like KMnO4,
if the contribution of the organic tritium to the total tritium is large.
Separation Methods
DISTILLATION. Tritium in water samples is essentially in the form of HOT. It can be removed
quantitatively from aqueous mixtures by distillation to dryness, which also separate it from other
radionuclides. Volatile iodine radionuclides are precipitated as silver iodide before distillation, if
they are present. The aqueous solution is usually distilled, however, from a basic solution of
potassium permangenate, which will oxidize radionuclides, such as iodine and carbon, and
oxidize organic compounds that might interfere with subsequent procedures, liquid scintillation
counting, for example. Charcoal can also be added to the distillation mixture as an additional
measure to remove organic material. Contaminating tritium in soil samples can be removed by
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distillation from similar aqueous mixtures. All tritium in soil samples might not be recovered by
this method, however, if the tritium is tightly bound to the soil matrix. Tritium also can be
removed by distillation of an azeotrope mixture formed with toluene or cyclohexane. In some
procedures, tritium is initially separated by distillation and then concentrated (enriched) by
electrolysis in an acid or base solution. Recovery of tritium from the electrolytic cell for analysis
is accomplished by a subsequent distillation.
DECOMPOSITION. Organically bound tritium in vegetation, food, and tissue samples can be
removed by combustion. The sample is freeze dried (lyophilized), and the water from the process
is collected in cold traps for tritium analysis. The remaining solid is collected as a pellet, which is
burned at 700 EC in a highly purified mixture of argon and oxygen in the presence of a copper(I)
oxide (Cu2O) catalyst, generated on a copper screen at the temperature of the process. Water
from the combustion process, containing tritium from the pellet, and water from the freezedrying process is analyzed for tritium by liquid scintillation counting.
Tritium in HOT can be reduced to TH by heating with metals, such as magnesium, zinc, or
calcium, and analyzed as a gas. Conversely, if tritium is present as HT or T2, it may be oxidized
to HOT by passing the gaseous sample over a platinum, palladium, or nickel catalyst in the
presence of air.
CONVERSION TO ORGANIC COMPOUNDS. Compounds that react readily with water to produce
hydrogen derivatives can be used to isolate and recover tritium that is present in the HOT form.
Organic compounds containing magnesium (Grignard reagents) with relatively low molecularweights will react spontaneously with water and produce a gaseous product containing hydrogen
from the water. Tritium from HOT in a water sample will be included in the gaseous sample. It is
collected after formation by condensation in a cold trap and vaporized into a gas tube for
measurement. Grignard reagents formed from butane, acetylene, and methane can be used in this
method. Tritiated butane is produced by the following chemical reaction:
C4H9MgBr + THO 6 C4H9T + Mg(OH)Br
Inorganic compounds can also be use to produce gaseous products:
Al4C3 + 3 HOT + 9 H2O 6 3 CH3T + 4 Al(OH)3
EXCHANGE. Methods to assess tritium in compounds take advantage of exchange reactions to
collect the radionuclide in a volatile substance that can be collected in a gas tube for measurement. Acetone is one compound that easily exchanges tritium in an acid or base medium and is
relatively volatile.

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Methods of Analysis
Tritium is collected primarily as HOT along with water (H 2O) by distillation and then determined
from its beta emission in a liquid scintillation system. No gamma rays are emitted. The distillation process is usually performed from a basic solution of potassium permangenate to oxidize
radionuclides and organic compounds, preventing them from distilling over and subsequently
interfering with counting. Charcoal can also be added to the distillation mixture as an additional
measure to remove organic material. Volatile iodine radionuclides can be precipitated as silver
iodide before distillation. Another distillation technique involves the use of cyclohexane to form
an azeotropic (low boiling point) mixture. This technique is sometimes used in analysis of biota
samples. Tritium may be analyzed, indirectly, by mass spectrometry of its progeny, 3He.
Compiled from: Choppin et al., 1995; Cotton and Wilkinson, 1988; DOE, 1994; Demange et
al., 2002; Duckworth, 1995; Greenwood and Earnshaw, 1984; Hampel, 1968; Hassinky and
Adloff, 1965; Kaplan, 1995; Lindsay, 1988; Mitchell, 1961; Passo and Cook, 1994; Surano et
al., 1992.
14.10.9.15 Uranium
Uranium, atomic number 92, is the last naturally occurring member of the actinide series and the
precursor to the transuranic elements. Three isotopes are found in nature, and uranium was the
active constituent in the salts whose study led to the discovery of radioactivity by Becquerel in
1896.
Isotopes
There are 19 isotopes of uranium with mass numbers ranging from 222 to 242. All isotopes are
radioactive with half-lives range ranging from microseconds to billions of years. Uranium-235
(0.72%) and 238U (99.27%) occur naturally as primordial uranium. Uranium-234 has a natural
abundance of 0.0055%, but is present as a part of the 238U decay natural decay chain. The 234U
that was formed at the time the Earth was formed has long since decayed. The half-lives of these
principal isotopes of uranium are listed below.
Isotope
234
235
238

Alpha Decay
Half-Life
2.46 × 105 years
7.04 × 108 years
4.48 × 109 years

Spontaneous Fission
Half-Life
1.42 × 1016 years
9.80 × 1018 years
8.08 × 1015 years

These isotopes have two different decay modes. Each decay mode has its own characteristic halflife. As seen above the alpha decay mode is the most significant, because it has the shortest halflife for each of these isotopes.
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Another isotope of uranium of significance is 232U (t½ . 69.8 y). It is used as a tracer in uranium
analyses and is also an alpha emitter so it can be determined concurrently with the major uranium
isotopes by alpha spectrometry.
Uranium-235 and artificially produced 233U are fissionable material on bombardment with slow
(thermal) neutrons. Other uranium radionuclides are fissionable with fast moving neutrons,
charged particles, high-energy photons, or mesons. Uranium-238 and 235U are both parents of
natural radioactive decay series, the uranium series of 238U that eventually decays with alpha and
beta emissions to stable 206Pb and the actinium series of 235U that decays to 207Pb.
Occurrence and Uses
Naturally occurring uranium is believed to be concentrated in the Earth’s crust with an average
concentration of approximately 4 ppm. Granite rocks contains up to 8 ppm or more, and ocean
water contains 0.0033 ppm. Many uranium minerals have been discovered. Among the better
known are uraninite, carnotite, adavidite, pitchblende, and coffinite. The latter two minerals are
important commercial sources of uranium. It is also found in phosphate rock, lignite, and
monazite sands and is commercially available from these sources. The artificial isotope, 233U, is
produced from natural 232Th by absorption of slow neutrons to form 233Th, which decays by the
emission of two beta particles to 233U.
Uranium is extracted from uranium minerals, ores, rocks, and sands by numerous chemical
extraction (leaching) processes. The extraction process is sometimes preceded by roasting the ore
to improve the processing characteristic of the material. The extraction process uses either an
acid/oxidant combination or sodium carbonate treatment, depending on the nature of the ore, to
convert the metal to a soluble form of the uranyl ion. Uranium is recovered from solution by
precipitating the uranate salt with ammonia or sodium hydroxide solution. Ammonium uranate is
known as “yellow cake.” The uranate salt is solubilized to give a uranyl nitrate solution that is
further purified by extraction into an organic phase to separate the salt from impurities and
subsequent stripping with water. It is precipitated as a highly purified nitrate salt that is used to
produce other uranium compounds—uranium trioxide (UO3) by thermal processing or uranium
dioxide (UO2) on reduction of the trioxide with hydrogen. Uranium tetrafluoride (UF4) is
prepared, in turn, from the dioxide by treatment with hydrogen fluoride. The metal is recovered
by fused-salt electrolysis in molten sodium chloride-calcium chloride or reduction with more
active metals such as calcium or magnesium (Ames Process) in an inert atmosphere at 1,000 EC.
Early in the twentieth century, the only use of uranium was in the production of a brown-yellow
tinted glass and glazes; it was a byproduct of the extraction of radium, which was used for
medicinal and research purposes. Since the mid-twentieth century, the most important use of
uranium is as a nuclear fuel, directly in the form of 233U and 235U, fissionable radionuclides, and
in the form of 238U that can be converted to fissionable 239Pu by thermal neutrons in breeder
reactors. Depleted uranium, uranium whose 235U content has been reduced to below about 0.2
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percent, the majority of waste from the uranium enrichment process, is used in shielded
containers to transport radioactive materials, inertial guidance devices, gyro compasses,
counterweights for aircraft control surfaces, ballast for missile reentry vehicles, fabrication of
armor-piercing conventional weapons, and tank armor plating. Uranium metal is used as a X-ray
target for production of high-energy X-rays, the nitrate salt as a photographic toner, and the
acetate is used in analytical chemistry.
Solubility of Compounds
Only a small number of the numerous uranium compounds are soluble in water. Except for the
fluorides, the halides of uranium (+3 and +4) are soluble, as are the chloride and bromide of
U(V) [UOX2] and the fluoride, chloride, and bromide of U(VI) [UO2X2]. Several of the uranyl
(UO2+2) salts of polyatomic anions are also soluble in water: the sulfate, bicarbonate, acetate,
thiocyanate, chromate, tungstate, and nitrate. The latter is one of the most water-soluble uranium
compounds.
Review of Properties
Uranium is a dense, malleable and ductile metal that exists in three allotropic forms: alpha, stable
to 688 EC where it forms the beta structure, which becomes the gamma structure at 776 EC. It is
a poor conductor of electricity. The metal absorbs gases and is used to absorb tritium. Uranium
metal tarnishes readily in an oxidation process when exposed to air. It burns when heated to 170
EC, and the finely divided metal is pyrophoric. Uranium slowly decomposes water at room
temperature, but rapidly at 100 EC. Under a flux of neutrons and other accelerated particles,
atoms of uranium are displaced from their equilibrium position in its metallic lattice. With high
temperatures and an accumulation of fission products, the metal deforms and swells, becoming
twisted, porous, and brittle. The problem can be avoided by using some of its alloys, particularly
alloys of molybdenum and aluminum.
Uranium forms a large number of binary and ternary alloys with most metals. It also forms
compounds with many metals: aluminum, bismuth, cadmium, cobalt, gallium, germanium, gold,
indium, iron, lead, magnesium, mercury, nickel, tin, titanium, zinc, and zirconium. Many binary
compounds of the nonmetals are also known: hydrides, borides, carbides, nitrides, silicides,
phosphides, halides, and oxides. Although other oxides are known, the common oxides are UO2,
UO3, and U3O8. Uranium reacts with acids to form the +4 salts and hydrogen. It is very reactive
as a strong reducing agent.
Uranium compounds are toxic at high concentrations. The physiological damage occurs to
internal organs, especially the kidneys. The radioactivity of natural uranium radionuclides is not
of great concern, although it is high for some artificial isotopes. Natural uranium in the
environment is considered a relatively low hazard, however, because of its very long half-life and
low toxicity at minute concentrations.
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Uranium in nature is almost entirely in the +4 and VI oxidation states. It occurs as the oxides,
UO2 and U3O8, in the solid state. In ground water under oxic conditions it exists as UO2+2 or
complexes of carbonate such as UO2(CO3)3!4. Complex formation increases its solubility under
all conditions in normal groundwater and even under fairly strong reducing conditions. The
amount associated with particulate matter is small in natural oxic waters. In some waters,
solubility may be limited, however, by formation of an uranyl silicate species. Uranium in
general is poorly absorbed on geologic media under oxic conditions, especially at moderate and
high concentrations and in the presence of high carbonate concentrations. A significant
adsorption occurs at pH above about 5 or 6 because of formation of hydrolytic complexes.
Reduction to the IV oxidation state would increase uptake in the environmental pH range.
Solution Chemistry
The radiochemistry of uranium is complicated because of the multiple oxidation states that can
exist in solution and the extensive complexation and hydrolytic reactions the ions are capable of
undergoing in solution. Four oxidation states are possible: +3, +4, (V) and (VI); the latter two
exist as oxycations: UO2+1 and UO2+2, respectively. Their stabilities vary considerably, and the +4
and +6 states are stable in solution under certain conditions; oxidation-reduction reagents are
used to form and maintain these ions in solution. Each ion has different chemical properties, and
those of the +4 and (VI) states have been particularly exploited to stabilize, solubilize, separate,
and collect uranium. The multiple possibilities of oxidation state, complexation, and hydrolysis
should be carefully considered when planning any radiochemical procedures.
OXIDATION-REDUCTION BEHAVIOR. The multiple oxidation states can be exploited during
separation procedures by taking advantage of their different chemical properties. Thorium can be
separated from uranium, for example, by oxidizing uranium in solution to the +6 oxidation state
with 30 percent hydrogen peroxide (H2O2) and precipitating thorium as the hydroxide; in the +6
state, uranium is not precipitated.
The U+3 ion is an unstable form of uranium, produced in perchlorate or chloride solutions by
reduction of UO2+2 electrochemically or with zinc amalgam. It is a powerful reducing agent, and
is oxidized to U+4 by chlorine or bromine. U+3 is slowly oxidized by water with the release of
hydrogen, and oxygen from air causes rapid oxidation. Aqueous solutions are red-brown and are
stable for several days in 1 M hydrochloric acid, especially if kept cold; rapid oxidation occurs in
more concentrated acid solutions.
The tetrapositive uranous ion, U+4, is produced by dissolving water-soluble salts of the ion in
solution, dissolving uranium metal with sulfuric or phosphoric acid, reduction of UO2+1 during its
disproportionation reaction, reduction of UO2+2 by Cr+2 or Ti+3, or oxidation of U+3. The tetrapositive ion is green in solution. The ion is stable, but slowly oxidizes by oxygen from air to the +6
state.
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The UO2+1 ion (V) is extremely unstable in solution and exist only as a transient species,
disproportionating rapidly to U+4 and UO2+2 according to the following reaction in the absence of
complicating factors (k = 1.7×106):
2 UO2+1 + 4 H+1 W UO2+2 + U+4 + 2 H2O
Maximum stability is observed in the pH range 2–4 where the reaction is considerably slower.
Solutions of UO2+1 are prepared by the dissolution of UCl5 or reduction of UO2+2 ions
electrochemically or with U+4 ions, hydrogen, or zinc amalgam.
Uranium(VI) is generally agreed to be in the form of the dioxo or uranyl ion, UO2+2. As the only
oxidation state stable in contact with air, it is very stable in solution and difficult to reduce.
Because of its exceptional stability, the uranyl ion plays a central role in the radiochemistry of
uranium. It is prepared in solution by the dissolution of certain water-soluble salts: nitrate,
halides, sulfate, acetate, and carboxylates; by dissolution of uranium(VI) compounds; and
oxidation of lower-oxidation state ions already in solution, U+4 with nitric acid for example. Its
solutions are yellow in color.
COMPLEXATION. Uranium ions form numerous complex ions, and the solution chemistry of
uranium is particularly sensitive to complexing agents present. Complex-ion chemistry is very
important, therefore, to the radiochemical separation and determination of uranium.
Complexation, for example, provides a method to prevent the removal of uranium ions or its
contaminants from solution and can influence the stability of ions in solution.
Among the oxidation states exhibited in solution, the tendency for formation of anionic
complexes is:
U+4 > UO2+2 > U+3 > UO2+1,
while the order of stability of the anionic complexes is represented by:
fluoride > nitrate > chloride > bromide > iodide > perchlorate > carbonate > oxalate > sulfate.
Numerous organic complexes form, including citrate, tartrate, and EDTA, especially with UO 2+2.
There is evidence for only a few complexes of U+3, cupferron and chloride for example. In
contrast, tetrapositive uranium, U+4, forms complexes with a wide variety of anions, and many
are stable: halides—including fluoride (up to eight ligands, UF 8!4)—chloride, and bromide;
thiocyanate; and oxygen-donors, nitrate, sulfates, phosphates, carbonate, perchlorate, and
numerous carboxylates: acetate, oxalate, tartrate, citrate, and lactate. The low charge on UO2+1
precludes the formation of very stable complexes. Fluoride (from hydrogen fluoride) is notable,
however, in its ability to displace oxygen from the ion, forming UF6!1—which inhibits
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disproportionation—and precipitating the complex ion from aqueous solution. The uranyl ion,
UO2+2, readily forms stable complexes with a large variety of inorganic and carboxylate anions
very similar to those that complex with U+4. In addition, numerous organic ligands besides
carboxylates are known that contain both oxygen and nitrogen as donor atoms. Complex-ion
formation must be considered, therefore, during precipitation procedures. Precipitation of
uranium ions is inhibited, for example, in solutions containing carbonate, tartrate, malate, citrate,
hydroxylamine, while impurities are precipitated as hydroxides, sulfides, or phosphates.
Conversely, uranium is precipitated with ammonia, while other ions are kept in solution as
complexes of EDTA.
HYDROLYSIS. Some uranium ions undergo extensive hydrolysis in aqueous solution. The
reactions can lead to formation of polymeric products, which form precipitates under certain
conditions. The tendency of the various oxidation states toward hydrolysis, a specific case of
complexation, is, therefore, in the same order as that of complex-ion formation (above).
Little data are available on the hydrolysis of U+3 ion because it is so unstable in solution.
Qualitative evidence indicates, however, that hydrolysis is about that expected for a +3 ion of its
size—a much weaker acid than most other metals ions of this charge. The U+4 ion is readily
hydrolyzed in solution, but exists as the unhydrolyzed, hydrated ion in strongly acidic solutions.
Hydrolysis begins at pH<1, starting with the U(OH)+3 species. As pH increases, several species
form progressively up to U(OH)5!1. The U(OH)+3 species predominates at high acidity and low
uranium concentrations, and the concentration of each species increases rapidly with the
temperature of the solution. In less acidic solutions and as the concentration of uranium
increases, a polymeric species forms, probably U6(OH)15+9. Hydrolytic complexes of high
molecular weight probably form subsequently, culminating in precipitation. Hydrolysis of the
UO2+1 ion has been estimated to be very low, consistent with the properties of a large, positive
ion with a single charge. Hydrolysis of UO2+2 begins at about pH 3 and is fairly complicated. In
very dilute solutions, the monomeric species, UO2(OH)+1, forms initially; but the dimerized
species, (UO2)2(OH)2+2, rapidly becomes the dominant form in solution, existing in a wide range
of uranium concentration and pH. As the pH increases, more complex polynuclear species
become prominent. The presence of complexing agents, such as chloride, nitrate, and sulfate ions
suppress hydrolysis to varying degrees.
Dissolution of Samples
Metallic uranium dissolves in nitric acid to form uranyl nitrate. Large amounts dissolve
moderately rapidly, but fine turnings or powder may react violently with nitric acid vapors or
nitrogen dioxide in the vapor. The presence of oxygen in the dissolution system tends to reduce
the oxides. The rate of dissolution of large amounts of uranium may be increased by the addition
of small amounts of sulfuric, phosphoric, or perchloric acids to the nitric acid solution. Other
common mineral acids such as sulfuric, phosphoric, perchloric, hydrochloric, and hydrobromic
acid are also used to dissolve uranium metal. Simple organic acids in hydrochloric acid dissolve
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the metal, and other solvent systems are used: sodium hydroxide and hydrogen peroxide,
bromine in ethyl acetate, and hydrogen chloride in ethyl acetate or acetone. Uranium compounds
are dissolved in numerous solvents and solvent combinations such as water, mineral acids,
organic solvents such as acetone, alcohols, and diethyl ether. Dissolution of uranium from
minerals and ores is accomplished by decomposition of the sample or leaching the uranium.
Grinding and roasting the sample facilitates recovery. Decomposition of the sample can be
accomplished with mineral acids or by fusion or a combination of the two processes. Hydrofluoric acid aids the process. The sample can be fused with sodium carbonate, sodium hydroxide,
sodium peroxide, sodium bisulfate, ammonium sulfate, lithium metaborate, and magnesium
oxide. The fused sample is dissolved in water or acid. Acid and alkaline mixtures are used to
leach uranium from minerals and ores. The procedures employ common mineral acids or alkaline
carbonates, hydroxides, and peroxides. Liquid biological samples may also be extracted to
remove uranium, or the solid sample can be ashed by a wet or dry process and dissolved in acid
solution. Wet ashing is carried out with nitric acid and completed with perchloric acid, but
extreme caution should be used when using perchloric acid in the presence of organic material.
Such mixtures have been known to detonate if the perchloric acid is allowed to dry out.
Separation Methods
PRECIPITATION AND COPRECIPITATION. There are a large number of reagents that will precipitate
uranium over a wide pH range. The number of reagents available coupled with the two possible
oxidation states of uranium in solution and the complexing properties of the ions provide many
opportunities to separate uranium from other cations and the two oxidation states from each
other. Precipitation can be inhibited, for example, by the presence of complexing agents that
form soluble complexes. Complexes that form weak complexes with uranium and strong
complexes with other cations allow the separation of uranium by its precipitation while the
complexed cations remain in solution. EDTA has been used in this manner to separate uranium
from many of the transition metals and alkaline earths. In contrast, uranium forms a very strong
soluble complex with carbonate, and this property has been used to keep uranium in solution
while ammonium hydroxide precipitates iron, titanium, zirconium, and aluminum. In a similar
manner, uranium is separated from other cations as they are precipitated as sulfides or phosphates. Common precipitating reagents include:
• Ammonium hydroxide, which precipitates uranium quantitatively at pH $ 4;
• Carbonate [however, it will form soluble anionic complexes with U(VI) at pH 5 to 11 while
many other metals form insoluble hydroxides];
• Peroxide;
• Oxalic acid, which completely precipitates uranium (+4) while U(VI) forms a soluble
complex;
• Iodide;
• Iodate;
• Phosphate for U(VI) over a wide pH range;
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• Sulfate;
• Cupferron, which precipitates uranium (+4) from an acidic solution but U(VI) from a neutral
solution; and
• 8-hydroxyquinoline, which forms a quantitatively precipitate with U(VI) only.
Coprecipitation of uranium is accomplished with several carriers. In the absence of carbonate, it
is quantitatively coprecipitated with ferric hydroxide at pH from 5 to 8. Aluminum and calcium
hydroxide are also employed to coprecipitate uranium. Uranium(VI), however, is only partially
carried by metal hydroxides in the presence of carbonate, and the amount carried decreases as the
concentration of carbonate increases. Small amounts of U(VI) coprecipatate with ceric and
thorium fluoride, calcium, zirconium, and aluminum phosphate, barium carbonate, thorium
hexametaphosphate, magnesium oxide, and thorium peroxide. Uranium (+4) is carried on ceric
sulfate, the phosphates of zirconium, bismuth, and thorium, lanthanum and neodymium fluoride,
ceric and zirconium iodates, barium sulfate, zirconium phosphate, and bismuth arsenate.
SOLVENT EXTRACTION. Liquid-liquid extraction is the most common method for the separation
of uranium in radioanalytical procedures. Extraction provides a high-recovery, one-batch process
that is more reproducible than other methods. With the development of extraction chromatography, solvent extraction has become a very efficient process for uranium separation. Many and
varied procedures are used to extract uranium from aqueous solutions, but the conditions can be
summarized as: (1) composition of the aqueous phase (form of uranium, type of acid present, and
presence of common cations and anions and of foreign anions); (2) nature of organic phase (type
and concentration of solvent and diluent); (3) temperature; and (4) time of equilibrium.
Extraction processes can be conveniently divided into three systems: those based on (1) oxygen
bonding, (2) chelate formation, and (3) extraction of anionic complexes.
Oxygen-bonding systems are more specific than those based on chelate formation. They employ
organic acids, ethers, ketones, esters, alcohols, organophosphates (phosphoesters), and nitroalkanes. Ethers are effective for the extraction of uranyl nitrate from nitric acid solutions. Cyclic
ethers are especially effective, and salting agents such as calcium nitrate increase the effectiveness. Methyl isobutyl ketone (MIBK or hexone) also effectively extracts uranium as the nitrate
complex. It has been used extensively by industry in the Redox process for extracting uranium
and plutonium from nuclear fuels. Aluminum hydroxy nitrate [AlOH(NO3)2] is an excellent
salting agent for the process and the extraction efficiency is increased by the presence of the
tetrapropylammonium cation [(C3H7)4N+1]. Another common system, used extensively in the
laboratory and in industrial process to extract uranium and plutonium from fission products,
known as the PUREX process, is used in most fuel reprocessing plants to separate the radionuclides. It employs TBP, tri-n-butyl phosphate [(C4H9)3PO], in a hydrocarbon solvent, commonly
xylene/toluene, as the extractant. The uranium fuel is dissolved in nitric acid, and uranium and
plutonium are extracted into a 30 percent TBP solution, forming a neutral complex, UO 2(TBP)2.
The organic phase is scrubbed with nitric acid solution to remove impurities, plutonium is
removed by back-extracting it as Pu+3 with a nitric acid solution containing a reducing agent, and
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uranium is removed with dilute nitric acid. A complexing agent can also be used as a stripping
agent. Trioctylphosphine oxide is 100,000 times more efficient in extracting U(VI). In both
cases, nitric acid is used both to form the uranium extracting species, uranyl nitrate, and as the
salting agent. Salting with aluminum nitrate produces a higher extraction efficiency but less
specificity for uranium. Specificity depends upon the salt used, its concentration, and the diluent
concentration.
Uranium is also extracted with select chelate forming agents. One of the most common systems
used for uranium is cupferron in diethyl ether or chloroform. Uranium(VI) is not extracted from
acidic media, so impurities soluble in the mixture under acidic conditions can be extracted first.
Uranium(VI) can be reduced to U+4 for subsequent extraction. Other chelating agents used to
extract uranium include 8-hydroxyquinoline, acetylacetone in hexone, or chloroform.
Amines with molecular weights in the 250 to 500 range are used to extract anionic complexes of
U(VI) from acidic solutions. The amine forms a salt in the acidic medium consisting of an
ammonium cation and complex anion, (C10H21)3NH+1 UO2(NO3)!1, for example. Selectivity of the
amines for U(VI) is in the order: tertiary > secondary > primary. An anionic extracting system
used extensively in laboratories and industry consists of triisooctyl amine (TIOA) in xylene/
toluene. Uranium is stripped with sodium sulfate or sodium carbonate solution. A number of
mineral and organic acids have been used with the system: hydrochloric, sulfuric, nitric,
phosphoric, hydrofluoric, acetic oxalic, formic, and maleic acid. Stripping is accomplished with
dilute acid solutions.
Extraction chromatography is a simple and relatively quick method for the separation of uranium
on a highly selective, efficient column system. One separation column consists of a triamylphosphate [(C5H11O)3PO] and diamylamylphosphonate (DAAP) [C5H11O)2(C5H11)PO] mixture in
an apolar polymeric matrix. In nitric acid, uranyl nitrate forms a complex with DAAP that is
soluble in triamylphosphate. Uranium can be separated in this system from many other metal
ions including thorium and the transuranium ions, plutonium, americium, and neptunium. It is
eluted from the column with the addition of oxalate to the eluent. Another extraction chromatography column uses CMPO dissolved in TBP and fixed on the resin matrix for isolation of
uranium in nitric acid. Elution occurs with the addition of oxalic acid to the eluent.
ION-EXCHANGE CHROMATOGRAPHY. Both cation- and anion-exchange chromatography have
been used to separate uranium from other metal ions. Both stable forms of uranium, uranium +4
and VI are exchanged onto cation-exchange resins. Uranium (+4) is more strongly exchanged,
and separation of U(VI) (UO2+2) is limited. On some cation-exchange columns, the ion also tends
to tail into other ion fractions during elution. Exchange increases with temperature, however, and
increasing the pH also increases exchange up to the beginning of formation of hydrolytic
precipitates at pH 3.8. In strong acid solutions, U(VI) is weakly absorbed compared to uranium
(+3 and +4) cations. Using complexing agents can increase specificity by elution of U(VI) with
common complex-forming anions, such as chloride, fluoride, nitrate, carbonate, and sulfate.
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Specificity also may be enhanced by forming EDTA, oxalate, acetate, or sulfate complexes with
cations in the analyte, producing a more pronounced difference in absorption of the ions on the
exchange resin. A general procedure for separating U(VI) from other metals using the first
method is to absorb U(VI) at pH of 1.5 to 2 and elute the metal with acetate solution.
Anion-exchange chromatography of uranium takes advantage of the stable anionic complexes
formed by the various oxidation states of uranium, especially U(VI), with many common anions.
Uranium(VI) forms both anionic or neutral complexes with acetate, chloride, fluoride, carbonate,
nitrate, sulfate, and phosphate. Strong anion-exchange resins are more selective and have a
greater capacity than weak exchangers whose use is more limited. Factors that affect the
separations include uranium oxidation state and concentration; type of anion and concentration;
presence and concentration of other metallic ions and foreign ions; temperature, resin, size,
porosity, and cross-linking. The various oxidation states of uranium and other metal ions
(particularly the actinides), the effect of pH on formation of complexes, and the net charge of the
column are all variables controlling the separation process.
A number of chromatographic systems are available for uranium separation on anion-exchange
resins. In hydrochloric acid, uranium is often exchanged and other cations are not. Uranium(VI)
can be exchanged from concentrated hydrochloric acid while alkali metals, alkaline earths, rare
earths, aluminum, yttrium, actinium, and thorium are washed off the column. In contrast,
uranium, molybdenum, bismuth, tin, technetium, polonium, plutonium and many transition
metals are exchanged on the column, and uranium is eluted exclusively with dilute hydrochloric
acid. Various oxidation states provide another method of separation. U+4 is separated from Pr+4
and Th+4 with 8 M hydrochloric acid. Thorium, plutonium, zirconium, neptunium, and uranium
can be separated individually by exchanging all the ions except thorium from concentrated
hydrochloric acid. Plutonium (+3) elutes with concentrated acid, zirconium at 7.5 M, Np+4 with 6
M hydrochloric acid and 5 percent hydroxylamine hydrochloride, and uranium at 0.1 M acid. U+4
can be separated from U(VI) because both strongly exchangefrom concentrated hydrochloric
acid, but they separate at 6 M acid because U+4 is not exchanged at that concentration.
Uranium(VI) exchanges strongly on an anion-exchange resin in dilute hydrofluoric acid, and the
exchange decreases with increasing acid concentration. Nitric acid provides an excellent method
to purify uranium, because uranium is more strongly exchanged from nitric acid/nitrate solutions
than from chloride/HCl solutions. More selectivity is achieved when acid concentration is low
and nitrate concentrations are high. Exchange is greatest when aluminum nitrate is use as the
source of nitrate. Ethyl alcohol increases exchange significantly.
ELECTRODEPOSITION. Electrochemical procedures have been used to separate metal ions from
uranium in solution by depositing them on a mercury cathode from a sulfuric acid solution, using
5 amps for one hour. Uranium is deposited at a cathode from acetate, carbonate, oxalate, formate,
phosphate, fluoride, and chloride solutions to produce a thin, uniform film for alpha and fission
counting. This is the primary use of electrodeposition of uranium in analytical work. In another
procedure, U(VI) is electroplated on a platinum electrode from the basic solution adjacent to the
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cathode that exists in a slightly acidic bulk solution. The conditions of the process should be
carefully controlled to obtain high yields and adherent coatings on the electrode.
VOLATILIZATION. Several halides of uranium and the uranyl ion are volatile and have the
potential for separation by sublimation or fractional distillation. Practically, however, their
volatility is not used to separate uranium in analytical procedures because of technical problems
or the high temperatures that are required for some procedures, but volatilization has been used
in industrial processes. Uranium hexafluoride and uranyl hexafluoride are volatile, and the
property is used to separate 235U from 238U in natural uranium isotope mixtures. Uranium tetrachloride and hexachloride are also volatile, and uranium has been isolated from phosphate rock
by heating with a mixture of chlorine and carbon monoxide at 800 EC and collecting the
tetrachloride.
Methods of Analysis
Uranium may be determined by fluorimetry. During the separation and purification process, the
sample is fused at 625 EC in a flux mixture containing potassium carbonate, sodium carbonate,
and sodium fluoride. The residue is exposed to light and its fluorescence is measured. Another
technique related to fluorescence is kinetic phosphorimetry analysis (KPA). Aqueous solutions of
the fully digested sample are exposed to a laser at a specific wavelength, and the
phosphorescence (at a different wavelength) intensity is measured.
Total uranium may be determined by gross alpha analysis. Individual radionuclides of uranium,
U, 235U, and 238U, can be determined by their alpha-particle emissions. Mass spectrometry also
can be used for longer-lived isotopes of uranium. Uranium radionuclides are collected by
evaporating the sample to dryness on a stainless steel planchet, by microprecipitation with a
carrier, such as lanthanum or cerium fluoride, or electrodeposition on a platinum or stainlesssteel disc. In each of these techniques, care must be taken to ensure that a single oxidation state is
achieved for the uranium prior to the collection technique. Total alpha activity is determined with
a gas-flow proportional counter or an alpha liquid scintillation system. Individual radionuclides
are measured by alpha spectrometry. Alpha emissions from 232U are used as a tracer to determine
chemical recovery.
234

Neutron activation analysis (NAA) was employed to determine uranium in the hydrogeochemical
samples from Savannah River Plants within the scope of the National Uranium Resource
Evaluation Program sponsored by DOE. Uranium was determined by cyclic activation and
delayed neutron counting of the 235U fission products. The method relied on absolute activation
techniques using the Savannah River Reactor Activation Facility. NAA, followed by delayedneutron detection, was commonly used to determine 235U.
Compiled from: Alfassi, 1990; Allard et al., 1984; Ahrland, 1986; Baes and Mesmer, 1976;
ASTM D5174; Bard, 1985; Booman and Rein, 1962; Choppin et al., 1995; Considine and
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Considine, 1983; Cotton and Wilkinson, 1988; CRC, 1998-99; DOE, 1990 and 1997; Echo
and Turk, 1957; EPA, 1973; Ehmann and Vance, 1991; Fritz and Weigel, 1995; Greenwood
and Earnshaw, 1984; Grindler, 1962; Hampel, 1968; Hassinsky and Adloff, 1965; Hochel,
1979; Katz et al., 1986; Katzin, 1986; SCA, 2001; Weigel, 1986.
14.10.9.16 Zirconium
Zirconium, atomic number 40, is a member of the second-row transition elements. It exhibits
oxidation states of +2, +3, and +4, and the +4 state is common in both the solid state and in
solution. It is immediately above hafnium in the periodic table, and both elements have very
similar chemical properties—more so than any other two elements in the periodic table. It is very
difficult, but not impossible, to prepare a sample of zirconium without the presence of hafnium.
Isotopes
There are twenty-nine isotopes of zirconium, including five metastable states, with mass numbers
from 81 through 104. Five are naturally occurring, 90Zr, 91Zr, 92Zr, 94Zr, and 96Zr. The remaining
isotopes have a half-life of milliseconds to days. The lower mass number isotopes decay
primarily by electron capture and the upper mass number isotopes are beta emitters. Zirconium95 (t1/2 . 64.0 d) and 97Zr (t1/2 . 16.9 h) are fission products and are beta emitters. Zirconium-93
(t1/2 . 1.53×106y) is a rare fission product, and 98Zr, and 99Zr are short-lived products with halflives of 30.7 s and 2.1 s, respectively. All are beta emitters.
Occurrence and Uses
Zirconium is one of the most abundant and widely distributed metals found in the Earth’s crust. It
is so reactive that it is found only in the combined state, principally in two minerals, zircon,
zircon orthosilicate (ZrSiO4), and baddeleyite, mostly zirconium dioxide (ZrO2). Zirkite is a
commercial ore that consists of both minerals. Hafnium is a minor constituent of all zirconium
minerals.
In the production of zirconium metal, zirconium sands, primarily zirconium dioxide, is passed
through an electrostatic separator to remove titanium minerals, a magnetic separator to remove
iron, ileminite, and garnet, and a gravity separator to remove the less dense silica. The recovered
zircon is heated with carbon in an arc furnace to form zirconium cyanonitride, an interstitial
solution of carbon, nitrogen, and oxygen (mostly carbon) in the metal. Silicon evaporates as
silicon monoxide (SiO), becoming silicon dioxide (SiO2) at the mouth of the furnace. The hot
zirconium cyanonitride is treated with chlorine forming volatile zirconium tetrachloride (ZrCl4),
which is purified by sublimation to remove, among other impurities, contaminating oxides. The
chloride is reduced in the Kroll process, along with liquid magnesium under conditions that
produce a metal sponge. The byproduct, magnesium chloride (MgCl2), is then removed by
melting the chloride, draining it off, and removing its residues by vacuum distillation. The
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zirconium sponge is crushed, melted into bars, arc-melted in an inert atmosphere, and formed
into ingots. For additional purification, the van Arkel-de Boer process removes all nitrogen and
oxygen. Crude zirconium is heated to 200 EC in an evacuated container containing a small
amount of iodine to form volatile zirconium tetraiodide (ZrI4). A tungsten filament is electrically
heated to 1,300 EC, decomposing the iodide and depositing zirconium on the filament. The
commercial grade of zirconium still contains up to three percent hafnium. To be used in nuclear
reactors, however, hafnium should be removed. Separation is usually accomplished by solvent
extraction of zirconium from an aqueous solution of zirconium tetrachloride as a complex ion
(phosphine oxide, for example), by ion-exchange, fractional crystallization of complex fluoride
salts, distillation of complexes of zirconium tetrachloride with phosphorus pentachloride or
phosphorus oxychloride, or differential reduction of the mixed tetrachlorides (zirconium
tetrachloride is more easily reduced to the nonvolatile trichloride than hafnium tetrachloride.
Zirconium-95 and 97Zr are fission products and are also produced by bombardment of naturally
occurring 94Zr and 96Zr, respectively, with thermal neutrons. Stable 90Zr is a product of the 90Sr
decay chain:
90
38 Sr

→

90
90
39Y + β → 40 Zr

+β

Zirconium metal and its alloys are highly resistant to corrosion and withstand streams of heated
water under high pressure. These properties, along with their low cross section for thermal
neutrons, make them an important material for cladding uranium fuel elements and as core armor
material in nuclear reactors. It is also used for making corrosive resistant chemical equipment
and surgical instruments and making superconducting magnets. Zirconium compounds are also
used in the ceramics industry as refractories, glazes, and enamels, in cores for foundry molds,
abrasive grits, and components of electrical ceramics. Crystals of zircon are cut and polished to
use in jewelry as simulated diamonds. They are also used in pyrotechnics, lamp filaments, in arc
lamps, cross-linking agents for polymers, components of catalysts, as bonding agents between
metal and ceramics and between ceramics and ceramics, as tanning agents, ion exchangers, and
in pharmaceutical agents as deodorants and antidotes for poison ivy. Zirconium-95 is used to
follow homogenization of oil products.
Solubility of Compounds
The solution properties of zirconium in water are very complex, mainly because of the formation
of colloids and the extensive hydrolysis and polymerization of the zirconium ion. hydrolysis and
polymerization are strongly dependent on the pH of the solution, concentration of the ion, and
temperature. The nitrate, chloride, bromide, iodide, perchlorate, and sulfate of zirconium are
soluble in acid solution, however.

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Review of Properties
Pure zirconium is a grey-white (silvery) lustrous metal with a density of 6.49 g/cm3. It exists in
two allotropic forms, alpha and beta, with a transition temperature of 870 EC. The alpha form is
stabilized by the common impurity oxygen. The amorphous powder is blue-black. Trace amounts
of common impurities (#1 percent), such as oxygen, nitrogen, and carbon, make the metal brittle
and difficult to fabricate. The metal is not considered to be a good conductor of heat and electricity, but compared to other metals it is soft, malleable, and ductile. Zirconium forms alloys with
most metals except mercury, the alkali metals, and the alkaline earths. It can absorb up to ten
percent oxygen and nitrogen. Zirconium is a superconductor at temperatures near absolute zero,
but its superconducting properties improve when the metal is alloyed with niobium and zinc.
Finely divided, dry zirconium (powder and chips) is pyrophoric and extremely hazardous. It is
hard to handle and store and should be moistened for safe use. Note, however, that both wetted
sponge and wet and dry stored scrap have been reported to spontaneously explode. Caution
should also be observed with waste chips produced from machining and cleaning (new)
zirconium surfaces. Both can be pyrophoric. In contrast, zirconium in the bulk form is extremely
resistant to corrosion at room temperature and remains bright and shiny in air. Resistance is
rendered by the formation of a dense, adherent, self-sealing oxide coating. The metal in this form
is resistant to acids, alkalis, and seawater. Without the coating, zirconium dissolves in warm
hydrochloric and sulfuric acids slowly; dissolution is more rapid in the presence of fluoride ions.
The metal is also resistant to high-pressure water streams and high-temperature steam. It also has
a low cross-section to thermal neutrons and is resistant to damage from neutron radiation. These
properties give pure zirconium (without hafnium) very useful as a fabrication material for nuclear
reactors. Zirconium metal alone, however, is not sufficiently resistant to hot water and steam to
meet the needs for use in a nuclear reactor. Alloyed with small percentages of tin, iron, nickel, or
chromium (Zircalloy), however, the metal meets the standards.
The coated metal becomes reactive when heated at high temperature ($ 500 EC) with nonmetals,
including hydrogen, oxygen, nitrogen, carbon, and the halogens, and forms solid solutions or
compounds with many metals. It reacts slowly with hot concentrated sulfuric and hydrochloric
acids, boiling phosphoric acid, and aqua regia. It is also attacked by fused potassium nitrate and
potassium hydroxide, but is nonreactive with aqueous alkali solutions. It is not reactive with
nitric acid. Hydrofluoric acid is the only reagent that reacts vigorously with zirconium.
Zirconium and its compounds are considered to have a low order of toxicity. Most handling and
testing indicate no level of toxicity, but some individuals seem to be allergic to zirconium
compounds. Inhalation of zirconium compound sprays and metallic zirconium dust have
produced inflammatory affects.
Very small quantities of 95Zr have been released to the environment from fuel reprocessing
facilities, atmospheric testing, and the Chernobyl accident. With a half-life of 64 days, the
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contamination of the environment is not significant. Zirconium lost from a waste repository
would be expected to move very slowly because of radiocolloidal attraction to surrounding soil
particles. Hydrolysis and polymerization renders most zirconium insoluble in natural water, but
absorption to suspended particles is expected to provide some mobility in an aqueous
environment.
Solution Chemistry
The only important oxidation state of zirconium ions in aqueous solution is +4. The solution
chemistry of zirconium is quite complex, nevertheless, because of the easy formation of colloids
and extensive hydrolysis and polymerization reactions that are strongly dependent on pH and ion
concentration.
COMPLEXATION. Zirconium ions form complexes with numerous substances: fluoride, carbonate,
borate, oxalate, and other dicarboxylic acids, among others. As a large, highly charged, spherical
ion, it exhibits high coordination numbers. One of the important chemical properties of zirconium ions in solution is the formation of a very stable hexafluorozirconate complex, ZrF!26 . For that
reason, hydrofluoric acid (HF) is an excellent solvent for the metal and insoluble zirconium
compounds. Unfortunately, the fluorocomplex interferes with most separation and determination
steps, and zirconium should be expelled by fuming with sulfuric or perchloric acid before
proceeding with analyses of other radionuclides. The addition of several milliliters of concentrated HF to a cool solution of zirconium carrier and sample will produce initial equilibration;
essentially all the zirconium is present in the +4 oxidation state as a fluoride complex. Note that
addition of HF to solutions above the azeotropic boiling point of the acid (120 EC) serves no
useful purpose and simply evaporates the HF.
Tartrate and citrate ions form stable complexes even in alkaline solutions, and zirconium
hydroxide will not precipitate in their presence (see hydrolysis below). Oxalate forms a complex
that is less stable. The ion, [Zr(C2O4)3]!2, is only stable in acid solution. On addition of base, the
complex is destroyed, and zirconium hydroxide precipitates. Sulfuric acid complexes in strongly
acidic solutions, forming Zr(SO4)!23 . In concentrated HCl solutions, ZrCl!26 is present.
Zirconium ions form chelate complexes with many organic compounds, usually through oxygen
atoms in the compounds. Typical examples are: acetylacetone (acac), EDTA, TTA, salicylic acid,
mandelic acid, cupferron, and 8-hydroxyquinoline.
HYDROLYSIS. Although Zr+4 has a large radius and any +4 cation is extensively hydrolyzed, Zr+4
appears to exist at low ion concentrations (approximately 10!4 M) and high pH. As the Zr+4
concentration increases and the concentration of H+1 decreases, however, hydrolysis and
polymerization occurs, and one or more polymeric species dominates in solution. Amorphous
hydrous oxides are precipitated near pH 2; they are soluble at high pH. Because of hydrolysis,
soluble salts (nitrate, sulfate, perchlorate, acetate, and halides) form acidic solutions when they
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dissolve. The reaction essentially seems to be a direct conversion to the tetranuclear
+8
ion. There is no convincing evidence for the existence of ZrO+2, thought at one
Zr4(OH)8(H2O)16
time to be present in equilibrium with numerous other hydrolysis products. It should be noted,
however, that freshly prepared solutions of zirconium salts might react differently from a solution
left standing for several days. Whatever the actual species in solution at any given time, the
behavior of Zr+4 depends on the pH of the solution, temperature, anion present, and age of
solution. In addition, zirconium compounds formed by precipitation from solution usually do not
have a constant composition because of their ease of hydrolysis. Even under exacting conditions,
it is difficult to obtain zirconium compounds of known, theoretical composition, and on aging,
hydrolysis products becomes more polymeric and polydisperse.
In acidic solutions, trace amounts of zirconium are strongly coprecipitated with most precipitates
in the absence of complexing ions, especially F!1 and C2O4!2 that form soluble complex ions.
In alkaline solutions, produced by the addition of hydroxide ions or ammonia, a white gelatinous
precipitate of zirconium hydroxide forms. Because the hydroxide is not amphoteric, it does not
dissolve in excess base. The precipitate is not a true hydroxide but a hydrated oxide, ZrO2 · nH2O
where n represents the variable nature of the water content. Freshly prepared zirconium hydroxide is soluble in acid; but as it dries, its solubility decreases. Precipitation is inhibited by tartrate
or citrate ions because Zr+4 forms complexes with these organic anions even in alkaline solutions
(see “Complexation,” on page 14-194, above).
In preparing zirconium solutions, it is wise to acidify the solution with the corresponding acid to
reduce hydrolysis and avoid precipitation of basic salts. During solubilization and radiochemical
equilibrium with a carrier, the tendency of zirconium ions to hydrolyze and polymerize even at
low pH should be kept in mind. Often, the formation of a strong complex with fluoride or TTA is
necessary.
RADIOCOLLOIDS. Radiocolloids of zirconium are adsorbed on practically any foreign matter (e.g.,
dirt, glass, etc.). Their formation can cause problems with dissolution, achieving radiochemical
equilibrium, and analysis. Generally, it is necessary to form a strong complex with fluoride (see
caution above) or TTA.
Dissolution of Samples
Metallic zirconium is dissolved in hydrofluoric acid, hot aqua regia, or hot concentrated sulfuric
acid. Hydrofluoric acid should be removed by fuming with sulfuric acid or perchloric acid
(caution), because fluoride interferes with most separation and analytical procedures. Zirconium
ores, rocks, and minerals are fused at high temperatures with sodium carbonate, potassium
thiosulfate, sodium peroxide, sodium tetraborate, or potassium hydrogen fluoride. The residue is
dissolved in dilute acid or water and might require filtration to collect a residue of zirconia
(impure ZrO2), which is dissolved in acid. As a minor constituent of natural sample or as a result
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M