r/EPA
United States Office of Air and Radiation EPA 402-R-04-002C
Environmental Protection Agency July 2004
UNDERSTANDING VARIATION IN
PARTITION COEFFICIENT, Kd, VALUES
Volume III:
Review of Geochemistry and Available Kd Values
for Americium, Arsenic, Curium, Iodine,
Neptunium, Radium, and Technetium
Case I: Kd = 1 mi/g
Continuous Source of Contamination
Case II: Kd =iOmi/g
Continuous Source of Contamination
Steady State
Flow
Steady State
Flow
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UNDERSTANDING THE VARIATION IN
PARTITION COEFFICIENT, Kd, VALUES
Volume III:
Review of Geochemistry and Available Kd Values
for Americium, Arsenic, Curium, Iodine,
Neptunium, Radium, and Technetium
July 2004
Interagency Agreement No. DW89937220-01-7
Project Officer
Ronald G. Wilhelm
Office of Radiation and Indoor Air
U.S. Environmental Protection Agency
Washington, D.C. 20460
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NOTICE
The following report is intended solely as guidance to EPA and other environmental professionals.
This document does not constitute rulemaking by the Agency, and cannot be relied on to create a
substantive or procedural right enforceable by any party in litigation with the United States. EPA may
take action that is at variance with the information, policies, and procedures in this document and may
change them at any time without public notice.
Reference herein to any specific commercial products, process, or service by trade name, trademark,
manufacturer, or otherwise, does not necessarily constitute or imply its endorsement, recommendation,
or favoring by the United States Government.
11
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FOREWORD
Understanding the long-term behavior of contaminants in the subsurface is becoming increasingly
important as the nation addresses groundwater contamination. Groundwater contamination is a
national concern as about 50 percent of the United States population receives its drinking water from
groundwater. It is the goal of the Environmental Protection Agency (EPA) to prevent adverse effects
to human health and the environment and to protect the environmental integrity of the nation's
groundwater.
Once groundwater is contaminated, it is important to understand how the contaminant moves in
the subsurface environment. Proper understanding of the contaminant fate and transport is necessary
in order to characterize the risks associated with the contamination and to develop, when necessary,
emergency or remedial action plans. The parameter known as the partition (or distribution) coefficient
(Kd) is one of the most important parameters used in estimating the migration potential of
contaminants present in aqueous solutions in contact with surface, subsurface and suspended solids.
This is the third volume in the series that describes: (1) the conceptualization, measurement, and
use of the partition coefficient parameter; and (2) the geochemical aqueous solution and sorbent
properties that are most important in controlling adsorption/retardation behavior of selected
contaminants. Volumes I and II were published in 1999. Volume I of this document focuses on
providing EPA and other environmental remediation professionals with a reasoned and documented
discussion of the major issues related to the selection and measurement of the partition coefficient for a
select group of contaminants. The selected contaminants investigated in Volume II of this document
include: chromium, cadmium, cesium, lead, plutonium, radon, strontium, thorium, tritium (3H), and
uranium. The contaminants discussed in Volume III include: americium, arsenic, curium, iodine,
neptunium, radium, and technetium. This three-volume report also addresses a void that has existed on
this subject in both this Agency and in the user community.
It is important to note that soil scientists and geochemists knowledgeable of sorption processes in
natural environments have long known that generic or default partition coefficient values found in the
literature can result in significant errors when used to predict the impacts of contaminant migration or
site-remediation options. Accordingly, one of the major recommendations of this report is that for
site-specific calculations, partition coefficient values measured at site-specific conditions are absolutely
essential.
For those cases when the partition coefficient parameter is not or cannot be measured, Volumes II
and III of this document: (1) provide a "thumb-nail sketch" of the key geochemical processes affecting
the sorption of the selected contaminants; (2) provide references to related key experimental and review
articles for further reading; (3) identify the important aqueous- and solid-phase parameters controlling
the sorption of these contaminants in the subsurface environment under oxidizing conditions; and (4)
identify, when possible, minimum and maximum conservative partition coefficient values for each
contaminant as a function of the key geochemical processes affecting their sorption.
in
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In addition, this publication is produced as part of ORIA's long-term strategic plan to assist in the
remediation of contaminated sites. It is published and made available to assist all environmental
remediation professionals in the cleanup of groundwater sources all over the United States.
Elizabeth A. Cotsworth, Director
Office of Radiation and Indoor Air
IV
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ACKNOWLEDGMENTS
Ronald G. Wilhelm from ORIA's Center for Radiation Site Cleanup was the project lead and EPA
Project Officer for this three-volume report. Project support was provided by EPA's Office of
Superfund Remediation and Technology Innovation (OSRTI).
EPA/ORIA wishes to thank the following people for their assistance and technical review
comments on various drafts of this report:
Paul M. Bertsch, Savannah River Ecology Laboratory
Patrick V. Brady, U.S. DOE, Sandia National Laboratories
Daniel I. Kaplan, Westinghouse Savannah River Company
David M. Kargbo, Temple University
Irma McKnight, U.S. EPA, Office of Radiation and Indoor Air
Andrew Sowder, Savannah River Ecology Laboratory
In addition, a special thanks goes to Lindsey Bender from ORIA's Radiation Protection Division,
for her contributions in the production of this document.
Principal authorship of this guide was provided by the Department of Energy's Pacific Northwest
National Laboratory (PNNL) under the Interagency Agreement Number DW89937220-01-07.
Lynnette Downing served as the Department of Energy's Project Officer for this Interagency
Agreement. PNNL authors involved in preparation of Volume III include:
Kenneth M. Krupka
R. Jeffrey Serne
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TO COMMENT ON THIS GUIDE
OR PROVIDE INFORMATION FOR FUTURE UPDATES
Send all comments/updates to:
U.S. Environmental Protection Agency
Office of Radiation and Indoor Air
Attention: Understanding Variation in Partition (Ka) Values
1200 Pennsylvania Avenue, N.W. (6608J)
Washington, D.C. 20460-2001
or
wilhelm.ron@epa.gov
VI
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ABSTRACT
This three-volume report describes the conceptualization, measurement, and use of the partition
(or distribution) coefficient, Kd, parameter, and the geochemical aqueous solution and sorbent
properties that are most important in controlling adsorption/retardation behavior of selected
contaminants. The report is provided for technical staff from EPA and other organizations who are
responsible for prioritizing site remediation and waste management decisions. Volumes I and II were
published by the EPA in 1999.1 Volume I focuses on the Kd concept and methods for measurement of
Kd values. Particular attention is directed at providing an understanding of: (1) the use of Kd values in
formulating retardation factor (Rf), (2) the difference between the original thermodynamic Kd
parameter derived from ion-exchange literature and its "empiricized" use in contaminant transport
codes, and (3) the explicit and implicit assumptions underlying the use of the Kd parameter in
contaminant transport codes. A conceptual overview of chemical reaction models and their use in
addressing technical defensibility issues associated with data from Kd studies is also presented.
Volumes II and III provide "thumb-nail sketches" of the important aqueous speciation,
coprecipitation/dissolution, and adsorption processes affecting the sorption of selected inorganic
contaminants under oxidizing conditions. The Kd values listed in the literature for these contaminants
are also summarized. The contaminants discussed in Volume II include chromium, cadmium, cesium,
lead, plutonium, radon, strontium, thorium, tritium (3H), and uranium. Volume III, which is an
extension of Volume II, includes reviews of the sorption of americium, arsenic, curium, iodine,
neptunium, radium, and technetium. However, due to the limited number of IQ adsorption studies for
these contaminates and the large uncertainty, conservative minimum and maximums were not included.
References to related key experimental and review articles are included for possible further reading.
Both volumes (EPA, 1999b,1999c) can be downloaded and printed over the Internet at:
http://www.epa.gov/radiation/cleanup/partition.htrn This is found in the EPA Radiation, Information,
Radiation Publications, Topical Publications, Protecting People and the Environment, Fate & Transport
sections of the web site.
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CONTENTS
NOTICE ii
FOREWORD iii
ACKNOWLEDGMENTS v
FUTURE UPDATES vi
ABSTRACT vii
LIST OF FIGURES xii
LIST OF TABLES xiii
1.0 Introduction 1.1
2.0 The Kd Model 2.1
3.0 Methods, Issues, and Criteria for Measuring Kd Values 3.1
3.1 Laboratory Batch Methods 3.1
3.2 Laboratory Flow-Through Method 3.2
3.3 Other Methods 3.2
3.4 Issues 3.3
4.0 Application of Chemical Reaction Models 4.1
5.0 Contaminant Geochemistry and h^ Values 5.1
5.1 General 5.1
5.2 Americium Geochemistry and Kd Values 5.3
5.2.1 Overview: Important Aqueous-and Solid-Phase Parameters
Controlling Retardation 5.3
5.2.2 General Geochemistry 5.4
5.2.3 Aqueous Speciation 5.4
5.2.4 Dissolution/Precipitation/Coprecipitation 5.5
5.2.5 Adsorption/Desorption 5.5
5.2.5.1 Guidance for Screening Calculations of Adsorption 5.5
5.2.5.2 General Adsorption Studies 5.7
5.2.5.3 Kd Studies for Americium on Soil Materials 5.8
5.2.5.4 Published Compilations Containing ^ Values for Americium 5.10
5.2.5.5 Kd Studies of Americium on Pure Mineral, Oxide, and
Crushed Rock Materials 5.14
5.3 Arsenic Geochemistry and ^ Values 5.14
5.3.1 Overview: Important Aqueous-and Solid-Phase Parameters
Controlling Retardation 5.14
5.3.2 General Geochemistry 5.14
viii
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5.3.3 Aqueous Speciation 5.17
5.3.4 Dissolution/Precipitation/Coprecipitation 5.18
5.3.5 Adsorption/Desorption 5.18
5.3.5.1 Guidance for Screening Calculations of Adsorption 5.18
5.3.5.2 General Adsorption Studies 5.19
5.3.5.3 Kd Studies for Arsenic on Soil Materials 5.21
5.3.5.4 Published Compilations Containing Kd Values for Arsenic 5.22
5.3.5.5 Kd Studies of Arsenic on Pure Mineral, Oxide, and
Crushed Rock Materials 5.23
5.4 Curium Geochemistry and Kd Values 5.24
5.4.1 Overview: Important Aqueous-and Solid-Phase Parameters
Controlling Retardation 5.24
5.4.2 General Geochemistry 5.24
5.4.3 Aqueous Speciation 5.24
5.4.4 Dissolution/Precipitation/Coprecipitation 5.25
5.4.5 Adsorption/Desorption 5.26
5.4.5.1 Guidance for Screening Calculations of Adsorption 5.26
5.4.5.2 General Adsorption Studies 5.27
5.4.5.3 ^ Studies for Curium on Soil Materials 5.27
5.4.5.4 Published Compilations Containing Kd Values for Curium 5.28
5.4.5.5 Kd Studies of Curium on Pure Mineral, Oxide, and
Crushed Rock Materials 5.30
5.5 Iodine Geochemistry and Kd Values 5.30
5.5.1 Overview: Important Aqueous-and Solid-Phase Parameters
Controlling Retardation 5.30
5.5.2 General Geochemistry 5.30
5.5.3 Aqueous Speciation 5.31
5.5.4 Dissolution/Precipitation/Coprecipitation 5.32
5.5.5 Sorption/Desorption 5.33
5.5.5.1 Guidance for Screening Calculations of Adsorption 5.33
5.5.5.2 General Adsorption Studies 5.34
5.5.5.3 Kd Studies of Iodine on Soil Materials 5.36
5.5.5.4 Published Compilations Containing Kd Values for Iodine 5.43
5.5.5.5 Kd Studies of Iodine on Pure Mineral, Oxide, and
Crushed Rock Materials 5.44
5.6 Neptunium Geochemistry and Kd Values 5.45
5.6.1 Overview: Important Aqueous-and Solid-Phase Parameters
Controlling Retardation 5.45
5.6.2 General Geochemistry 5.46
5.6.3 Aqueous Speciation 5.47
5.6.4 Dissolution/Precipitation/Coprecipitation 5.50
5.6.5 Adsorption/Desorption 5.51
5.6.5.1 Guidance for Screening Calculations of Adsorption 5.51
ix
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5.6.5.2 General Adsorption Studies 5.51
5.6.5.3 Kd Studies for Neptunium on Soil Materials 5.52
5.6.5.4 Published Compilations Containing Kd Values for Neptunium 5.55
5.6.5.5 Kd Studies of Neptunium on Pure Mineral, Oxide, and
Crushed Rock Materials 5.58
5.7 Radium Geochemistry and Kd Values 5.58
5.7.1 Overview: Important Aqueous-and Solid-Phase Parameters
Controlling Retardation 5.58
5.7.2 General Geochemistry 5.58
5.7.3 Aqueous Speciation 5.60
5.7.4 Dissolution/Precipitation/Coprecipitation 5.60
5.7.5 Adsorption/Desorption 5.61
5.7.5.1 Guidance for Screening Calculations of Adsorption 5.61
5.7.5.2 General Adsorption Studies 5.62
5.7.5.3 ^ Studies for Radium on Soil Materials 5.63
5.7.5.4 Published Compilations Containing Kd Values for Radium 5.64
5.7.5.5 ^ Studies of Radium on Pure Mineral, Oxide, and
Crushed Rock Materials 5.66
5.8 Technetium Geochemistry and Kd Values 5.66
5.8.1 Overview: Important Aqueous-and Solid-Phase Parameters
Controlling Retardation 5.66
5.8.2 General Geochemistry 5.67
5.8.3 Aqueous Speciation 5.67
5.8.4 Dissolution/Precipitation/Coprecipitation 5.67
5.8.5 Adsorption/Desorption 5.70
5.8.5.1 Guidance for Screening Calculations of Adsorption 5.70
5.8.5.2 General Adsorption Studies 5.70
5.8.5.3 Kd Studies for Technetium on Soil Materials 5.71
5.8.5.4 Published Compilations Containing Kd Values for Technetium 5.76
5.8.5.5 Kd Studies of Technetium on Pure Mineral, Oxide, and
Crushed Rock Materials 5.79
5.9 Conclusions 5.79
6.0 References 6.1
X
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Appendix A -Acronyms, Abbreviations, Symbols, and Notation A.1
A.1 Acronyms and Abbreviations A.2
A.2 List of Symbols for Elements and Corresponding Names A.4
A.3 List of Symbols and Notation A.5
Appendix B - Definitions B.1
Appendix C - Americium Adsorption Studies C.1
C.1 Adsorption Studies of Americium on Single Mineral Phases C.2
C.2 Kd Studies of Americium on Crushed Rock Materials C.5
Appendix D - Arsenic Adsorption Studies D.1
D.1 Arsenic Adsorption Studies D.2
Appendix E - Curium Adsorption Studies E.1
E.1 Adsorption Studies of Curium on Single Mineral Phases E.2
E.2 Kd Studies of Curium on Crushed Rock Materials E.3
Appendix F - Iodine Adsorption Studies F.1
F.1 Adsorption Studies of Iodine on Single Mineral Phases F.2
F.2 Kd Studies of Iodine on Crushed Rock Materials F.6
Appendix G - Neptunium Adsorption Studies G.1
G.1 Adsorption Studies of Neptunium on Single Mineral Phases G.2
G.2 Kd Studies of Neptunium on Crushed Rock Materials G.7
Appendix H - Radium Adsorption Studies H.1
H.1 Adsorption Studies of Radium on Single Mineral Phases H.2
H.2 Kd Studies of Radium on Crushed Rock Materials H.4
Appendix I - Technetium Adsorption Studies 1.1
1.1 Adsorption Studies of Technetium on Single Mineral Phases 1.2
1.2 Kd Studies of Technetium on Crushed Rock Materials 1.2
XI
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LIST OF FIGURES
Page
Figure 5.1. Calculated aqueous speciation for Am(lll) as a function of pH 5.5
Figure 5.2. K,, values (ml/g) for Am(lll) adsorption on soil reported
by Routson et a/. (1975, 1977) (solid squares) and Sanchez et a/.
(1982) (solid triangles) 5.6
Figure 5.3. Eh-pH stability diagram for dominant arsenic aqueous
species at 25°C 5.15
Figure 5.4. Calculated aqueous speciation for As(V) as a function of pH 5.16
Figure 5.5. Calculated aqueous speciation for Cm(lll) as a function of pH 5.25
Figure 5.6. Eh-pH stability diagram for dominant iodine aqueous
species at 25°C 5.32
Figure 5.7. Eh-pH stability diagram for dominant neptunium aqueous
species at 25°C 5.48
Figure 5.8. Calculated aqueous speciation for Np(V) as a function of pH 5.49
Figure 5.9. Eh-pH stability diagram for the dominant technetium aqueous
species at 25°C 5.68
Xll
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LIST OF TABLES
Table 5.1 Estimated mean composition of river water of the world
from Hem (1985) 5.3
Table 5.2. Americium(lll) aqueous species 5.4
Table 5.3. Measured Kd values (ml/g) for americium as a function of pH
for Hudson River estuary environment [Sanchez et al. (1982)] 5.8
Table 5.4. Measured americium ^ values (ml/I) and soil properties for
soils studied by Nishita et al. (1981) 5.9
Table 5.5. Properties of soils used in Kd measurements by
Routson etal. (1975, 1977) 5.10
Table 5.6. Americium Kd values (ml/g) listed by Thibault et al.
(1990, Tables 4 to 8) 5.11
Table 5.7. Americium Kd values (ml/g) listed by McKinley and
Scholtis (1993, Tables 1, 2, and 4) from sorption databases used
by different international organizations for performance assessments
of repositories for radioactive wastes 5.13
Table 5.8. Measured arsenic ^ values (ml/g) based on analyses of an
arsenic-contaminated aquifer at a Superfund Site (Mariner et al., 1996) 5.22
Table 5.9. Measured curium Kd values (ml/l) and soil properties for soils
studied by Nishita etal. (1981) 5.28
Table 5.10. Curium ^ values (ml/g) listed by Thibault et al.
(1990, Tables 4 to 8) 5.29
Table 5.11. Kd values (ml/g) for the adsorption of iodine to sediments
from the Hanford Site in southeastern Washington state
(Kaplan etal., 1998a,1998b) 5.38
Table 5.12. Median Kd values (ml/g) for iodide and iodate measured in
deionized water [Yoshida et al. (1998)] 5.40
Table 5.13. Geometric mean ^ values (ml/g) for iodine measured in
untreated river water under oxic conditions (Bird and Schwartz, 1996) 5.41
Xlll
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Table 5.14. Values of Kd (ml/g) measured by Muramatsu et al. (1990) for the
sorption of I" and I0g on soil and selected soil components of an andosol 5.43
Table 5.15. Iodine ^ values (ml/g) listed by Thibault et al.
(1990, Tables 4 to 8) 5.44
Table 5.16. Iodine Kd values (ml/g) listed by McKinley and Scholtis
(1993, Tables 1, 2, and 4) from sorption databases used by
different international organizations for performance assessments
of repositories for radioactive wastes 5.45
Table 5.17. Np(V) aqueous species 5.49
Table 5.18. Neptunium(V) Kd values (ml/g) measured for three sediments
by Kaplan etal. (1996) 5.52
Table 5.19. Measured Np(V) Kd values (ml/I) and soil properties for soils
studied by Nishita et al. (1981) 5.53
Table 5.20. Properties of soils used in Kd measurements by
Routson etal. (1975, 1977) 5.54
Table 5.21. Neptunium Kd values (ml/g) measured for Washington
and South Carolina soil samples in Ca(N03)2 and NaN03 solutions
by Routson etal. (1975, 1977) 5.54
Table 5.22. Neptunium Kd values (ml/g) listed by Thibault et al.
(1990, Tables 4 to 8) 5.56
Table 5.23. Neptunium Kd values (ml/g) listed by McKinley and Scholtis
(1993, Tables 1, 2, and 4) from sorption databases used by
different international organizations for performance assessments
of repositories for radioactive wastes 5.57
Table 5.24. Ionic radii (A) for alkaline earth elements (Molinari
and Snodgrass, 1990) 5.59
Table 5.25. Radium h^ values (ml/g) as function of calcium
concentration [Nathwani and Phillips, 1979b)] 5.64
Table 5.26. Properties of soil samples for which Kd values are given in
Nathwani and Phillips (1979b) 5.65
Table 5.27. Radium h^ values (ml/g) measured by Serne (1974) for sandy,
arid soil samples from Utah 5.65
XIV
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Table 5.28. Radium Kd values (ml/g) listed by Thibault et al.
(1990, Tables 4 to 8) 5.66
Table 5.29. Technetium ^ values (ml/g) measured for Hanford sediments
under oxidizing conditions (Kaplan et al. (1998a, 1998b) 5.73
Table 5.30. Average technetium Kd values (ml/g) (based on three
replicates) for technetium on Hanford sediments after 21.5 days
from Gee and Campbell (1980) 5.75
Table 5.31. Technetium Kd values (ml/g) listed by Thibault et al.
(1990, Tables 4 to 8) 5.77
Table 5.32. Technetium ^ values (ml/g) listed by McKinley and Scholtis
(1993, Tables 1, 2, and 4) from sorption databases used by
different international organizations for performance assessments
of repositories for radioactive wastes 5.78
Table 5.33. Selected chemical and transport properties of the contaminants 5.80
Table 5.34. Distribution of dominant contaminant species at 3 pH values
for an oxidizing water described in Tables 5.1 5.81
Table 5.35. Some of the more important aqueous- and solid-phase
parameters affecting contaminant sorption 5.82
Table C.1 Americium adsorption studies on pure mineral and oxide phases C.3
Table C.2 Americium adsorption studies on crushed rock and related materials .... C.5
Table D.1 Arsenic adsorption studies on pure mineral and oxide phases D.2
Table E.1 Curium adsorption studies on pure mineral and oxide phases E.2
Table E.2 Curium adsorption studies on crushed rock and related materials E.3
Table F.1 Iodine adsorption studies on pure minerals and oxide phases F.3
Table F.2 Iodine adsorption studies on crushed rock and related materials F.7
Table G.1 Neptunium adsorption studies on pure mineral and oxide phases G.2
Table G.2 Measured Np(V) Kd values (ml/I) for fracture-coating minerals
as a function of total dissolved solids (TDS) and normal versus low
oxygen conditions G.4
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Table G.3 Neptunium adsorption studies on crushed rock and
related materials G.7
Table H.1 Radium adsorption studies on pure mineral and oxide phases H.3
Table H.2 Radium adsorption studies on crushed rock and related materials H.5
Table 1.1 Technetium adsorption studies on pure mineral and oxide phases 1.3
Table 1.2 Technetium adsorption studies on crushed rock and related materials 1.5
XVI
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1.0 Introduction
The objective of the report is to provide a reasoned and documented discussion on the technical issues
associated with the measurement and selection of partition (or distribution) coefficient, Kd/'2 values
and their use in formulating the retardation factor, Rf. The contaminant retardation factor (Rf) is the
parameter commonly used in transport models to describe the chemical interaction between the
contaminant and geological materials (i.e., soil, sediments, rocks, and geological formations, henceforth
simply referred to as soils3). It includes processes such as surface adsorption, absorption into the soil
structure, precipitation, and physical filtration of colloids. Specifically, it describes the rate of
contaminant transport relative to that of groundwater. This report is provided for technical staff from
EPA and other organizations who are responsible for prioritizing site remediation and waste
management decisions. The three-volume report describes the conceptualization, measurement, and
use of the Kd parameter; and geochemical aqueous solution and sorbent properties that are most
important in controlling the adsorption/retardation behavior of a selected set of contaminants.
This review is not meant to assess or judge the adequacy of the Kj approach used in modeling tools for
estimating adsorption and transport of contaminants and radionuclides. Other approaches, such as
surface complexation models, certainly provide more robust mechanistic approaches for predicting
contaminant adsorption. However, as one reviewer noted, "Kd's are the coin of the realm in this
business." For better or worse, the Kj model is an integral part of current methodologies for modeling
contaminant and radionuclide transport and risk analysis.
The Kd concept, its use in fate and transport computer codes, and the methods for the measurement of
Kd values are discussed in detail in Volume I and briefly introduced in Chapters 2 and 3 in Volumes II
(EPA, 1999c) and III. Particular attention is directed at providing an understanding of: (1) the use of
Kd values in formulating Rf, (2) the difference between the original thermodynamic Kd parameter
derived from the ion-exchange literature and its "empiricized" use in contaminant transport codes, and
(3) the explicit and implicit assumptions underlying the use of the Kd parameter in contaminant
transport codes.
As typically used in fate and contaminant transport calculations, the K^ is defined as the ratio of the
contaminant concentration associated with the solid to the contaminant concentration in the
surrounding aqueous solution when the system is at equilibrium. Soil chemists and geochemists
knowledgeable of sorption processes in natural environments have long known that generic or default
Kd values can result in significant error when used to predict the impacts of contaminant migration or
site-remediation options. To address some of this concern, modelers often incorporate a degree of
Throughout this report, the term "partition coefficient" will be used to refer to the K^ "linear isotherm" sorption
model. It should be noted, however, that the terms "partition coefficient" and "distribution coefficient" are used
interchangeably in the literature for the Kd model.
A list of acronyms, abbreviations, symbols, and notation is given in Appendix A. A list of definitions is given in
Appendix B.
The terms "sediment" and "soil" have particular meanings depending on one's technical discipline. For example,
the term "sediment" is often reserved for transported and deposited particles derived from soil, rocks, or biological
material. "Soil" is sometimes limited to referring to the top layer of the earth's surface, suitable for plant life. In this
report, the term "soil" was selected with concurrence of the EPA Project Officer as a general term to refer to all
unconsolidated geologic materials.
MP &
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conservatism into their calculations by selecting limiting or bounding conservative Kj values. For
example, the most conservative Kd value from the perspective of off-site risks due to contaminant
migration through subsurface natural soil and groundwater systems is to assume that the soil has little
or no ability to slow (retard) contaminant movement (i.e., a minimum bounding Kd value). Conse-
quently, the contaminant would travel in the direction and at the rate of water. Such an assumption
may in fact be appropriate for certain contaminants such as tritium, but may be too conservative for
other contaminants, such as thorium or plutonium, which react strongly with soils and may migrate 102
to 106 times more slowly than the water.
On the other hand, when estimating the risks and costs associated with on-site remediation options, a
maximum bounding Kd value provides an estimate of the maximum concentration of a contaminant or
radionuclide sorbed to the soil. Conservatism for remediation calculations would tend to err on the
side of underestimating the extent of contaminant desorption that would occur in the aquifer once
pump-and-treat or soil flushing treatments commenced. Such an estimate would provide an upper
limit to time, money, and work required to extract a contaminant from a soil. This would be
accomplished by selecting a Kd from the upper range of literature values.
In some instances because of long groundwater flow paths, extremely arid site characteristics, or
presence of impermeable soils, the final results of risk and transport calculations for some
contaminants may be insensitive to the Kd value even when selected within the range of technically-
defensible, limiting minimum and maximum Kd values. However, for most situations that are sensitive
to the selected Kd value, site-specific Kd values are essential for obtaining defensible risk and transport
predictions.
The Kd is usually a measured parameter that is obtained from laboratory experiments. The 5 general
methods used to measure Kd values are reviewed in Volume I (EPA, 1999b). These methods include
the batch laboratory method, the column laboratory method, field-batch method, field modeling
method, and soil organic carbon/water partition coefficient ( Koc) method. Volume I (EPA, 1999b)
identifies what ancillary information is needed regarding the adsorbent (soil), solution (contaminated
ground-water or process waste water), contaminant (concentration, valence state, speciation
distribution), and laboratory details (spike addition methodology, phase separation techniques, contact
times). The advantages, disadvantages, and, perhaps more importantly, the underlying assumptions of
each method are also presented.
A conceptual overview of geochemical modeling calculations and computer codes as they pertain to
evaluating Kd values and modeling of adsorption processes is discussed in detail in Volume I (EPA,
1999b) and briefly described in Chapter 4 of Volumes II (EPA, 1999c) and III. The use of
geochemical codes in evaluating aqueous speciation, solubility, and adsorption processes associated
with contaminant fate studies is reviewed. This approach is compared to the traditional calculations
that rely on the constant Kd construct. The use of geochemical modeling to address quality assurance
and technical defensibility issues concerning available Kd data and the measurement of Kd values are
also discussed. The geochemical modeling review includes a brief description of the EPA's
MINTEQA2 geochemical code and a summary of the types of conceptual models it contains to
quantify adsorption reactions. The status of radionuclide thermodynamic and contaminant adsorption
model databases for the MINTEQA2 code is also reviewed.
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The main focus of Volumes II (EPA, 1999c) and III is to: (1) provide a "thumb-nail sketch" of the key
geochemical processes affecting the sorption of a selected set of contaminants; (2) provide references
to related key experimental and review articles for further reading; (3) identify the important aqueous-
and solid-phase parameters controlling the sorption of these contaminants in the subsurface
environment; and (4) identify, when possible, minimum and maximum conservative Kd values for each
contaminant (as a function key geochemical processes affecting their sorption). The contaminants
chosen for the first phase of this project, Volume II (EPA, 1999c), included reviews on cadmium,
cesium, chromium, lead, plutonium, radon, strontium, thorium, tritium (3H), and uranium. This
document (Volume III) represents the second phase of this project, and is an extension of Volume II.
Volume III provides "thumb-nail sketches" of the key geochemical processes affecting the sorption of
americium, arsenic(V) (as arsenate), curium, iodine [-1 (as iodide) and +5 (iodate)], neptunium, radium,
and technetium(VII) (as pertechnetate) to soils.
The selection of these contaminants by EPA and Pacific Northwest National Laboratory (PNNL)
project staff was based on two criteria. First, the contaminant had to be of high priority to site
remediation or risk assessment activities of EPA, DOE, and/or NRC. Second, because the available
funding precluded a review of all contaminants that met the first criteria, a subset was selected to
represent categories of contaminants based on their chemical behavior. The six nonexclusive
categories are:
• Cations - americium, cadmium, cesium, curium, neptunium, plutonium, radium, strontium,
thorium, and uranium.
• Anions - arsenic(V) (as arsenate), chromium(VI) (as chromate), iodine [-1 (as iodide) and +5
(iodate)], technetium(VII) (as pertechnetate) and uranium(VI) complexes (e.g., uranyl carbonate
complexes).
• Radionuclides - americium, cesium, curium, iodine, neptunium, plutonium, radium, radon,
strontium, technetium, thorium, tritium (3H), and uranium.
• Conservatively transported contaminants - tritium (3H) and radon.
• Nonconservatively transported contaminants - other than tritium (3H) and radon.
• Redox sensitive elements - arsenic, chromium, iodine, neptunium, plutonium, technetium, and
uranium.
The general geochemical behaviors discussed in this report can be used by analogy to estimate the
geochemical interactions of similar elements for which data are not available. For example,
contaminants present primarily in anionic form, such as Cr(VI), tend to adsorb to a limited extent to
soils. Thus, one might generalize that other anions, such as nitrate, chloride, and U(VI)-anionic
complexes, would also adsorb to a limited extent. Literature on the adsorption of these three solutes
show no or very little adsorption.
The concentration of contaminants in groundwater is controlled primarily by the amount of contam-
inant present at the source; rate of release from the source; hydrologic factors such as dispersion,
advection, and dilution; and a number of geochemical processes including aqueous geochemical
processes, adsorption/desorption, precipitation, and diffusion. To accurately predict contaminant
transport through the subsurface, it is essential that the important geochemical processes affecting
contaminant transport be identified and, perhaps more importantly, accurately described in a
mathematically and scientifically defensible manner. Dissolution/precipitation and
adsorption/desorption are usually the most important processes affecting contaminant interaction with
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soils. Dissolution/precipitation is more likely to be the key process where chemical nonequilibrium
exists, such as at a point source, an area where high contaminant concentrations exist, or where steep
pH or oxidation-reduction (redox) gradients exist. Adsorption/desorption will likely be the key process
controlling contaminant migration in areas where chemical steady state exist, such as in areas far from
the point source. Diffusion flux spreads solute via a concentration gradient (i.e., Pick's law). Diffusion
is a dominant transport mechanism when advection is insignificant, and is usually a negligible transport
mechanism when water is being advected in response to various forces.
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2.0 The Kd Model
The simplest and most common method of estimating contaminant retardation is based on the
partition (or distribution) coefficient, FQ. The Kd parameter is a factor related to the partitioning of a
contaminant between the solid and aqueous phases. It is an empirical unit of measurement that
attempts to account for various chemical and physical retardation mechanisms that are influenced by a
myriad of variables. The Kd metric is the most common measure used in transport codes to describe
the extent to which contaminants are sorbed to soils. It is the simplest, yet least robust model available.
A primary advantage of the Kd model is that it is easily inserted into hydrologic transport codes to
quantify reduction in the rate of transport of the contaminant relative to groundwater, either by
advection or diffusion. Technical issues, complexities, and shortcomings of the Kd approach to
describing contaminant sorption to soils are summarized in detail in Chapter 2 of Volume I (EPA,
1999b). Particular attention is directed at issues relevant to the selection of K^ values from the
literature for use in transport codes.
The partition coefficient, Kd, is defined as the ratio of the quantity of the adsorbate adsorbed per mass
of solid to the amount of the adsorbate remaining in solution at equilibrium. For the reaction
A + Cj = Aj, (2.1)
the mass action expression for Kd is
„ _ Mass of Adsorbate Sorbed _ Aj
Mass of Adsorbate in Solution Cj
where A = free or unoccupied surface adsorption sites
Q = total dissolved adsorbate remaining in solution at equilibrium
Aj = amount of adsorbate on the solid at equilibrium.
The Kd is typically given in units of ml/g. Describing the Kd in terms of this simple reaction assumes
that A is in great excess with respect to Q and that the activity of A; is equal to 1.
Chemical retardation, Rf, is defined as,
Rf = - . (2.3)
c
where vp = velocity of the water through a control volume
vc = velocity of contaminant through a control volume.
The chemical retardation term does not equal unity when the solute interacts with the soil; almost
always the retardation term is greater than 1 due to solute sorption to soils. In rare cases, the
retardation factor is actually less than 1, and such circumstances are thought to be caused by anion
exclusion (See Volume I, Section 2.7). Knowledge of the Kd and of media bulk density and porosity
for porous flow, or of media fracture surface area, fracture opening width, and matrix diffusion
attributes for fracture flow, allows calculation of the retardation factor. For porous flow with saturated
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moisture conditions, the Rf is defined as
ne
(2.4)
where pb = porous media bulk density (mass/length3)
ne = effective porosity of the media at saturation.
The Kd parameter is valid only for a particular adsorbent and applies only to those aqueous chemical
conditions (e.g., adsorbate concentration, solution/electrolyte matrix) in which it was measured. Site-
specific Kd values should be used for site-specific contaminant and risk assessment calculations.
Ideally, site-specific Kd values should be measured for the range of aqueous and geological conditions
in the system to be modeled. However, literature-derived Kd values are commonly used for screening
calculations. Suitable selection and use of literature-derived Kd values for screening calculations of
contaminant transport is not a trivial matter. Among the assumptions implicit with the Kd construct
is: (1) only trace amounts of contaminants exist in the aqueous and solid phases, (2) the relationship
between the amount of contaminant in the solid and liquid phases is linear, (3) equilibrium conditions
exist, (4) equally rapid adsorption and desorption kinetics exists, (5) it describes contaminant
partitioning between one sorbate (contaminant) and one sorbent (soil), and (6) all adsorption sites are
accessible and have equal adsorption binding energies. Many of these assumptions are not met for
groundwater/soil environments. Thus, literature-derived Kd values should be used only to predict
transport in systems similar to those used in the laboratory and field to measure the Kd. Variation in
either the soil or aqueous chemistry of a system can result in extremely large differences in Kd values.
A more robust approach than using a single Kd to describe the partitioning of contaminants between
the aqueous and solid phases is the parametric-K^ model. This model varies the Kd value according to
the chemistry and mineralogy of the system at the node being modeled. The parametric-Kd value,
unlike the constant-Kd value, is not limited to a single set of environmental conditions. Instead, it
describes the sorption of a contaminant in the range of environmental conditions used to create the
parametric-Kd equations. These types of statistical relationships are devoid of causality and therefore
provide no information on the mechanism by which the radionuclide partitioned to the solid phase,
whether it be by adsorption, absorption, or precipitation. Understanding these mechanisms is
extremely important relative to estimating the mobility of a contaminant.
When the parametric-Kd model is used in the transport equation, the code must also keep track of the
current value of the independent variables at each point in space and time to continually update the
concentration of the independent variables affecting the Kd value. Thus, the code must track many
more parameters and some numerical solving techniques (such as closed-form analytical solutions). It
can no longer be used to perform the integration necessary to solve for the Kd value and/or retardation
factor, Rf. Generally, computer codes that can accommodate the parametric-Kd model use a chemical
subroutine to update the Kd value used to determine the RF, when called for by the main transport
code. The added complexity in solving the transport equation with the parametric-Kd sorption model
and its empirical nature may be the reasons this approach has been used sparingly.
Mechanistic models explicitly accommodate for the dependency of K^ values on contaminant concen-
tration, charge, competing ion concentration, variable surface charge on the soil, and solution species
distribution. Incorporating mechanistic adsorption concepts into transport models is desirable because
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the models become more robust and, perhaps more importantly from the standpoint of regulators and
the public, scientifically defensible. However, truly mechanistic adsorption models are rarely, if ever,
applied to complex natural soils. The primary reason for this is because natural mineral surfaces are
very irregular and difficult to characterize. These surfaces consist of many different microcrystalline
structures that exhibit quite different chemical properties when exposed to solutions. Thus,
examination of the surface by virtually any experimental method yields only averaged characteristics of
the surface and the interface.
Less attention will be directed to mechanistic models because they are not extensively incorporated into
the majority of EPA, DOE, and NRC modeling methodologies. The complexity of installing these
mechanistic adsorption models into existing transport codes is formidable. Additionally, these models
also require a more extensive database collection effort than will likely be available to the majority of
EPA, DOE, and NRC contaminant transport modelers. A brief description of the state of the science
of mechanistic adsorption modeling is presented in Volume I (EPA, 1999b) primarily to provide a
paradigm for sorption processes. Readers should note, that since the completion of Volumes I and II
(EPA, 1999b, 1999c), Brown et al. (1999) published an extensive review of the interactions of metal
oxide surfaces with aqueous solutions and microbial organisms. This review includes a detailed
discussion of the theory and thermodynamic models for describing the adsorption of chemical species
to oxide surfaces.
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3.0 Methods, Issues, and Criteria for Measuring Kd Values
There are five general methods used to measure Kd values: the batch laboratory method, laboratory
flow-through (or column) method, field-batch method, field modeling method, and Koc method. The
Koc method is specific to adsorption of organics. These methods and the associated technical issues are
described in detail in Chapter 3 of Volume I (EPA, 1999b). Each method has advantages and
disadvantages, and perhaps more importantly, each method has its own set of assumptions for
calculating Kd values from experimental data. Consequently, it is not only common, but expected that
Kd values measured by different methods will produce different values.
3.1 Laboratory Batch Method
Batch tests are commonly used to measure Kd values. A batch test is conducted by spiking a solution
with the element of interest, mixing the spiked solution with a solid for a specified period of time,
separating the solution from the solid, and measuring the concentration of the spiked element
remaining in solution. The concentration of contaminant associated with the solid is determined by the
difference between initial and final contaminant concentration in solution. The primary advantage of
the batch method is that such experiments can be completed quickly for a wide variety of elements and
chemical environments. The primary disadvantage of the batch technique for measuring Kd is that it
does not necessarily reproduce the chemical reaction conditions that take place in the real environment.
For instance, in a soil column, water passes through at a finite rate and both reaction time and degree
of mixing between water and soil can be much less than those occurring in a laboratory batch test.
Consequently, Kd values from batch experiments can be high relative to the extent of sorption
occurring in a real system, and thus result in an estimate of contaminant retardation that is too large.
Another disadvantage of batch experiments is that they do not accurately simulate desorption of the
radionuclides or contaminants from a contaminated soil or solid waste source. The Kd values are
frequently used with the assumption that adsorption and desorption reactions are reversible. This
assumption is contrary to most experimental observations that show that the desorption process is
appreciably slower than the adsorption process, a phenomenon referred to as hysteresis. The rate of
desorption may even go to zero, yet a significant mass of the contaminant remains sorbed on the soil.
Thus, use of Kd values determined from batch adsorption tests in contaminant transport models is
generally considered to provide estimates of contaminant remobilization (release) from soil that are too
large (i.e., estimates of contaminant retention that are too low).
Inherent in the Kd "linear isotherm" adsorption model is the assumption that adsorption of the
contaminant of interest is independent of its concentration in the aqueous phase. Partitioning of a
contaminant on soil can often be described using the Kd model, but typically only for low contaminant
concentrations as would exist some distance away (far field) from the source of contamination. It is
common knowledge that contaminant adsorption on soils can deviate from the linear relationship
required by the Kd construct. This is possible for conditions as might exist in leachates or groundwater
near waste sources where contaminant concentrations are large enough to affect the saturation of
surface adsorption sites. In these latter situations, non-linear isotherm models [see Section 2.3.3 in
Volume I (EPA, 1999b)] are used to describe the case where sorption relationships deviate from
linearity.
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3.2 Laboratory Flow-Through Method
Flow-through column experiments are intended to provide a more realistic simulation of dynamic field
conditions and to quantify the movement of contaminants relative to groundwater flow. It is the
second most common method of determining Kd values. The basic experiment is completed by
passing a liquid spiked with the contaminant of interest through a soil column. The column experi-
ment combines the chemical effects of sorption and the hydrologic effects of groundwater flow
through a porous medium to provide an estimate of retarded movement of the contaminant of interest.
The retardation factor (a ratio of the velocity of the contaminant to that of water) is measured directly
from the experimental data. A Kd value can be calculated from the retardation factor. It is frequently
useful to compare the back-calculated Kd value from these experiments with those derived directly
from the batch experiments to evaluate the influence of limited interaction between solid and solution
imposed by the flow-through system.
One potential advantage of the flow-through column studies is that the retardation factor can be
inserted directly into the transport code. However, if the study site contains different hydrological
conditions (e.g., porosity and bulk density) than the column experiment, then a Kd value needs to be
calculated from the retardation factor. Another advantage is that the column experiment provides a
much closer approximation of the physical conditions and chemical processes occurring in the field site
than a batch sorption experiment. Column experiments permit the investigation of the influence of
limited spatial and temporal (nonequilibrium) contact between solute and solid have on contaminant
retardation. Additionally, the influence of mobile colloid facilitated transport and partial saturation can
be investigated. A third advantage is that both adsorption or desorption reactions can be studied. The
predominance of one mechanism of adsorption or desorption over another cannot be predicted a priori
and therefore generalizing the results from one set of laboratory experimental conditions to field
conditions is never without some uncertainty. Ideally, flow-through column experiments would be
used exclusively for determining Kd values, but equipment cost, time constraints, experimental
complexity, and data reduction uncertainties discourage more extensive use. Another important issue
for column studies is that a flow model (e.g., piston flow or a mobile/immobile flow system) must be
assumed to calculate a Kd value. A different result may be obtained depending on which flow model is
chosen. The Kd values derived from column studies are also conditional on the flow rate; generally, the
faster the flow rate, lower the calculated Kd.
3.3 Other Methods
Less commonly used methods include the Koc method, in-situ batch method, and the field modeling
method. The Koc method is a very effective indirect method of calculating Kd values, however, it is
only applicable to hydrophobic organic compounds. The Koc method becomes increasingly inaccurate
as the organic contaminant is increasingly hydrophilic, because hydrophilic organic compounds can also
adsorb to mineral surfaces, in addition to partitioning to organic matter. The in-situ batch method
requires that paired soil and groundwater samples be collected directly from the aquifer system being
modeled and then measuring directly the amount of contaminant on the solid and liquid phases. The
advantage of this approach is that the precise solution chemistry and solid phase mineralogy existing in
the study site is used to measure the Kd value. However, this method is not used often because of the
analytical problems associated with measuring the exchangeable fraction of contaminant on the solid
phase. Finally, the field modeling method of calculating Kd values uses groundwater monitoring data
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and source term data to calculate a Kd value. The accuracy of the field modeling method in highly
dependent on the accuracy of the other input parameters used in the calculation. For instance, if the
dispersion value or the hydraulic conductivity is incorrect, it is very unlikely that the true Kd value will
be calculated. One key drawback to this technique is that it is very model dependent. Because the
calculated Kd values are model dependent and highly site specific, the derived Kd values should not be
used for contaminant transport calculations at other sites.
3.4 Issues
A number of issues exist concerning the measurement of Kd values and the selection of Kd values from
the literature. These issues are discussed with supporting references in Volume I (EPA, 1999b), and are
briefly summarized below. These issues include: using simple versus complex systems to measure Kj
values, field variability, the "gravel issue," the "colloid issue," and contaminant solubility limits. Soils
are a complex mixture containing solid, gaseous, and liquid phases. Each phase contains several
different constituents. The use of simplified systems containing single mineral phases and aqueous
phases with one or two dissolved species has provided valuable paradigms for understanding sorption
processes in more complex, natural systems. However, the Kd values generated from these simple
systems are generally of little value for importing directly into transport models. Values for transport
models should be generated from geologic materials from or similar to the study site.
The "gravel issue" is the problem that transport modelers face when converting laboratory-derived Kd
values based on experiments conducted with the <2-mm fraction into values that can be used in
systems containing particles >2 mm in size (Kaplan eta!., 2000c). No standard methods exist to
address this issue. There are many subsurface soils dominated by cobbles, gravel, or boulders. To base
the Kd values on the <2-mm fraction, which may constitute only a low percent of the soil volume but is
the most chemically reactive fraction, may grossly overestimate the actual Kd of the aquifer. Two
general approaches have been proposed to address this issue. The first is to assume that all particles
>2-mm have a Kd = 0 ml/g. Although this assumption is incorrect; i.e., we know that cobbles, gravel,
and boulders do in fact sorb contaminants, the extent to which sorption occurs on these larger particles
may be small. The second approach is to normalize laboratory-derived Kd values by soil surface area.
Theoretically, this latter approach is more satisfying because it permits some sorption to occur on the
>2-mm fraction and the extent of the sorption is proportional to the surface area. The underlying
assumptions in this approach are that the mineralogy is similar in the less than 2- and greater than
2-mm fractions and that the sorption processes occurring on the smaller particles are similar to those
that occur on the larger particles.
Spatial variability provides additional complexity to understanding and modeling contaminant retention
to subsurface soils. The extent to which contaminants partition to soils changes as field mineralogy
and chemistry change. Thus, a single Kd value is almost never sufficient for an entire study site and
should change as chemically important environmental conditions change. Three approaches used to
vary Kd values in transport codes are the Kd look-up table approach, the parametric-Kd approach, and
the mechanistic Kd approach. The extent to which these approaches are presently used and the ease of
incorporating them into a flow model varies greatly. Parametric-Kd values typically have limited
environmental ranges of application. Mechanistic Kd values are limited to uniform solid and aqueous
systems with little application to heterogeneous soils existing in nature. The easiest and the most
common variable-Kd model interfaced with transport codes is the look-up table. In Kd look-up tables,
separate Kd values are assigned to a matrix of discrete categories defined by chemically important
3.3
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ancillary parameters. No single set of ancillary parameters, such as pH and soil texture, is universally
appropriate for defining categories in Kd look-up tables. Instead, the ancillary parameters must vary in
accordance to the geochemistry of the contaminant. It is essential to understand fully the criteria and
process used for selecting the values incorporated in such a table. Differences in the criteria and
process used to select Kd values can result in appreciably different Kd values. Examples are presented
in this volume.
Contaminant transport models generally treat the subsurface environment as a two-phase system in
which contaminants are distributed between a mobile aqueous phase and an immobile solid phase (e.g.,
soil). An increasing body of evidence indicates that under some subsurface conditions, components of
the solid phase may exist as mobile colloids1 that may be transported with the flowing water.
Subsurface mobile colloids originate from (1) the dispersion of surface or subsurface soils,
(2) decementation of secondary mineral phases, and (3) homogeneous precipitation of groundwater
constituents. Association of contaminants with this additional mobile phase may enhance not only the
amount of contaminant that is transported, but also the rate of contaminant transport. Most current
approaches to predicting contaminant transport ignore this mechanism not because it is obscure or
because the mathematical algorithms have not been developed, but because little information is
available on the occurrence, the mineralogical properties, the physicochemical properties, or the
conditions conducive to the generation of mobile colloids. There are two primary problems associated
with studying colloid-facilitated transport of contaminants under natural conditions. First, it is difficult
to collect colloids from the subsurface in a manner which minimizes or eliminates sampling artifacts.
Second, it is difficult to unambiguously delineate between the contaminants in the mobile-aqueous and
mobile-solid phases.
Some contaminants, such as americium, curium, and others, are very insoluble under certain
groundwater conditions. Therefore, care must be taken not to exceed their solubilities when measuring
their Kd values using laboratory batch and flow-through column techniques. Values of Kd determined
under such conditions will overestimate the retardation due to the adsorption of the contaminant.
Investigators must carefully analyze their results to insure that the Kd values were not measured at
oversaturated conditions. If batch Kd measurements are completed for a range of initial contaminant
concentrations for a fixed set of geochemical conditions (e.g., pH), the final dissolved concentrations of
the contaminant may eventually reach a constant value for those solution concentrations exceeding the
solubility of the contaminant. The resulting sorption isotherm will be a vertical line at the solubility
limit when plotting the final concentrations of sorbed contaminant (y-axis) as a function of final
concentrations of dissolved contaminant (x-axis) [see "Isotherm Adsorption Models" in Section 2.3.3
in Volume I (EPA, 1999b)]. For Kd values measured using the flow-through column technique with
constant step input [see Section 3.2.3 in Volume I (EPA, 1999b)], the ratio of the contaminant
concentration in the effluent (Ceff) to that in the input at time 0 (zero) (C0) for the effluent (or break-
through) curve will never achieve a value of 1.0 when contaminant precipitation has occurred in the soil
column. For flow-through measurements with a pulse input, the total mass of contaminant that is
determined in the collected effluent after Ceff returns to zero will not equal the total mass injected into
A colloid is any fine-grained material, sometimes limited to the particle-size range of <0.00024 mm (i.e., smaller than
clay size), that can be easily suspended (Bates and Jackson, 1980). In its original sense, the definition of a colloid
included any fine-grained material that does not occur in crystalline form. The geochemistry of colloid systems is
discussed in detail in sources such as Yariv and Cross (1979) and the references therein.
3.4
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the column due to irreversible precipitation. Moreover, the shape of effluent break-through curve
(Ceff/C0) will not be very Gaussian in form and will have a protracted tail.
When planning Kd studies and analyzing the results, investigators can use chemical reaction computer
models, like EPA's MINTEQA2 geochemical code, to estimate from thermodynamic principles
solubility limits for a contaminant as a function of pH, redox, concentrations of complexing ligands,
and temperature [see Section 5 in Volume I (EPA, 1999b)]. The results of these modeling calculations
can be used to set limits for the maximum initial contaminant concentrations to be used in the Kd
measurements and/or to alert investigators that solubility limits may have been exceeded during the
course of study. An example of this type of application is described in Section 5.2.4.2 in Volume I
(EPA, 1999b).
It is incumbent upon the transport modeler to understand the strengths and weaknesses of the
different Kd methods, and perhaps more importantly, the underlying assumption of the methods in
order to properly select Kd values from the literature. The Kd values reported in the literature for any
given contaminant may vary by as much as six orders of magnitude. An understanding of the
important geochemical processes and knowledge of the important ancillary parameters affecting the
sorption chemistry of the contaminant of interest is necessary for selecting appropriate Kd value(s) for
contaminant transport modeling.
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4.0 Application of Chemical Reaction Models
Computerized chemical reaction models based on thermodynamic principles may be used to calculate
processes such as aqueous complexation, oxidation/reduction, adsorption/desorption, and mineral
precipitation/dissolution for contaminants in soil-water systems. The capabilities of a chemical
reaction model depend on the models incorporated into its computer code and the availability of
thermodynamic and/or adsorption data for aqueous and mineral constituents of interest. Chemical
reaction models, their utility to understanding the solution chemistry of contaminants, and the
MINTEQA2 model in particular are described in detail in Chapter 5 of Volume I (EPA, 1999b).
Chemical reaction models can be used to support evaluations of Kd values and related contaminant
migration and risk assessment modeling predictions. Most models include options for calculating
(1) the distribution of complexed and uncomplexed aqueous species for a specified water composition,
(2) the dissolved concentration of an element based on the solubility of solids containing that element,
and in some instances, (3) the mass of an element adsorbed by ion exchange or surface complexation
onto a single, pure mineral phase. Results from aqueous complexation calculations indicate the ionic
state and composition of the dominant aqueous species for a dissolved contaminant present in a soil-
water system. This information may in turn be used to substantiate the conceptual model being used
for calculating the adsorption of a particular contaminant. Solubility calculations provide a means of
predicting technically defensible maximum concentration limits for contaminants as a function of key
composition parameters (e.g., pH) for any specific soil-water system. These values may provide more
realistic bounding values for the maximum concentration attainable in a soil-water system when doing
risk assessment calculations. Modeling computations can also be used to examine initial and final
geochemical conditions associated with laboratory Kd measurements to determine if the measurements
were affected by processes such as mineral precipitation, which might have compromised the derived
Kd values.
Although Kd values cannot be predicted a priori with chemical reaction models, modeling results can
provide aqueous speciation and solubility information that is exceedingly valuable in the evaluation of
Kd values selected from the literature and/or measured in the laboratory. Moreover, some models
include electrostatic adsorption submodels that may be used to estimate the changes in the composition
of the aqueous phase due to adsorption onto a selected mineral phase such as hematite or amorphous
iron oxyhydroxide. These results in turn can be used to back calculate a Kd value.
The MINTEQA21 computer code is an equilibrium chemical reaction model. It was developed with
EPA funding by originally combining the mathematical structure of the MINEQL code with the
thermodynamic database and geochemical attributes of the WATEQ3 code. The MINTEQA2 code
includes submodels to calculate aqueous speciation/complexation, oxidation-reduction, gas-phase
equilibria, solubility and saturation state (i.e., saturation index), precipitation/dissolution of solid phases,
and adsorption. The most current version of MINTEQA2 available from EPA is compiled to execute
on a personal computer (PC) using the MS-DOS computer operating system. The MINTEQA2
software package includes PRODEFA2, a computer code used to create and modify input files for
MINTEQA2.
Since the publication of Volumes I and II, the EPA has released a new version, Version 4.0, of the MINTEQ
software package (EPA, 1999a). The software package and documentation are available free on the EPA Internet site:
http://www.epa.gov/ceampubl/softwdos.htm
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The MINTEQA2 code contains an extensive thermodynamic database for modeling the speciation and
solubility of contaminants and geologically significant constituents in low-temperature, soil-water
systems. Of the contaminants selected for consideration in this project [americium, arsenic, cadmium,
cesium, chromium, curium, iodine, lead, neptunium, plutonium, radon, radium, strontium, technetium,
thorium, tritium (3H), and uranium], the MINTEQA2 thermodynamic database contains speciation and
solubility reactions for arsenic, including the valence states As (III) and As(V); chromium, including the
valence states Cr(II), Cr(III), and Cr(VI); cadmium; iodide, lead; strontium; and uranium, including the
valence states U(III), U(IV), U(V), and U(VI).
Individual users can supplement and/or update the MINTEQA2 database using thermodynamic
constants listed in published sources. Particularly noteworthy are several extensive, critical reviews of
thermodynamic data for radionuclides completed since 1992 by the Nuclear Energy Agency (NEA)
Thermodynamic Database Project of the Organisation for Economic Co-operation and Development
(OECD).1 These reviews were conducted by international teams of experts, peer reviewed prior to
publication, and commercially published in hard cover. These excellent sources include the reviews of
thermodynamic data for americium (Silva eta!., 1995), neptunium (Lemire eta!., 2001), plutonium
(Lemire eta!., 2001), technetium (Hard eta!., 1999) and uranium (Grenthe eta!., 1992).2 The
thermodynamic data in Grenthe eta!. (1992) supersede the uranium thermodynamic database currently
available with MINTEQA2.
The MINTEQA2 code includes seven adsorption model options. The non-electrostatic adsorption
models include the activity Kdct, activity Langmuir, activity Freundlich, and ion exchange models. The
electrostatic adsorption models include the diffuse layer, constant capacitance, and triple layer models.
The MINTEQA2 code does not include an integrated database of adsorption constants and reactions
for any of the seven models. These data must be supplied by the user as part of the input file
information.
Chemical reaction models, such as the MINTEQA2 code, cannot be used a priori to predict a partition
coefficient, Kd, value. The MINTEQA2 code may be used to calculate the chemical changes that result
in the aqueous phase from adsorption using the more data intensive, electrostatic adsorption models.
The results of such calculations in turn can be used to back calculate a Kd value. However, the user
must make assumptions concerning the composition and mass of the dominant sorptive substrate, and
supply the adsorption parameters for surface-complexation constants for the contaminants of interest
and the assumed sorptive phase. The EPA (EPA, 1992, 1996) has used the MINTEQA2 model and
this approach to estimate Kd values for several metals under a variety of geochemical conditions and
metal concentrations to support various waste disposal issues. The EPA in its "Soil Screening
Guidance" determined MINTEQA2-estimated Kd values for barium, beryllium, cadmium, Cr(III),
The Nuclear Energy Agency (NEA) is a specialized agency within the Organization for Economic Co-operation and
Development (OECD), which is an intergovernmental organization of industrialized countries, based in Paris, France.
Funding for the OECD/NEA Thermodynamic Database Project was provided from 17 organizations in 12 countries,
including the U.S. Department of Energy (DOE). More information regarding the OECD/NEA Thermodynamic
Database Project is available on the Internet at: http://www.nea.fr/html/dbtdb/cgi-bin/tdbdocproc.cgi
The reviews of thermodynamic data by Rard et al. (1999) and Lemire et al. (2001) were published in hard cover and
on companion compact disks (CDs).
4.2
-------
Hg(II), nickel, silver, and zinc as a function of pH assuming adsorption on a fixed mass of iron oxide
(EPA, 1996; RTI, 1994). The calculations assumed equilibrium conditions, and did not consider redox
potential or metal competition for the adsorption sites. In addition to these constraints, EPA (1996)
noted that this approach was limited by the potential sorbent surfaces that could be considered and
availability of thermodynamic data. Their calculations were limited to metal adsorption on iron oxide,
although sorption of these metals to other minerals, such as clays and carbonates, is well known.
Typically, the data required to derive the values of adsorption parameters that are needed as input for
adsorption submodels in chemical reaction codes are more extensive than information reported in a
typical laboratory batch Kd study. If the appropriate data are reported, it is likely that a user could hand
calculate a composition-based Kd value from the data reported in the adsorption study without the
need of a chemical reaction model.
4.3
-------
5.0 Contaminant Geochemistry and Kd Values
The important geochemical factors affecting the sorption1 of americium (Am), arsenic (As), curium
(Cm), iodine (I), neptunium (Np), radium (Ra), and technetium (Tc) are discussed in this chapter. The
objectives of this chapter are to: (1) provide a "thumb-nail sketch" of the key geochemical processes
affecting sorption of these contaminants, (2) provide references to key experimental and review articles
for further reading, (3) identify the important aqueous- and solid-phase parameters controlling
contaminant sorption in the subsurface environment, and (4) discuss the availability of sorption data
and Kd values for each contaminant. Unlike the contaminants reviewed in Volume II (EPA, 1999c),
the availability of Kd values and/or our understanding of the adsorptive behavior as a function of key
geochemical factors is more restricted for each contaminant reviewed in Volume III. These limitations
precluded selection of minimum and maximum conservative Kd values as a function of key
geochemical factors as was done for the contaminants in Volume II.
5.1 General
Important chemical speciation, (co)precipitation/dissolution, and adsorption/desorption processes for
each contaminant are discussed. Emphasis of these discussions is directed at describing the general
geochemistry that occurs in oxic environments containing low concentrations of organic carbon
located far from a point source (i.e., in the far field). These environmental conditions comprise a large
portion of the contaminated sites of concern to the EPA, DOE, and/or NRC. We found it necessary
to focus on the far-field, as opposed to near-field, geochemical processes for two main reasons. First,
the near field frequently contains very high concentrations of salts, acids, bases, and/or contaminants
which often require unusual chemical or geochemical considerations that are quite different from those
in the far field. Secondly, the differences in chemistry among various near-field environments varies
greatly, further compromising the value of a generalized discussion. Some qualitative discussion of the
effect of high salt conditions and anoxic conditions are presented for contaminants whose sorption
behavior is profoundly affected by these conditions.
The distribution of aqueous species for each contaminant was calculated for an oxidizing environment
containing the water composition listed in Table 5.1 and the chemical equilibria code MINTEQA2
(Version 3.10, Allison et a/., 1991). The water composition in Table 5.1 is based on a "mean
composition of river water of the world" estimated by Hem (1985). We use this chemical composition
simply as a proxy for the composition of a shallow groundwater. Obviously, there are significant
differences between surface waters and groundwater, and considerable variability in the concentrations
of various constituents in surface and groundwater. For example, the concentrations of dissolved gases
and complexing ligands, such as carbonate, may be less in a groundwater as a result of infiltration of
surface water through the soil column. Additionally, the redox potential of groundwater, especially
deep groundwater, will likely be more reducing than surface water. As explained later in this chapter,
the adsorption and solubility of certain contaminants and radionuclides may be significantly different
under reducing groundwater conditions compared to oxidizing conditions. However, as explained in
When a contaminant is associated with a solid phase, it is commonly not known if the contaminant is adsorbed
onto the surface of the solid, absorbed into the structure of the solid, precipitated as a three-dimensional molecular
coating on the surface of the solid, or absorbed into organic matter. "Sorption" will be used in this report as a generic
term devoid of mechanism to describe the partitioning of aqueous phase constituents to a solid phase. Sorption is
frequently quantified by the partition (or distribution) coefficient, K^.
5.1
-------
the Foreword, it was necessary to limit the scope of this review to oxidizing conditions. Use of the
water composition in Table 5.1 does not invalidate the aqueous speciation calculations discussed later
in this chapter relative to the behavior of the selected contaminants in oxidizing and transitional
groundwater systems. The calculations demonstrate what complexes might exist for a given
contaminant in any oxidizing water as a function of pH and the specified concentrations of each
inorganic ligand. If the concentration of a complexing ligand, such as phosphate, is less for a site-
specific groundwater compared to that used for our calculations, then aqueous complexes containing
that contaminant and ligand may be less important for that water. Importantly, the water composition
in Table 5.1 has a low ionic strength and contains no natural (e.g., humic or fulvic acids1) or
anthropogenic (e.g., EDTA) organic ligands.
Throughout this chapter, particular attention will be directed at identifying the important aqueous- and
solid-phase parameters controlling retardation2 of contaminants by sorption in soil. This information
was used to guide the review and discussion of published Kd values according to the important
chemical, physical, and mineralogical characteristics or variables. Perhaps more importantly, we chose
parameters that are readily available. For instance, particle size and pH data are often available, whereas
such parameters as abundance of iron oxide or surface area are not as frequently available.
"Humic and fulvic acids are breakdown products of cellulose from vascular plants. Humic acids are defined as the
alkaline-soluble portion of the organic material (humus) which precipitates from solution at low pH and are generally of
high molecular weight. Fulvic acids are the alkaline-soluble portion which remains in solution at low pH and is of lower
molecular weight" (Gascoyne, 1982).
Retarded or attenuated (i.e., nonconservative) transport means that the contaminant moves slower than water
through geologic material. Nonretarded or nonattenuated (i.e., conservative) transport means that the contaminant
moves at the same rate as water.
5.2
-------
Table 5.1. Estimated mean composition of river
water of the world from Hem (1985).
Dissolved Constituent
Silica, as H4SiO4
Ca
Mg
Na
K
Inorganic Carbon, as CO3
SO4
Cl
F
NO3
P04
Total Concentration1
mg/1
20.8
15
4.1
6.3
2.3
57
11
7.8
1
1
0.0767
mol/1
2.16 xlQ-4
3.7 x 10-4
1.7 xlO"4
2.7 x 10~4
5.9xl05
9.5 xlO"4
l.lxlO4
2.2 x 10~4
5 x 10~5
2 x 10~5
8.08 x 10~7
1 Most values from this table were taken from Hem (1985: Table 3,
column 3). Mean concentrations of total dissolved fluoride and
phosphate are not listed in Hem (1985, Table 3). The concentration
of dissolved fluoride was taken from Hem (1985, p. 120) who states
that the concentration of total dissolved fluoride is generally less than
1.0 mg/1 for most natural waters. Hem (1985, p. 128) lists 25 |ag/l
for average concentration of total dissolved phosphorous in river
water estimated by Meybeck (1982). This concentration of total
phosphorus was converted to total phosphate (PO4) listed above.
5.2 Americium Geochemistry and Kd Values
5.2.1 Overview: Important Aqueous- and Solid-Phase Parameters Controlling
Retardation
Americium is a transuranic (actinide) element, and can exist in the +3 oxidation state in natural waters.
Over the pH range of most natural waters, dissolved americium III (Am(III)) is present primarily as the
uncomplexed cation Am3+ in moderately to highly acidic conditions, and aqueous americium carbonate
complexes in near neutral to alkaline pH conditions. Americium readily sorbs to soil, mineral, and
crushed rock materials, and exhibits high Kd values. Americium is therefore considered to be immobile
in soil environments. However, the tendency of americium to strongly adsorb to soil particles indicates
that there is potential for colloid-facilitated transport of americium. Concentrations of dissolved
americium may be controlled by precipitation of hydroxide or carbonate solids in some systems.
5.3
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Therefore, some sorption measurements resulting in very high Kd values may reflect americium
precipitation reactions.
5.2.2 General Geochemistry
Americium [Am, atomic number (Z) = 95] has 19 isotopes (Tuli, 2000). The atomic masses of these
isotopes range from 231 to 249. Most americium isotopes have short half lives of minutes or less, and
only three americium isotopes have half lives longer than a few days. The americium isotopes present
in radioactive wastes include the radionuclides Am-241 (t,/2 = 432.2 years) and Am-243 (t,/2 =
7,370 years) (Onishi eta!., 1981). Americium found in the environment occurs primarily from decay of
Pu-241 (t,/2 = 14.29 years) in nuclear fallout and waste steams from nuclear fuel reprocessing plants.
Americium can exist in the +3, +4, +5, and +6 valence states. The +3 state is the most stable valence
state and the one most important in environmental systems. The higher oxidation states are strong
oxidizing agents, and stable only in systems containing no oxidizable compounds (Ames and Rai, 1978).
The environmental behavior of americium has been reviewed by Silva and Nitsche (1995), Coughtrey
etd. (1984), Onishi etaL (1981), and Ames and Rai (1978).
5.2.3 Aqueous Speclatlon
Silva eta!. (1995) have published an extensive, detailed review of the chemical thermodynamics of
americium aqueous species and solids. Moulin eta!. (1988, 1992) review the aqueous speciation of
Am(III) in natural waters and in the presence of humic substances in natural waters, respectively.
Studies indicate that Am(III) may form strong complexes with humic substances. Possible inorganic
aqueous species of Am(III) include those listed in Table 5.2.
Table 5.2. Americium(lll) aqueous species.
Aqueous Species
Am3+, AmOH2+< Am(OH)+, Am(OH)°3 (aq)
AmCO+, Am(CO3)2, Am(CO3)f
AmH2POf • Am(H2PO4)2, Am(H2PO4)°3 (aq), Am(H2PO4)4,
AmSO;, Am(SO4)2
AmCl2+' AmCl2, AmF2+< AmF+, AmF°3 (aq),
AmNO^'Am(NO3)2
The distribution of aqueous species for Am(III) (see Figure 5.1) as a function of pH was calculated for
an oxidizing environment containing the water composition listed in Table 5.1. The speciation
calculations indicate that the uncomplexed ion Am3+ is the dominant aqueous species for moderately to
highly acidic conditions. At near neutral to alkaline pH conditions, Am(III) carbonate complexes will
dominate the aqueous speciation of Am(III). Aqueous complexes, such as Am(CO3)f, will be
increasingly important with increasing concentrations of dissolved carbonate at these pH conditions.
5.4
-------
At highly alkaline pH values, Am(III) hydroxyl complexes, such as Am(OH)3 (aq), may become more
important than the Am(III) carbonate species.
5.2.4 Dissolution/Precipitation/Coprecipitation
Concentrations of dissolved Am(III) in soil environments may be controlled by the precipitation of
solids such as Am(OH)3 and AmOHCO3, and Am2(CO3)3, especially at near neutral and alkaline pH
conditions (Felmy et a/., 1990; Vitorge, 1992; Silva, 1984; and others). With increasing pH and
dissolved carbonate concentrations, AmOHCO3, will be the likely solubility control for dissolved
americium. Viorge (1992) used thermodynamic calculations to predict the stability domains of these
americium solids as a function of pH and dissolved carbonate.
Figure 5.1. Calculated aqueous speciation for Am(lll) as
a function of pH. [Americium(lll) aqueous speciation
was calculated based on a total dissolved concentration
of americium of 1x10'12 mol/l (2.43x10'7 mg/l),
and the water composition in Table 5.1.]
100
Am(OH)3u (aq)
AmOH2+
6
7
8
10
PH
5.2.5 Adsorption/Desorption
5.2.5.1 Guidance for Screening Calculations of .Adsorption
Most sorption studies indicate that Am(III) readily sorbs to minerals, crushed rock, and soil materials,
and is therefore considered one of the most immobile actinide elements in the environment.
Americium adsorption studies published prior to 1984 have been reviewed by Coughtrey et al. (1984),
Onishi et al. (1981), and Ames and Rai (1978). Some of these early studies and more recent studies of
Am(III) adsorption are summarized in the sections below. Americium(III) exhibits large Kd values
often in the range of 1,000 to >100,000 ml/g. However, the concentrations of dissolved americium
may be controlled in some geochemical systems by precipitation of hydroxide or carbonate solids. The
5.5
-------
reader should therefore be cautious because some sorption measurements resulting in very high Kd
values may have been affected by americium precipitation reactions.
A limited number of Kd studies was identified during this review for the adsorption of Am(III) on soil
as a function of key geochemical parameters such as pH. The limited availability of such values
precluded development of Kd look-up tables of conservative minimum and maximum Kd values for
Am(III). Of the studies reviewed below, only Routson et al. (1975, 1977) and Sanchez et al. (1982)
report Kd values and corresponding pH values for Am(III) adsorption on soil (Figure 5.2). For the pH
range from 4 to 10, it is suggested that a Kd of 4 ml/g be used as a minimum Kd value for screening
calculations of americium transport in soils. This value was reported for pH 7.8 by Routson et al.
(1975, 1977) and is the lowest Kd value that they gave for experiments conducted with very to
moderately dilute calcium and sodium electrolyte solutions. The other Kd values reported by Routson
et al. (1975, 1977) for these solution concentrations ranged from 6 ml/g at pH 6.2 to 1,200 ml/g at
pH 4.1 and 7.4.
Figure 5.2. ^ values (ml/g) for Am(lll) adsorption on soil
reported by Routson et al. (1975, 1977) (solid squares)
and Sanchez et al. (1982) (solid triangles).
250,000
200,000
=? 150,000
100,000
<
50,000
A
A A
AA
AA
AA
L_
6
7
PH
10
The solid line segments in Figure 5.2 connect the maximum Kd values reported at pH values of 4, 6,
and 10 by Sanchez et al. (1982). The Kd values corresponding to integer pH values between 4 to 10 are,
respectively, 5,600, 16,500, 27, 300, 76,700, 126,000, 176,000, and 225,000 ml/g based on straight line
extrapolations between these 3 Kd values from Sanchez et al. (1982). The reader may want to consider
these values as conservative maximum Kd values for Am(III) adsorption on soil.
A large uncertainty however is associated with these roughly estimated, conservative minimum and
5.6
-------
maximum Kd values due the limited number of Kd adsorption studies for Am(III) on soil. Accordingly,
one of the major recommendations of this report is that for site-specific calculations, partition
coefficient values measured at site-specific conditions are absolutely essential.
Based on studies of Am(III) on single mineral phases and crushed rock, the adsorption of Am(III) is
strongly pH dependent, and increases with increasing pH with peak adsorption occurring between pH
values of 5 and 6 (see studies discussed below). This observed pH dependence is expected, because the
dominant aqueous species of americium in the pH range of natural waters are primarily Am3+ and
cationic carbonate complexes at acidic and basic pH values, respectively (see Figure 5.1).
Americium(III) is more mobile at low to moderate pH values where the net surface charge on minerals
becomes more positive and in high ionic strength solutions. Adsorption of Am(III) might decrease in
the pH range from approximately 8.5 to 10 due to the dominance of the anionic complex Am(CO3)2
(Figure 5.1).
5.2.5.2 General .Adsorption Studies
Sheppard et al. (1979) used batch equilibration experiments to study the sorption of americium and
curium to colloidal-size soil particles that are potentially diffusible in soil/water systems. The
experiments were conducted with distilled water and 14 soils from Muscatine, Illinois; Hanford,
Washington; Barnwell, South Carolina; Idaho Falls, Idaho; and Paradise and Placerville, California.
Centrifugation measurements indicated that much of the Am-241 was retained by the colloidal-size soil
particles. The sorption of Am-241 on the soil particles was not complete after 4 to 6 months, and
percentage of radionuclide retained by the colloidal-size fraction decreased with time. Sheppard et al.
(1979) found it difficult to correlate the sorption results with chemical and physical properties of the
soils. They attributed this to the lack of precise distribution ratios, competition with cation exchange
reactions, and complexation with humic and fulvic acid materials. The results also indicated that the
sorption behavior of Am-241 and curium (Cm-244) under these experimental conditions was identical.
Sheppard et al. (1979) suggested that colloids of clay and humic acids are potentially important
processes for the transport of actinides in soil/water systems.
Other studies have shown the potential importance of colloid-facilitated transport of americium in soil
systems. Laboratory studies by Penrose et al. (1990) predicted that the movement of americium and
plutonium would be limited to less than a few meters through a shallow aquifer within the site of the
Los Alamos National Laboratory (LANL) which is in a semiarid region. However, both actinides were
detected in monitoring wells as far as 3,390 m down gradient from the point source. Almost all of the
americium and plutonium in the groundwater at the 3,390 m well were associated with colloids 0.025 to
0.45 |^m in diameter. Similarly, the results of laboratory measurements using site-specific soils and a
two-phase solute transport code indicated that americium, curium, plutonium, and uranium would
migrate less than 10 m in the F-Area of the Savannah River Site (Kaplan et al., 1994). The
contaminants however were found associated with groundwater colloids 1,200 m away from the point
source. Colloid-facilitated migration of contaminants is reviewed in Section 2.7 in Volume I (EPA,
1999b). The reader is cautioned that importance of colloid-facilitated migration, especially in aquifer
systems that do not involve fracture flow of groundwater, is still the subject of debate.
Means et al. (1978) studied the mineralogy of the adsorbents for americium in soil near a disposal trench
at the nuclear waste burial grounds at the Oak Ridge National Laboratory in Oak Ridge, Tennessee.
5.7
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Their analyses indicated that americium concentrations were highly correlated to the concentrations of
manganese oxides, even though the concentrations of iron oxyhydroxides and organic carbon were
significantly greater than those of manganese oxides. Means et al. (1978) expected americium to be
adsorbed by iron oxyhydroxides, but their correlation coefficients for americium sorption versus total
iron oxyhydroxides and organic carbon matter were low.
5.2.5.3 K., Studies for Ameridum on Soil Materials
Carroll et al. (1999) investigated the sorption of radionuclides in waters and sediments from the Arctic
Sea. Field and laboratory Kd batch measurements were made for the sorption of americium using
freshly collected sediment and water samples from stations in the Novaya Zemlya Trough and on the
continental shelf of the Ob/Yenisey Rivers. The Kj values measured for americium ranged from 7xl03
to I.lxl06ml/g.
Nakayama and Nelson (1988) measured the pH dependency of Kd values for the sorption of 243Am and
CM-244 onto sediment. They found that changes in pH had a significant effect on the adsorption of
these 2 radionuclides. However, the relative adsorption behaviors of these two radionuclides to each
other were not significantly different as a function of pH.
Table 5.3. Measured ^ values (ml/g) for americium as a
function of pH for Hudson River
estuary environment [Sanchez et al. (1982)].
pH
4.0
6.0
7.0
8.0
10.0
Station Location at Hudson River Estuary for Sediment Samples
MPO.l
0.43xl04
2.91xl04
8.24xl04
13.7xl04
18.9xl04
MP 18.6
O.llxlO4
1.77xl04
4.5xl04
21.0xl04
22.5xl04
MP 43.3
0.43xl04
2.67xl04
6.12xl04
10.43xl04
16.7xl04
MP 59.8
0.56xl04
2.73xl04
6.5xl04
6.8xl04
15.3xl04
Burton et al. (1986) studied the sorption of americium on intertidal sediment from the Ravenglass
Estuary in seawater from the North Sea. The Kd values were of the magnitude 103-104 ml/g. Results
from desorption experiments conducted with seawater diluted with river water indicated that the Kj
values were essentially constant at salinities from 34 to 2-5%, and decreased by 2 orders of magnitude
at salinities less than 2-5%. Burton et al. (1986) attributed this decrease in Kj in part to a decrease in
pH.
Sanchez et al. (1982) used the batch method to measure Kd values for americium for a variety of
freshwater, estuarine, and marine environments. The adsorption Kd values ranged from 6.76xl04 to
2.27xl06 ml/g for all environments studied. Experiments conducted as a function of pH (Table 5.3.)
5.8
-------
indicated that the Kd values increased with increasing pH. The americium Kd values increased by a
factor of 5 over the pH range 4 to 6, and continued to increase at high pH values. Sanchez et al. (1982)
found no apparent effect of salinity on the americium Kd values.
Table 5.4. Measured americium Kd values (ml/1) and soil
properties for soils studied by Nishita et al. (1981).
Soil
Type
Silt Clay Loam
Sharp sburg
Sandy Loam
Malbis
Sandy Loam
Lyman
Silty Clay
Holtsville
Loam
Aiken
Silt Loam
Yolo
Muck
Egbert
pH1
5.9
5.3
5.0
7.8
6.0
6.7
7.2
OM2
Fraction
(%)
2.8
2.4
5.7
0.6
8.4
2.5
40.8
CEC3
(meq/100
K)
20
15
15
30
15
25
60
Free Fe
Oxides
(%)
1.29
1.65
1.52
1.20
5.29
2.41
1.57
Mn4
Fraction
(%)
0.06
0.05
0.04
0.04
0.10
0.08
0.10
Extract
pH
5.41
6.56
4.39
5.71
4.58
6.17
7.12
8.04
5.71
6.72
6.12
6.98
7.14
7.54
Kd
(ml/g)
29,800
17,280
9,635
8,063
1,549
182
35,630
47,230
21,870
10,660
23,870
20,210
7,266
5,529
1 Saturated paste. 4 In 4 M HNO3 extract.
2 Organic matter.
3 Cation Exchange Capacity
Nishita et al. (1981) studied the extractability of Am-241 from several types of soils as a function of pH.
The extractability of americium was considered to parallel the solubilization of aluminum, iron, and/or
manganese hydrous oxides, which are important adsorbents for dissolved contaminants in soil systems
(Nishita et al., 1981). The americium Kd values and soil characteristics determined for these soils by
Nishita et al. (1981) are listed in Table 5.4. At pH values more acidic than those studied, Nishita et al.
(1981) suggested that americium would be present in ionic form and highly mobile. At higher pH
values, americium formed hydroxyl complexes that were readily sorbed by the soils.
Routson et al. (1975, 1977) used batch equilibration experiments to measure Kd values for Am-241 on
2 soils as a function of the concentrations of dissolved calcium and sodium. The soil samples were
selected to represent a range of weathering intensities. For arid conditions in the western United States,
a sandy (coarse-textured), low-exchange capacity soil was selected from a low rainfall area in eastern
Washington. For humid conditions in southeastern United States, a moderate-exchange capacity soil
5.9
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was selected from South Carolina. Properties of the soils used for these measurements are listed in
Table 5.5.
Table 5.5. Properties of soils used in h^ measurements
by Routson et al. (1975, 1977).
Soil
Washington
South Carolina
CaCO3
(mg/g)
0.8
<0.2
Silt
(%)
10.1
3.6
Clay
(%)
0.5
37.2
CEC
(meq/100 g)
4.9
2.5
pH
7.0
5.1
The Kd values were measured in 0.002, 0.02, 0.05, 0.10, and 0.20 M Ca(NO3)2 solutions, and 0.015,
0.030, 0.30, 0.75, and 3.0 M NaNO3 solutions. The pH values for selected samples of the Am-241
solutions in the calcium and sodium systems were 6.9 and 4.1 for the Washington soil, and 7.1 and 6.1
for the South Carolina soil. The Kd values decreased with increasing concentrations of dissolved
calcium and sodium. For the solution concentrations used in these experiments, the Kd values for Am-
241 on the South Carolina soil ranged from 67 to 1.0 ml/g as a function of dissolved calcium, and 280
to 1.6 ml/g as a function of dissolved sodium. For the Washington soil, the Kd values were
> 1,200 ml/g, and were independent of the concentrations of dissolved calcium and sodium. Their
calculated Kd values ranged from 1,200 to 8,700 ml/g. The sorption of Am-241 on the Washington
soil was greater than that expected by investigators.
5.2.5.4- Published Compilations Containing Kj Values for Amendum
Because the references in this section are often cited or used for comparison in other publications, the
following summaries are provided for completeness. It is recommended that the reader review the
original reference and the references cited therein to understand the procedures and sources of the Kd
values used for each compilation. The compilations do not typically consider important factors that
contribute to variability in sorption, such as pH. Moreover, in cases where very large Kd values are
listed, there is a risk that the original Kd measurement may have included precipitated components.
Baes and Sharp (1983) present a simple model developed for order-of-magnitude estimates for leaching
constants for solutes in agricultural soils. As part of this model development, they reviewed and
determined generic default values for input parameters, such as Kd. A literature review was completed
to evaluate appropriate distributions for Kd values for various solutes, including americium. Because
Baes and Sharp (1983) are cited frequently as a source of Kd values in other published Kd reviews (e.g.,
Looney et al., 1987, Sheppard and Thibault, 1990), the americium Kd values listed by Baes and Sharp are
reported here for completeness. Based on the distribution that Baes and Sharp determined for the Kd
values for cesium and strontium, they assumed a lognormal distribution for the Kd values for all other
elements in their compilation. Baes and Sharp listed an estimated default Kd of 810 ml/g for
americium based on 46 Kd values that ranged from 1.0 to 47,230 ml/g for agricultural soils and clays
5.10
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over the pH range 4.5 to 9.0. Their compiled Kd values represent a diversity of soils, pure clays (other
Kd values for pure minerals were excluded), extracting solutions, measurement techniques, and
experimental error.
Looney et al. (1987) tabulated estimates for geochemical parameters needed for environmental
assessments of waste sites at DOE's Savannah River Plant in South Carolina. Looney et al. list Kd
values for several metal and radionuclide contaminants based on values that they found in one to five
published sources. For americium, Looney et al. list a "recommended" Kd of 100 ml/g, and a range
from 1 to 100,000 ml/g. Looney et al. note that their recommended values are specific to the Savannah
River Plant site, and they must be carefully reviewed and evaluated prior to use in assessments at other
sites.
Table 5.6. Americium Kd values (ml/g) listed by
Thibault ei al. (1990, Tables 4 to 8).
Soil Type
Sand
Silt
Clay
Organic
Kd Values (ml/g)
Geometric
Mean
1,900
9,600
8,400
112,000
Number of
Observations
29
20
11
5
Range
8.2 - 300,000
400 - 48,309
25 - 400,000
6,398 - 450,000
Thibault et al. (1990) (also see Sheppard and Thibault, 1990) present a compilation of soil Kd values
prepared to support radionuclide migration assessments for a Canadian geologic repository for spent
nuclear fuel in Precambrian Shield plutonic rock. Thibault et al. collected Kd values from other
compilations, journal articles, and government laboratory reports for important elements, such as
americium, that would be present in the nuclear fuel waste inventory. The americium Kd values listed
in Thibault et al. (1990) are included in Table 5.6. Thibault et al. (1990) describe the statistical methods
used for analysis of the compiled Kd values. The range for the Kd values used to calculate the
"geometric mean" cover several orders of magnitude. Readers are cautioned against using "geometric
mean" values or any other form of averaged Kd values as "default" Kd values, as such values are usually
calculated from data compiled from different investigators for different soil systems. These mean or
average values do not represent any particular environmental system and geochemical conditions. As
discussed in Volume I (EPA, 1999b), the variation observed in the literature for Kd values for a
contaminant is due to differences in sorption mechanisms, geochemical conditions, soil materials, and
methods used for the measurements.
McKinley and Scholtis (1993) compare radionuclide Kd sorption databases used by different
international organizations for performance assessments of repositories for radioactive wastes. The
americium Kd values listed in McKinley and Scholtis (1993, Tables 1, 2, and 4) are listed in Table 5.7.
5.11
-------
The reader should refer to sources cited in McKinley and Scholtis (1993) for details regarding their
source, derivation, and measurement. Radionuclide Kd values listed for cementitious environments in
McKinley and Scholtis (1993, Table 3) are not included in Table 5.7. The organizations listed in the
tables include: AECL (Atomic Energy of Canada Limited); GSF (Gesellschaft fur Strahlen- und
Umweltforschung m.b.H., Germany); IAEA (International Atomic Energy Agency, Austria); KBS
(Swedish Nuclear Safety Board); NAGRA [Nationale Genossenschaft fur die Lagerung radioaktiver
Abfalle (Swiss National Cooperation for Storage of Radioactive Waste), Switzerland]; NIREX (United
Kingdom Nirex Ltd.); NRC (U.S. Nuclear Regulatory Commission); NRPB (National Radiological
Protection Board, United Kingdom); PAGIS [Performance Assessment of Geological Isolation
Systems, Commission of the European Communities (CEC), Belgium; as well as PAGIS SAFIR
(Safety Assessment and Feasibility Interim Report]; PSE (Projekt Sicherheitsstudien Entsorgung,
Germany); RIVM [Rijksinstituut voor Volksgezondheid en Milieuhygience (National Institute of Public
Health and Environment Protection), Netherlands]; SKI [Statens Karnkraftinspektion (Swedish
Nuclear Power Inspectorate)]; TVO [Teollisuuden Voima Oy (Industrial Power Company), Finland];
and UK DoE (United Kingdom Department of the Environment).
5.12
-------
Table 5.7. Americium Kd values (ml/g) listed by McKinley and
Scholtis (1993, Tables 1, 2, and 4) from sorption databases
used by different international organizations for performance
assessments of repositories for radioactive wastes.
Organization
AECL
GSF
IAEA
KBS-3
NAGRA
NIREX
NRC
NRPB
PAGIS
PAGIS SAFIR
PSE
RIVM
SKI
TVO
UK DoE
Argillaceous (Clay)
Sorbing
Material
Bentonite-Sand
Sediment
Pelagic Clay
Bentonite
Bentonite
Clay
Clay Muds tone
Clay, Soil Shale
Clay
Bentonite
Clay
Sediment
Sandy Clay
Bentonite
Bentonite
Lake Sediment
Coastal Marine
Water
Kd
(ml/g)
300
2,000,000
2,000,000
29,400
5,000
70
5,000
10,000
80
2,000
600
10,000
1,000
7,000
2,900
1,000
100,000
1,000,000
Crystalline Rock
Sorbing
Material
Granite
Granite
Granite
Granite
Basalt
Tuff
Granite
Crystalline
Rock, Reducing
Crystalline Rock
Kd
(ml/g)
1,000
5,000
5,000
300
50
100
5,000
3,000
400
Soil/Surface Sediment
Sorbing
Material
Soil/ Sediment
Soil/ Sediment
Soil/ Sediment
Soil/ Sediment
Soil/ Sediment
Soil/ Sediment
Kd
(ml/g)
5,000
3,000
8,800
100,000
800
5.13
-------
5.2.5.5 Kj Studies ofA^meridum on Pure Mineral, Oxide, and Crushed Rock Materials
Numerous adsorption studies have been conducted of americium on pure minerals, oxide phases, and
other geologic-related materials. The Kd values listed in these studies are not necessarily relevant to
the mobility and sorption of americium in soils. However, they are listed in Appendix C for
completeness. The references cited in Appendix C are listed in the main reference list in Chapter 6.
The potential value of the references that they cite and the sorption processes that they discuss is left
to the reader. The studies of americium sorption on crushed rock were conducted typically as part of
national research programs to investigate the feasibility of geological disposal of high-level radioactive
waste (HLW).
5.3 Arsenic Geochemistry and Kd Values
5.3.1 Overview: Important Aqueous- and Solid-Phase Parameters Controlling
Retardation
Arsenic is a known carcinogen, and occurs in natural systems in the +5 (arsenate) and +3 (arsenite)
valence states. Arsenic(III) (As(III)) is more mobile (adsorbs less) and is many times more toxic than
As(V). The pH and redox conditions are the two most important geochemical factors affecting the
mobility of arsenic in the environment. Sorption studies indicate that the concentrations of dissolved
As(V) and As (III) are controlled by adsorption on iron and aluminum oxides and clays. The
adsorption of As(V) is high and independent of pH at acidic pH values, and decreases with increasing
pH in the range of 7 to 9. Arsenic(V) adsorption will be decreased in soils with high phosphate
concentrations because of anion competition. Iron-reducing bacteria may cause arsenic mobilization
from soils as a consequence of reductive dissolution of iron oxyhydroxide adsorbents. Sulfate-
reducing bacteria, in addition, may promote arsenic reduction by producing hydrogen sulfide.
5.3.2 General Geochemistry
Arsenic [As, atomic number (Z) = 33] is a known carcinogen. Arsenic contamination of groundwater
may result from a variety of sources, such as weathering of rocks, mining activities, discharges of
industrial waste, and application of arsenical herbicides and pesticides. Arsenic exhibits a complex
geochemical behavior. It can occur in the valence states -3, 0, +3, and +5 in natural systems.
Arsenic(V) (arsenate) and As (III) (arsenite) are the main valence states of arsenic in natural waters
under oxidizing and reducing conditions, respectively. Elemental (metal) arsenic [As(0)J occurs rarely,
and As (-III) exists only at extremely low redox potential (i.e., Eh values). Arsenic(III) is more mobile
(adsorbs less) and many times more toxic than As(V). Although As(V) should be the dominant
valence state in oxidizing waters based on thermodynamic considerations, the results of some studies
indicate that As (III) concentrations exceed those of As(V) in some surface waters (Korte and
Fernando, 1991). Arsenic(III) may be more prevalent than generally believed in such environments
due to oxidation/reduction disequilibrium and biotic processes.
It should be noted that only three published studies containing Kd values for arsenic sorption on soil
were identified during the course of this review. This is an important finding given the increasing
concern over arsenic contamination in the environment and associated risks to plants, animal, and
human health. We suspect that the limited availability of Kd values for arsenic may be related to the
Kd approach being used prior to the 1990's primarily in risk and site performance assessments
5.14
-------
associated with radionuclides and their disposal and potential migration in the subsurface
environment.
Figure 5.3. Eh-pH stability diagram for the dominant arsenic
aqueous species at 25°C. [Diagram based on a total
concentration of 10"6 mol/l dissolved arsenic.]
1
.5
'5T
"o
3x
.c
LJJ
0
-.5
(
~"\ H3As04° (aq)
*^^~\^
H2AsO4 ^-^
•"""""-- r~"~-~
^xx ^~^~
x~xx
Xx-vx HAsO42
H3As03°(aq) \
J
NN
""-^^ XX AS°43
"*^"*^ ^\^^ 1 x
N"""'v-» T^^ Xx^
""""""-» ! ! "^x^
~x~*. ' """~x
-^ 1 x^
MASU3 ~x^
Illlllllllll """t"^
) 2 4 6 8 10 12 1
PH
4
Arsenic contamination of drinking water is a major concern recognized in areas of the United States,
and internationally in many countries. Arsenic contamination of groundwater in Bangladesh has been
highly publicized, and has generally been attributed to a geologic origin. The following models
(WHO, 1999) have been proposed for how groundwater contamination occurs:
• Pyrite oxidative dissolution - Pumping of air or water containing dissolved oxygen into the
subsurface results in the dissolution of arsenic-containing pyrite, and mobilization of arsenic.
• Oxyhydroxide reductive dissolution - Iron oxyhydroxide particles containing sorbed arsenic
are exposed to a reducing geochemical environmental that results in the dissolution of the
iron oxyhydroxides and mobilization of the sorbed arsenic.
5.15
-------
Figure 5.4. Calculated aqueous speciation for As(V) as a
function of pH. [Arsenic(V) aqueous speciation was
calculated based on a total concentration of dissolved
arsenic of 6.7x10"7 mol/l (50 ppm) and the water
composition in Table 5.1.]
.g
.9
"w
b
1
CD
100
80
60
40
20
0
H2As04"
HaAsCV (aq)
HAs04
2-
As04
4567
PH
8
9 10
The mechanisms responsible for arsenic mobility in groundwater in Bangladesh and West Bengal
have discussed by many, including Nickson et al. (2001).
Various aspects of the behavior of arsenic in natural systems are reviewed by Sadiq (1997), Korte and
Fernando (1991), Cullen and Reimer (1989), Crecelms et al. (1986), Rai et al. (1984), Woolson (1983),
Braman (1983). Biotic processes are potentially of considerable significance to the environmental
chemistry of arsenic (see Section 5.3.3.1). The review by Cullen and Reimer (1989) is particularly
noteworthy because of its magnitude (51 pages), level of detail, and extensive reference list
(458 references cited). Reviews of arsenic geochemistry and behavior in fossil fuel combustion
residues (Rai et al., 1987) and geothermal systems (Ballantyne and Moore, 1988) have also be
published. Korte and Fernando (1991) review in detail the behavior of As (III) in groundwater.
Remediation of arsenic-contaminated environments has also been studied. Recently published
studies, for example, include arsenic immobilization in contaminated soil by the addition of ferrous
sulfate and water, followed by the addition of Ca(OH)2, Portland cement, and water (Voigt and
Brantley, 1996), and formation of solid calcium arsenate (Both and Brown, 1999); and in situ removal
of dissolved arsenate and arsenite in contaminated groundwater by introduction of metallic iron
fillings (zero valent iron) (Su and Puls, 2001; Lackovic et al., 2000).
5.16
-------
5.3.3 Aqueous Speciation
Sadiq and Lindsay (1981) reviewed the thermodynamic data for arsenic aqueous and solids species.
They concluded that the majority of thermodynamic constants for arsenic were not reliable. We have
not identified any compilation of arsenic thermodynamic constants more current than Sadiq and
Lindsay (1981).
The aqueous speciation of dissolved arsenic has been reviewed in detail by Sadiq (1997) and Cullen
and Reimer (1989). The stable forms of dissolved arsenic in natural waters are primarily a series of
aqueous As(V) (arsenate, AsO^) and As (III) (arsenite, AsOf) oxyanion species under oxidizing and
reducing conditions, respectively. Figure 5.3 is an Eh-pH diagram that shows the dominant aqueous
species of arsenic as a function of pH and Eh (redox potential) at 25°C with respect to the
thermodynamic stability of water.
The known aqueous species of As(V) include AsO^, HAsO^, H2AsO4, and H3AsO° (aq) (Wagman
et a/., 1982; Sadiq, 1997). The distribution of these As(V) aqueous species (Figure 5.4) as a function
of pH was calculated for an oxidizing environment containing the water composition listed in
Table 5.1. The calculations indicate that the oxyanions H2AsO4 and HAsOj" dominate the aqueous
speciation of As(V) at pH values less than and greater than approximately pH 7, respectively.
Because As(V) will be present primarily as anions in oxidizing waters, As(V) is expected to adsorb to
geologic materials under acidic conditions and not to adsorb to any significant extent at neutral and
alkaline pH conditions. Similarly, the dominate As (III) aqueous species are H3AsO° (aq) and H2AsO3
at pH values less than and greater than approximately pH 9, respectively.
Arsenic is not expected to form aqueous complexes with other inorganic complexing ligands, such as
dissolved sulfate, because dissolved arsenic is present in anionic form. However, the results of Kim et
al. (2000) indicate that carbonation reactions with arsenic sulfide minerals may be an important
process in leaching arsenic into groundwater under anaerobic conditions. Kim et al. (2000)
investigated the role of bicarbonate in leaching arsenic from a sandstone into groundwater. They
determined that the release of arsenic from arsenic sulfides in the sandstone was strongly related to
the bicarbonate concentrations in the leachate. Kim et al. (2000) proposed that the released As (III)
was converted to As (III) carbonate complexes with the presumed compositions As(CO3)2,
As(CO3)(OH)2, and/or AsCO3. Kim et al. (2000) expected that once formed, the As (III) carbonate
complexes would be stable under anaerobic groundwater conditions at acidic to neutral pH. No
published stability constants are known to exist for such aqueous complexes.
Organoarsenic species, such as methyl arsonic acid (CH3AsO(OH)2) and dimethyl arsenic acid
[(CH3)2AsOOH], can be formed by biologically-mediated methylation in the environment, and have
been detected in soils and aquatic systems. Because their formation requires the involvement of
anaerobic bacteria, the literature refers to the origin of these organoarsenic compounds as
biomethylation. Methylated arsenic species typically comprise only a small fraction of the available
arsenic (Sanders, 1980). Thermodynamic data are not available for such compounds, and no stability
diagrams have been published regarding their range of stability (Cullen and Reimer, 1989). The
chemistry, formation, and behavior of these organoarsenic compounds are reviewed by Cullen and
Reimer (1989) and Sadiq (1997).
5.17
-------
5.3.4 Dissolution/Precipitation/Coprecipitation
Although it can be argued, concentrations of dissolved As(V) in natural waters and soils are not likely
controlled by formation of arsenic solids unless elevated concentrations of arsenic exist, such as in
mining waste waters, industrial waste streams, or fossil fuel combustion residues. The formation of
arsenic solids that would limit the solubility of arsenic in waste waters, natural waters, and soils has
been the subject of considerable study and conjecture. For example, see Lumsdon eta!. (2001), Davis
(2000), Tempel eta!. (2000), Rochette eta!. (1998); Sadiq (1997), Fruchter eta!. (1990), and Crecelms et
al. (1986). Searches of mineralogy databases available on the Internet indicate that there are more
than 490 known minerals that contain arsenic as a primary component. These minerals include
arsenic sulfides, arsenates, and arsenides, and are present in unique geologic settings such as ore
deposits and geothermal settings. Many attempts at identifying possible solubility controls for
dissolved As(V) and As (III) have relied on geochemical solubility calculations. Because
thermodynamic constants and kinetic data exist for a very limited number of arsenic solids (Sadiq,
1997), and even these data are not well established, arsenic thermodynamic-based solubility
calculations are problematic at best.
Several investigators have suggested Ba3(AsO4)2 as a solubility-limiting solid for As(V). However,
Crecelius et al. (1986) noted that there is no physical evidence for the existence of such a solid in
geologic systems and thus its thermodynamic properties are in error or kinetic constraints inhibit its
precipitation.
Under reducing conditions in the presence of dissolved sulfide, arsenic sulfide solids, such as
orpiment may control the solubility of dissolved As (III). Moreover, pyrite (FeS2) has been found to
contain a significant mass of coprecipitated (i.e., absorbed) arsenic in its crystal structure \_e.g., see
Schreiber et al. (2000)], and to be a source and/or solubility control for dissolved As (III).
5.3.5 Adsorption/Desorption
5.3.5.1 Guidance for Screening Calculations of Adsorption
Only three published studies containing Kd values for arsenic sorption on soil were identified during
the course of our review. These studies are discussed in the sections below. The studies by Cornett
et al. (1992) and Mok and Wai (1990) involve sampling environments affected by ore milling
operations and list extremely high Kd values (>1000's ml/g) for arsenic sorption on sediment. These
high Kd values are surprising given the large mobility expected for As(V) and As (III) based on the
results of other published sorption studies at all but the very acidic pH values. The results from
Cornett et al. (1992) and Mok and Wai (1990) are therefore suspect and were not considered further
in this review.
The limited availability of Kd values for arsenic on soil (i.e., 1 study) precluded calculation of Kd look-
up tables for arsenic as a function of important geochemical parameters such as pH. Because
dissolved As(V) will exist primarily as anionic aqueous species in the pH range of 4 tolO, a Kd value
of 0 ml/g is suggested as a conservative minimum value for site screening calculations of the
maximum extent of off-site migration of As(V) in soil. However, one of the major recommendations
of this report, as noted previously, is that for site-specific calculations, partition coefficient values
measured at site-specific conditions are absolutely essential.
5.18
-------
Given the limited availability of As(V) Kd values for soil, readers may want to consider using
geochemical models [see Section 5 in Volume I (EPA, 1999b)] to estimate the mass of adsorbed
As(V). This approach was recently taken by Lumsdon et al. (2001). Characterization studies indicated
that hydrous ferric oxide (HFO) was an important mineral in the soils at the site being studied.
Assuming that HFO was the dominant adsorbent for As(V) in these soils, Lumsdon et al. (2001) used
the generalized two-layer surface complexation model (GTLM) and As (III)-HFO surface
complexation parameters listed in Dzombak and Morel (1990) to estimate the mass of adsorbed
As(V). Based on the HFO content of the soils, Kj values could then be estimated from the modeling
results.
The results of sorption studies of arsenic on single mineral phases indicate that the concentrations of
dissolved As(V) and As (III) are controlled by adsorption on iron and aluminum oxides and clays.
Adsorption of arsenic exhibits a marked pH dependency. While adsorption of As(V) is high and
independent of pH at acidic pH values, adsorption decreases with increasing pH in the range of 7 to
9.
At a given pH, the adsorption of As(V) is typically greater than As (III). Arsenic(V) and dissolved
phosphate compete for adsorption sites, and the adsorption of arsenic will be reduced in soils with
high phosphate concentrations. Microbial-mediated reactions may affect arsenic mobilization in soils.
Studies have shown that iron-reducing bacteria may cause arsenic mobilization from soils as a
consequence of reductive dissolution of iron oxyhydroxide adsorbents. Microbial-mediated reactions
can cause the precipitation and dissolution of minerals, and thus affect the mobility of contaminants
in aqueous environments. The interactions between microbes and minerals are reviewed in detail
elsewhere, such as Banfield and Nealson (1997). Additionally, sulfate-reducing bacteria may promote
arsenic reduction by producing hydrogen sulfide.
5.3.5.2 General Adsorption Studies
Sorption studies conducted with soils indicate that arsenic sorption on soil is a function of the iron
oxyhydroxide and clay contents (Manning and Goldberg, 1997b; Wauchope and McDowell, 1984;
Livesey and Huang, 1981; Wauchope, 1975, Jacobs et al., 1970; and others). Manning and Goldberg
(1997) measured the adsorption of As(V) and As (III) on three arid-zone soils from California as a
function of varying arsenic concentrations, pH, and ionic strength. The greatest extent of As(V) and
As (III) adsorption was measured for the soil having the highest citrate-dithionite extractable iron and
percent clay content. Manning and Goldberg (1997) noted that the arsenic adsorption behavior for
this soil was similar to the arsenic adsorption behavior that they measured for goethite. The sorption
measurements also indicated that As(V) adsorbed more strongly than As (III) under most conditions.
Wauchope and McDowell (1984) measured the sorption of As(V) on 14 lake sediments collected in
the Delta Mississippi River flood plain. Wauchope and McDowell (1984) found that the extent of
arsenic adsorption was related to the hydrous oxide and clay contents of the sediments. Based on
their results, Wauchope and McDowell (1984) proposed that the clay content of sediment was the
best predictor for arsenic adsorption in soils and dilute water-sediment mixtures. Wauchope (1975)
used batch reaction experiments to study the sorption of As(V) on 16 soils from the Mississippi River
alluvial flood plain. Wauchope (1975) determined that the sorption of As(V) was strongly correlated
with the iron oxide and clay contents of the soils. Based on measurements of 24 soils representing a
range of chemical and physical properties from throughout Wisconsin, Jacobs et al. (1970) showed
that the mass of sorbed arsenic increased with an increase of the free Fe2O3 content of soil.
5.19
-------
Carrillo and Drever (1998) conducted a series of sorption experiments for As(V) and As (III) on soils
taken from an aquifer in the San Antonio-El Triunfo mining area in Baja California peninsula,
Mexico. The aquifer material consisted of quartz, feldspar, calcite, chlorite, illite, and
magnetite/hematite. The experiments were conducted using background electrolyte solutions of 0.1
and 0.01 M NaNO3 solutions, and pH ranges of 4 to 13 and 6 to 12 as adjusted with HC1 and NaOH
solutions. The observed percent of sorbed total arsenic, As(V), and As (III) versus pH agreed with
that expected for the adsorption of arsenic on iron oxyhydroxides. Sorption of arsenic on the aquifer
material was high at acidic pH values, and low at alkaline pH values. Arsenic sorption decreased
quickly in the pH range 6 to 9.
Sorption studies have also been conducted to determine the effect on arsenic sorption by the
presence of other anions, such as dissolve phosphate, sulfate, nitrate, and chloride (Reynolds et al.,
1999; Livesey and Huang, 1981). The results of these studies indicate that dissolved phosphate
competes with arsenic for adsorption sites, and thus suppressing arsenic adsorption and enhancing
the mobility of arsenic. Livesey and Huang (1981) found that dissolved sulfate, nitrate, and chloride
present at concentrations present in saline soils had little effect on arsenic adsorption.
As noted previously, As (III) is more mobile and toxic than As(V), and thus redox conditions affect
the relative sorption behavior of As(V) and As (III). Biotic processes affecting the reduction of As(V)
to As (III) have recently been the subject of considerable study. Recently published studies confirm
that microbial processes can affect the mobilization of arsenic in environmental systems by a variety
of direct and indirect oxidation/reduction processes (e.g., Ahman et al., 1997; Cummings et al., 1999;
Jones etal., 2000; Langner and Inskeep, 2000; Zobrist et al., 2000). Ahman et al. (1997) studied the
mobilization of solid-phase arsenic in sediment microcosms from the Aberjona Watershed. Sediment
suspensions were incubated with iron arsenate, which was used as an analogue for sedimentary
arsenic, to investigate the potential for arsenic mobilization by microbial reactions in these sediment
samples. The results indicated that arsenic could be mobilized by microbial-mediated reductive
dissolution of the iron arsenate solids. Arsenic(V) released from the dissolution of the iron arsenate
was also reduced to As (III) under these reducing conditions.
Ahman et al. (1997) also determined that the microbial dissolution/reduction processes could be
prevented by sterilization of the sediment suspensions, which helped to confirm that the processes
were microbially mediated.
Cummings et al. (1999) studied arsenic mobilization by dissimilatory iron-reducing bacteria (DIRB).
Using the dissimilatory iron-reducing bacterium Shewanella alga strain BrY, their experiments showed
that arsenic could be mobilized from crystalline ferric arsenate (the mineral scorodite, FeAsO4'2H2O)
and from As(V) sorbed to sediments from Lake Coeur d'Alene, Idaho. Arsenic mobilization was the
result of dissimilatory (i.e., respiratory) reduction of Fe(III), to Fe(II). Analyses of iron and arsenic in
the solid phases and aqueous solutions indicated that arsenic and iron were present as As(V) and
Fe(II) in both the solid and aqueous phases. Cummings et al. (1999) concluded that even in the
absence of reduction to As (III), As(V) could be mobilized from sediments by biotic reactions of
dissimilatory iron-reducing bacteria.
Jones et al. (2000) investigated the mechanisms controlling the rates of microbial-mediated reduction
of dissolved and adsorbed arsenic. The experiments were conducted with microorganisms obtained
from fine loamy agricultural soil that contained naturally elevated concentrations of arsenic. The
results of microbial-mediated mobilization experiments conducted in the presence of the iron oxides
5.20
-------
goethite or ferrihydrite indicated that the extent of arsenic solubilization depended on the arsenic
surface coverage and on the surface area or crystallinity of the iron oxide phase. During microbial-
mediated reduction, the rate of arsenic desorption from ferrihydrite was determined to be two orders
magnitude greater than that from goethite at similar concentrations of dissolved arsenic. Jones et al.
(2000) proposed that this difference in solubilization rates was due to differences in the rates of
reductive dissolution of the iron oxide phases.
Langner and Inskeep (2000) studied the effect of microbial reduction of dissolved As(V) on the
desorption of As(V) sorbed to ferrihydrite in the absence of reductive dissolution of the Fe(III)-oxide
solid phase. The reduction of As(V) was investigated using a glucose-fermenting Clostridium sp. strain
CN8 enriched from an arsenic contaminated soil. Despite rapid reduction of dissolved As(V) to
As (III), the sorbed arsenic remained primarily as As(V), and desorption of As(V) was too slow to
cause a significant increase in dissolved concentrations of arsenic over the 24-day experiment. Based
on their results, Langner and Inskeep (2000) proposed that the reduction of dissolved As(V) plays a
minor role in the remobilization of As(V) sorbed to iron oxyhydroxide phases, and arsenic release
from contaminated soils may proceed faster via microbial-mediated reductive dissolution of the iron
oxyhydroxides adsorbents.
Zobrist et al. (2000) studied the ability of the baterium Sulfurospirillum barnesii strain SES-3 to reduce
As(V) adsorbed on ferrihydrite and aluminum hydroxide. Their results indicated that cell suspension
of S. barnesii were able to reduce As(V) to As (III) when As(V) is present in solution or when adsorbed
on ferrihydrite or aluminum hydroxide. The results from the experiments conducted with aluminum
hydroxide show that reduction of adsorbed As(V) does not require microbial-remediated reductive
dissolution of ferrihydrite because the aluminum hydroxide does not undergo reductive dissolution.
5.3.5.3 Kj Studies for Arsenic on Soil Materials
Only three published studies containing Kd values for arsenic sorption on soil were identified during
the course of our review. Only one of the identified published studies listing Kd values for arsenic
sorption on soil is reviewed below. In agreement with the other studies of arsenic on single mineral
phases and soils (see above), the results of Mariner et al. (1996) indicate a limited extent of arsenic
adsorption at alkaline pH values. The other two Kd studies (Cornett et al., 1992; Mok and Wai, 1990)
were for environments affected by ore milling operations, and list extremely high Kd values
(>1000's ml/g) for arsenic sorption on sediment. These high Kd values are surprising given the
mobility expected for As(V) and As (III) based on the results of other published sorption studies at all
but the very acidic pH values. Review of the studies by Cornett et al. (1992) and Mok and Wai (1990)
indicate that the reported sorption values were desorption Kd studies, which are typically greater than
adsorption Kd values, and were possibly affected by the presence of arsenic in particulate form (i.e.,
arsenic minerals) in the sediment as opposed to being adsorbed to iron oxyhydroxide mineral
coatings or clays.
Mariner et al. (1996) investigated the effects of high pH on arsenic mobility in a shallow sandy aquifer
at the Commencement Bay Superfund Site in Tacoma, Washington. The groundwater plume
contaminated with arsenic derived from a chemical plant is characterized by high pH and high silica
concentrations. Values of Kd (Table 5.8) calculated from arsenic analyses of core sediment and pore
water samples (i.e., in situ Kd values) at this site decrease by at least an order of magnitude as the pH
increased from 8.5 to 11.
5.21
-------
Table 5.8. Measured arsenic Kd values (ml/g) based on
analyses of an arsenic-contaminated aquifer at a Superfund
Site (Mariner et a/., 1996).
Core
S
S
S
S
S
S
S
S
S
S
Depth
(m)
6.2
6.4
6.5
6.7
6.8
7.0
7.1
7.3
7.9
8.0
Pore
Water pH
8.48
8.39
8.42
8.37
8.30
8.41
8.39
9.73
10.4
10.5
Kd
(ml/g)
0.92
1.44
1.76
2.14
2.63
1.57
1.80
0.31
0.19
0.12
Core
T
T
T
T
T
T
T
T
T
T
Depth
(m)
6.4
6.5
6.7
7.0
7.2
7.3
7.5
7.7
7.8
8.0
Pore
Water pH
9.94
10.1
9.89
10.0
10.4
10.6
10.8
10.8
11.0
10.8
Kd
(ml/g)
6.46
3.91
1.97
0.44
0.42
0.21
0.14
0.69
0.19
0.28
5.3.5.4 Published Compilations Containing Kj Values for Arsenic
Because the references in this section are often cited or used for comparison in other publications,
the following summaries are provided for completeness. It is recommended that the reader review
the original reference and the references cited therein to understand the procedures and sources of
the Kd values used for each compilation. The compilations do not distinguish between oxidation
states for those contaminants that are redox sensitive or consider other important factors that
contribute to variability in sorption, such as pH. Moreover, in cases where very large Kd values are
listed, there is a risk that the original Kd measurement may have included precipitated components.
Baes and Sharp (1983) present a simple model developed for order-of-magnitude estimates for
leaching constants for solutes in agricultural soils. As part of this model development, they reviewed
and determined generic default values for input parameters, such as Kd. A literature review was
completed to evaluate appropriate distributions for Kd values for various solutes, including arsenic.
Because Baes and Sharp (1983) are cited frequently as a source of Kd values in other published Kd
reviews (e.g., Looney et al., 1987), the arsenic Kd values listed by Baes and Sharp are reported here for
completeness. Based on the distribution that Baes and Sharp determined for the Kd values for
cesium and strontium, they assumed a lognormal distribution for the Kd values for all other elements
in their compilation. Baes and Sharp listed an estimated default Kd of 6.7 ml/g for As(V) based on
37 arsenic Kd values from 1.9 to 18 ml/g for agricultural soils and clays in the pH range 4.5 to 9.0.
Their compiled Kd values represent a diversity of soils, pure clays (other Kd values for pure minerals
were excluded), extracting solutions, measurement techniques, and experimental error.
5.22
-------
Looney et al. (1987) tabulated estimates for geochemical parameters needed for environmental
assessments of waste sites at DOE's Savannah River Plant in South Carolina. Looney et al. list Kd
values for several metal and radionuclide contaminants based on values that they found in 1-5
published sources. For arsenic, Looney et al. list a "recommended" Kd of 3.2 ml/g, and a range of 1
to 10 ml/g. Looney et al. note that their recommended values are specific to the Savannah River
Plant site, and must be carefully reviewed and evaluated prior to use in assessments at other sites.
5.3.5.5 Kj Studies of Arsenic on Pure Mineral, Oxide, and Crushed Rock Materials
Numerous adsorption studies have been conducted of arsenic on pure minerals and oxide phases.
The Kd values listed in these studies are not necessarily relevant to the mobility and sorption of
arsenic in soils. However, they are listed in Appendix D for completeness. The references cited in
Appendix D are listed in the main reference list in Chapter 6. The potential value of the references
that they cite and the sorption processes that they discuss is left to the reader.
A significant number of arsenic adsorption studies (Appendix D) have been conducted on iron
oxyhydroxides and clay minerals. The results of these studies demonstrate that
adsorption/desorption and coprecipitation reactions on these minerals can control the concentrations
of dissolved arsenic. The focus of most of these studies, especially those with iron oxyhydroxides
such as ferrihydrite, was primarily on understanding the mechanisms for As(V) and As (III)
adsorption. The adsorption of both As(V) and As (III) on iron oxyhydroxides and clay minerals
varies with pH. As expected for the adsorptive behavior of anions, arsenic is strongly sorbed under
acid conditions where the net surface charge of most minerals is positive. The extent of As(V)
adsorption then decreases with increasing pH in the range of 7 to 9. At alkaline pH conditions,
As(V) is expected to be highly mobile because of the negative net surface charge on most minerals.
The laboratory studies also indicate that As(V) adsorbs more strongly than As (III) to iron and
manganese oxides. Dissolved As (V) and phosphate compete for adsorption sites, and the sorption
of arsenic will be reduced in soils with high phosphate concentrations. Concentrations of other
anions, such as dissolved chloride, nitrate, and sulfate, appear to have little effect on arsenic
adsorption.
No sorption studies of arsenic on crushed rock materials were identified during our review. This was
expected, because studies of contaminant (radionuclide) sorption on crushed rock materials were/are
motivated by performance and risk assessment studies conducted in support of national programs for
the geologic disposal of nuclear waste and long-lived arsenic radioisotopes are not found in nuclear
waste.
5.23
-------
5.4 Curium Geochemistry and Kd Values
5.4.1 Overview: Important Aqueous- and Solid-Phase Parameters Controlling
Retardation
Curium is a transuranic (actinide) element, and can exist in the +3 oxidation state in natural waters.
Curium(IV) is not stable in solutions because of self-radiation reactions (Onishi eta!., 1981). A very
limited number of laboratory and field studies of the aqueous speciation, solubility, and sorption
behavior of Cm(III) exist. However, Cm(III) geochemistry is expected and widely accepted to be
very similar to that of Am(III) and trivalent lanthanide elements, such as europium (III), Eu(III).
Compared to other actinides, Cm(III) and Am(III) are considered to be immobile in soil
environments, and both exhibit high Kd values. However, the tendency of curium, like americium, to
strongly adsorb to soil particles indicates that there is potential for colloid-facilitated transport of
curium. The concentrations of dissolved curium may be controlled by hydroxide or carbonate solids
in some systems. Therefore, some sorption measurements resulting in very high Kd values may
reflect curium precipitation reactions.
5.4.2 General Geochemistry
Curium [Cm, atomic number (Z) = 96] has 21 isotopes (Tuli, 2000). The atomic masses of these
isotopes range from 232 to 252. Most curium isotopes have relatively short half lives of minutes or
less. Although both 242Cm (t,/2 = 162.8 days) and CM-244 (t,/2 = 18.1 years) are released to the
environment from nuclear facilities, CM-244 is probably more significant due to its longer half life
(Coughtrey eta!., 1984).
Balsley et a!. (1997) Curium can exist in the +3 and +4 valence states, but the +3 state is the dominant
valence state in natural waters. Curium(IV) is not stable in solutions because of self-radiation
reactions (Onishi eta!., 1981). Because the electronic structures of the Cm(III), Am(III), and trivalent
lanthanides [rare earth elements (REE)] are similar, the environmental behavior (i.e., aqueous
speciation, solubility, and sorption) of Cm(III) is expected to be very similar to those of Am(III) (see
Section 5.2) and trivalent lanthanide elements, such as Eu(III). This analogy is well established; for
example, see Choppin (1989). The environmental behavior of curium has been reviewed by Silva and
Nitsche (1995), Coughtrey eta!. (1984), Onishi etaL (1981), and Ames and Rai (1978).
5.4.3 Aqueous Speciation
Curium(III) should form complexes with inorganic ligands present in natural waters. However, the
thermodynamic data for curium aqueous species are limited and not well established. We did not
identify any published reviews of thermodynamic properties of curium aqueous species and solids.
With respect to the hydrolysis of curium, Baes and Mesmer (1976) suggest that a good analogy is the
hydrolysis of the lanthanide elements. Curium aqueous and solid compounds are not listed in the
technically recognized compilation of thermodynamic data by Wagman eta!. (1982) from the National
Bureau of Standards [NBS, currently NIST (National Institute of Standards and Technology].
Wimmer eta!. (1992) determined the equilibrium constants for several Cm(III) hydrolytic and
carbonate aqueous complexes by time-resolved laser-induced fluorescence spectroscopy. Figure 5.5
shows the aqueous speciation for Cm(III) based on the constants listed by Wimmer eta!. (1992).
5.24
-------
By analogy, the aqueous speciation of dissolved curium is likely to be similar to those of Am(III) and
trivalent lanthanide elements, such as Eu(III). As shown in Section 5.2, the aqueous speciation of
Am(III) is dominated by the uncomplexed ion Am3+ at acidic to near neutral pH conditions, Am(III)
carbonate complexes at neutral to high pH conditions, and Am(III) hydroxyl complexes at very high
pH conditions.
Figure 5.5. Calculated aqueous speciation for Cm(lll) as a
function of pH. [Curium(lll) aqueous speciation was calculated
based on a total dissolved concentration of curium of 1x10"12 mol/l
(2.43x10"7 mg/l), and the water com position in Table 5.1]
100
o 80
.a
£ 60
CD
40
S. 20
0
Cm
3+
6 7
PH
8
10
Recent aqueous speciation studies of Cm(III) include, for example, those on aqueous
complexation by dissolved carbonate (Fanghanel eta!., 1998); sulfate (Paviet eta!., 1996);
fluoride (Aas eta!., 1999); chloride (Konnecke eta!., 1997); and humic substances (Hummel et
a!., 1999; Panak eta!., 1996; Shin eta!., 1995; Kim eta!., 1993).
5.4.4 Dissolution/Precipitation/Coprecipitation
No studies pertaining to solubility controls for dissolved curium in natural environments were
identified during the course of this review. Given the widely accepted idea that the environmental
behavior of curium is analogous to Am(III), and trivalent lanthanides, such as Eu(III), solubility
controls for curium may be similar to those for Am(III). As discussed in Section 5.2, concentrations
of dissolved Am(III) in soil environments may be controlled by the precipitation of solids such as
Am(OH)3 and AmOHCO3.
5.25
-------
5.4.5 Adsorption/Desorption
5.4.5.1 Guidance for Screening Calculations ofA-dsorbtion
Relative to the other contaminants reviewed in Volume II (EPA, 1999c) and this volume, there are
very few experimental and field studies of the sorption behavior and mobility of curium in
environmental systems. The results of curium adsorption studies published prior to 1981 have been
reviewed by Ames and Rai (1978) and Onishi et al. (1981). More current studies are reviewed in the
sections below. All of the available sorption studies indicate that Cm(III) readily sorbs to minerals,
crushed rock, and soil materials, and is therefore considered one of the most immobile actinide
elements in the environment.
The limited number of Kd adsorption studies for Cm(III) in soils prevents calculation of Kd look-up
tables. However, the sorption behavior of Cm(III) is very similar to that of Am(III) (see
Section 5.2) and trivalent lanthanide elements, such as Eu(III). Guidance given above for Kd values
for Am(III) in Section 5.2 can be used for screening calculations of Cm(III) migration in soils.
Given the absence of definitive maximum and minimum Kd values for Cm(III) as a function of the
key geochemical parameters, such as pH, EPA suggests that Kd values measured for Cm(III) at site-
specific conditions are thus essential for site-specific contaminant transport calculations and
conceptual models. One of the major recommendations of this report is that for site-specific
calculations, partition coefficient values measured at site-specific conditions are absolutely essential.
The geochemical behavior of Eu(III) has been determined to be similar to the trivalent actinides,
such as Cm(III). Therefore, the following generalities made by Clark et al. (1998) regarding the
adsorption of Eu(III) should also apply to Cm(III) and Am(III):
• European (III) adsorption increases with increasing pH.
• European (III) adsorption is sensitive to the ionic strength at low concentrations of total
Eu(III) and to calcium at higher concentrations of total europium.
• Ion exchange is an important sorption mechanisms for Eu(III), especially at pH values less
than 4.5.
• Mobility of Eu(III) in the environment increases at low to moderate pH values and in high
ionic strength solutions.
Available curium sorption studies indicate that sorption of curium is strongly pH dependent and
increases with increasing pH with peak adsorption occurring between pH values of 5 and 6. The
observed pH dependence is expected, because the dominant aqueous species of curium in the pH
range of natural waters are primarily cations such as Cm3+ and Cm(III) carbonate complexes at
acidic and basic pH values, respectively.
Compared to other actinides, Cm(III) is considered to be immobile in soil environments, and
exhibits high Kd values. Because the concentrations of dissolved curium may be controlled by
hydroxide or carbonate solids in some systems, some sorption measurements resulting in very high
Kd values may have been affected by curium precipitation reactions.
Because Cm(III) strongly adsorbs to soil particles, there is potential for colloid-facilitated transport
of curium. Laboratory measurements by Kaplan et al. (1994) using site-specific soils and a 2-phase
5.26
-------
solute transport code indicated that americium, curium, plutonium, and uranium would migrate less
than 10 meters (m) in the F-Area of the Savannah River Site. The contaminants however were
found associated with colloidal-size particles in groundwater 1,200 m away from the point source.
Colloid-facilitated migration of contaminants is reviewed in Section 2.7 in Volume I (EPA, 1999b).
The reader is cautioned that the importance of colloid-facilitated migration, especially in aquifer
systems that do not involve fracture flow of groundwater, is still the subject of debate.
5.4.5.2 General .Adsorption Studies
Sheppard et al. (1979) used batch equilibration experiments to study the sorption of curium to
colloidal-size soil particles that are potentially diffusible in soil/water systems. The experiments
were conducted with distilled water and 14 soils from Muscatine, Illinois; Hanford, Washington;
Barnwell, South Carolina; Idaho Falls, Idaho; and Paradise and Placerville, California.
Centrifugation measurements indicated that much of the CM-244 was retained by the colloidal-size
soil particles. Sheppard et al. (1979) found it difficult to correlate the sorption results with chemical
and physical properties of the soils. They attributed this to the lack of precise distribution ratios,
competition with cation exchange reactions, and complexation with humic and fulvic acid materials.
The results also indicated that the sorption behavior of Am-241 and CM-244 under these
experimental conditions was identical. Sheppard et al. (1979 ) suggest that colloids of clay and humic
acids are potentially important processes for the transport of actinides in soil/water systems.
Means et al. (1978) studied the mineralogy of the adsorbents for Cm-244 in soil near a disposal
trench at nuclear waste burial grounds at the Oak Ridge National Laboratory in Oak Ridge,
Tennessee. Their analyses indicated that curium concentrations were highly correlated to the
concentrations of manganese oxides, even though the concentrations of iron oxyhydroxides and
organic carbon were significantly greater than those of manganese oxides.
5.4.5.3 Kj Studies for Curium on Soil Materials
Nakayama and Nelson (1988) measured the pH dependency of Kd values for Am-243 and Cm-244
sorption onto sediment. Changes in pH were found to have a significant effect on the adsorption of
these two radionuclides. However, variation of pH did not cause a significant difference in the
adsorption behaviors of these two radionuclides relative to each other. Sibly and Wurtz (1986)
measured the Kd values for Cm-244 on freshwater and estuary sediments. They found that Kd
values previously reported for americium were equivalent to their measured values for curium.
Adsorption of curium was pH dependent and increased from pH 4 to pH 7. The measured
adsorption remained high in low organic, carbonate-buffered systems. It should be noted that the
measured adsorption of several lanthanide trivalent cations on amorphous Fe(OH)3, hematite, and
magnetite were also determined to be strongly dependent on pH (Music and Ristic, 1988).
Nishita et al. (1981) studied the extractability of Cm-244 from several types of soils as a function of
pH. The extractability of curium was considered to parallel the solubilization of aluminum, iron,
and/or manganese hydrous oxides, which are important adsorbents for dissolved contaminants in
soil systems (Nishita et al., 1981). The curium Kd values and soil characteristics determined for these
soils by Nishita et al. (1981) are listed in Table 5.9. Under highly acidic conditions, Nishita et al.
(1981) concluded that curium was present in ionic form and was mobile. At higher pH values,
curium formed hydroxyl complexes that were readily sorbed by the soils.
5.27
-------
Table 5.9. Measured curium Kd values (ml/1) and soil
properties for soils studied by Nishita et al. (1981).
Soil
Type
Silt Clay Loam
Sharpsburg
Sandy Loam
Malbis
Sandy Loam
Lyman
Silty Clay
Holts ville
Loam
Aiken
Silt Loam
Yolo
Muck
Egbert
pH1
5.9
5.3
5.0
7.8
6.0
6.7
7.2
OM2
Fraction
(%)
2.8
2.4
5.7
0.6
8.4
2.5
40.8
CEC3
(meq/100 g)
20
15
15
30
15
25
60
Free Fe
Oxides
(%)
1.29
1.65
1.52
1.20
5.29
2.41
1.57
Mn4
Fraction
(%)
0.06
0.05
0.04
0.04
0.10
0.08
0.10
Extract
pH
5.41
6.56
4.39
5.71
4.58
6.17
7.12
8.04
5.71
6.72
6.12
6.98
7.14
7.54
Kd
(ml/g)
23,350
13,860
8,523
6,809
1,374
186
36,260
51,900
30,920
15,020
21,780
17,090
6,772
5,056
1 Saturated paste. 4 In 4 M HNO3 extract.
2 Organic matter.
3 Cation exchange capacity.
5.4.5.4 Published Compilations Containing Kj Values for Curium
Because the references in this section are often cited or used for comparison in other publications,
the following summaries are provided for completeness. It is recommended that the reader review
the original reference and the references cited therein to understand the procedures and sources of
the Kd values used for each compilation. The compilations do not typically consider important
factors that contribute to variability in sorption, such as pH. Moreover, in cases where very large Kd
values are listed, there is a risk that the original Kd measurement may have included precipitated
components.
Baes and Sharp (1983) present a simple model developed for order-of-magnitude estimates for
leaching constants for solutes in agricultural soils. As part of this model development, they reviewed
and determined generic default values for input parameters, such as Kd. A literature review was
completed to evaluate appropriate distributions for Kd values for various solutes, including curium.
Because Baes and Sharp (1983) are cited frequently as a source of Kd values in other published Kd
reviews (e.g., Looney et al., 1987, Sheppard and Thibault, 1990), the curium Kd values listed by Baes
and Sharp are reported here for completeness. Based on the distribution that Baes and Sharp
5.28
-------
determined for the Kd values for cesium and strontium, they assumed a lognormal distribution for
the Kd values for all other elements in their compilation. Baes and Sharp listed an estimated default
Kd of 3,300 ml/g for curium based on 31 Kd values that ranged from 93.3 to 51,900 ml/g for
agricultural soils and clays over the pH range 4.5 to 9.0. Their compiled Kd values represent a
diversity of soils, pure clays (other Kd values for pure minerals were excluded), extracting solutions,
measurement techniques, and experimental error.
Looney et al. (1987) tabulated the estimates for geochemical parameters needed for environmental
assessments of waste sites at DOE's Savannah River Plant in South Carolina. Looney et al. list Kd
values for several metal and radionuclide contaminants based on values that they found in 1-5
published sources. For curium, Looney et al. list a "recommended" Kd of 3,160 (103'5) ml/g, and a
range of 100 to 1,000,000 ml/g. Looney et al. note that their recommended values are specific to the
Savannah River Plant site, and they must be carefully reviewed and evaluated prior to use in
assessments at other sites.
Thibault et al. (1990) (also see Sheppard and Thibault, 1990) present a compilation of soil Kd values
prepared to support radionuclide migration assessments for a Canadian geologic repository for spent
nuclear fuel in Precambrian Shield plutonic rock. Thibault et al. collected Kd values from other
compilations, journal articles, and government laboratory reports for important elements, such as
curium, that would be present in the nuclear fuel waste inventory. The curium Kd values listed in
Thibault etal (1990) are included in Table 5.10. Thibault et al. (1990) describe the statistical methods
used for analysis of the compiled Kd values. The range for the Kd values used to calculate the
"geometric mean" cover several orders of magnitude. Readers are cautioned against using
"geometric mean" values or any other form of averaged Kd values as "default" Kd values, as such
values are usually calculated from data compiled from different investigators for different soil
systems. These mean or average values do not represent any particular environmental system and
geochemical conditions. As discussed in Volume I (EPA, 1999b), the variation observed in the
literature for Kd values for a contaminant is due to differences in sorption mechanisms, geochemical
conditions, soil materials, and methods used for the measurements.
Table 5.10. Curium Kd values (ml/g) listed in
Thibault etal. (1990, Tables 4 to 8).
Soil Type
Sand
Silt
Clay
Organic
Kd Values (ml/g)
Geometric
Mean
4,000
18,000
6,000
6,000
Number of
Observations
2
4
1
1
Range
780 - 22,970
7,666 - 44,260
5.29
-------
5.4.5.5 Kj Studies of Curium on Pure Mineral, Oxide, and Crushed Rock Materials
Few adsorption studies have been conducted of curium on pure minerals, oxide phases, and other
geologic-related materials. The Kd values listed in these studies are not necessarily relevant to the
mobility and sorption of curium in soils. However, they are listed in Appendix E for completeness.
The references cited in Appendix E are listed in the main reference list in Chapter 6. The potential
value of the references that they cite and the sorption processes that they discuss is left to the reader.
The studies of curium sorption on crushed rock were conducted typically as part of national research
programs to investigate the feasibility of geological disposal of high-level radioactive waste (HLW).
5.5 Iodine Geochemistry and Kd Values
5.5.1 Overview: Important Aqueous- and Solid-Phase Parameters Controlling
Retardation
Although the environmental chemistry of iodine is normally assumed to be simple and well known,
recent studies suggest that the fate and mobility of iodine in environmental systems may be more
complex than expected. This complexity is caused by the multiple redox states of iodine that may
exist under oxidizing conditions. The -1, +5, and molecular I2 oxidation states are those most
relevant for iodine in environmental systems. In most aqueous environments, iodine is present as
the iodide ion, I". In marine and highly oxidizing environments such as surface waters and some
oxygenated shallow ground waters, iodine may be present in the +5 oxidation state as the iodate ion,
IO3. Under oxidizing, acidic conditions, molecular 1° (aq) may form from the reduction of IO3 or
oxidation of T. Some sorption studies suggest (Section 5.5.5) that these iodine oxidation/reduction
reactions may affect the observed sorption behavior of iodine in soils and the reactions may be
affected by the organic content and/or microbial processes in soils and sediments. Although iodine
is a primary component of several naturally occurring minerals, the formation of such minerals
represent unique geological conditions and is not likely to occur in soils due to the low
concentrations of iodine typically present in the environment.
Sorption of iodine species appears to be controlled in part by soil organic matter and in part by iron
and aluminum oxides, with adsorption of iodine becoming increasingly important under more acid
conditions. Although the extent of sorption is typically low, especially in systems containing little or
no organic matter, I" and IO3 are sorbed to a measurable extent by soils and some oxide and sulfide
minerals at near neutral and alkaline pH conditions. The adsorption behavior of IO3 also appears to
be appreciably different from that of I", in that IO3 sorbs much more strongly than I" to soil and
mineral surfaces. Mechanisms causing this sorption behavior of iodine at near neutral and alkaline
pH conditions are not completely understood. Some have proposed that this observed adsorption
behavior in soils may be a result of the oxidation of Tand/or reduction of IO3 to the more reactive
molecular 1° (aq) and/or its hydrolysis products.
5.5.2 General Geochemistry
Iodine [I, atomic number (Z) = 53] is a halide element, like fluorine, chlorine, and bromine. Iodine
is known to occur in the -1, +1, +3, +5, and +7 oxidation states. There are 37 reported isotopes of
iodine (Tuli, 2000). Only 1 isotope, 127I is stable. Isotopes 131I and 129I are generated during the
operation of nuclear power plants, reprocessing of nuclear fuels, and testing of nuclear weapons.
5.30
-------
Due to its long half life (ti/2 — 1.57xl07y), 129I is the only iodine isotope of major environmental
concern. Fission products 131I ((t,/2 = 8.02 days) and 125I (t,/2 = 59.4 days) are often short term
disposal hazards. Most of the remaining iodine isotopes have half lives of a few hours to fractions
of a second. The environmental behavior of iodine has been reviewed by others, such as Lieser and
Stemkopff (1989), Whitehead (1984), Coughtrey et al. (1983), and Ames and Rai (1978).
5.5.3 Aqueous Speciation
Iodine is known to occur in several oxidations. Of these, the -1, +5, and molecular I2 oxidation
states are those most relevant to environmental systems. Figure 5.6 is an Eh-pH diagram that shows
the dominant aqueous species of iodine as a function of pH and Eh (redox potential) at 25°C with
respect to the thermodynamic stability of water.
In most aqueous environments, iodine is present in the -1 valence state as the iodide ion, I". The
stability range of I" extends almost over the entire pH and Eh range for the thermodynamic stability
of water. In marine and highly oxidizing environments such as surface waters and some oxygenated
shallow ground waters, iodine may be present in the +5 oxidation state as the iodate ion, IO3
(Figure 5.6). Under oxidizing, acidic conditions, the Eh-pH diagram indicates that molecular 1° (aq)
may form from the reduction of IO3 or oxidation of I". Some iodine sorption studies suggest
(Section 5.5.5) that these iodine oxidation/reduction reactions may affect the observed sorption
behavior of iodine in soils and the reactions may be affected by the organic contents and/or
microbial processes in soils and sediments. For example, studies, such as Skogerboe and Wilson
(1981), indicate that fulvic acid derived from soil is cable of reducing molecular 1° (aq) and IO3 to I"
under conditions generally characteristic of natural waters.
The volatilization of iodine from soil to the atmosphere may occur as a result of both chemical and
microbiological processes (Whitehead, 1984). The chemical processes generally result in molecular
iodine or hydrogen iodide, and the microbiological processes yield organic compounds, such as
methyl iodide. Methyl iodide is not strongly retained by soil components and is only slightly soluble
in water (Whitehead, 1984).
Vovk (1988) completed a review of the literature regarding the influence of radiolysis on the
migration of iodine in the geosphere, including published studies of iodine liberation from uranium
ore into groundwater and iodine in formation waters (i.e., brines in oilfields and gasfields) from deep
strata. A variety of other iodine species are known to form radiolytically, such as HOI° (aq) and OI".
Although these species are unstable, they may be important in some chemical reactions that affect
iodine in the environment. Radiolytic reactions lead to conversion of non-volatile and relatively
unreactive species, such as I" and IO3, into volatile forms, such as molecular I2, which are fairly
reactive and react rapidly with organic compounds to form organic iodide species (Vovk, 1988).
5.31
-------
Figure 5.6. Eh-pH stability diagram for dominant iodine
aqueous species at 25°C. [Diagram based on a total
concentration of 10"6 mol/l dissolved iodine.]
.5
o
LU
0
-.5
HI03° (aq)
0
4 6 8 10
PH
12
14
5.5.4 Dissolution/Precipitation/Coprecipitation
Iodine can be found as a primary component in some rare, naturally occurring minerals that are
associated with evaporite1 and brine deposits (Johnson, 1994; Doner and Lynn, 1977). The iodine is
commonly present in substitution for other halogen elements, such as chloride and bromide. lodate
is typically associated with sulfate- or nitrate-type minerals. For example, the northern Chilean
nitrate deposits contain iodine minerals, such as bruggenite [Ca(IO3)2'H2O], lautarite [Ca(IO3)2], and
dietzeite [Ca2H2O(IO3)2(CrO4)] (Johnson, 1994). Such minerals are expected to be highly soluble in
geochemical systems. Therefore, the formation of iodine-containing minerals is not likely to be
important in most soil systems due to the low concentrations of iodine typically present in the
environment.
An evaporite is a sedimentary rock composed principally of minerals precipitated from a saline solution as a result of
extensive or total evaporation of the solution.
5.32
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5.5.5 Adsorption/Desorption
5.5.5.1 Guidance for Screening Calculations ofA-dsorbtion
Sorption of iodine species appears to be due in part to soil organic matter and in part to iron and
aluminum oxides, with adsorption of iodine becoming increasingly important under acidic
conditions. Because iodine is present as either the anions I" or IO3 in most soils, conventional
wisdom suggests that their adsorption on soils and most individual mineral phases should be
negligible at near neutral and alkaline pH conditions, and increase as pH values become more acidic.
Sorption studies of iodine published prior to 1976 are reviewed by Onishi et al. (1981) and Ames and
Rai (1978). More recent studies are reviewed in the sections below. The majority of these studies
pertain to the adsorption of iodide.
Numerous studies have been conducted in which Kd values for iodide adsorption on soil have been
reported along with associated pH values and/or soil organic carbon contents. These include the Kd
measurements by Kaplan et al. (1996, 1998a, 1998b, 2000a, 2000b), Fukui et al. (1996), Bird and
Schwartz (1996), Sheppard et al. (1995), Serne et al. (1993), Muramatsu et al. (1990), Sheppard and
Thibault (1988), and Gee and Campbell (1980) (see reviews below). However, these data are too
scattered or limited in scope to determine Kj look-up tables of conservative minimum and
maximum Kd values for iodide as a function of key geochemical parameters such as pH and organic
carbon content of soil. One of the key recommendations of this report is that for site-specific
calculations, partition coefficient values measured at site-specific conditions are absolutely essential.
Based on the studies by Kaplan and coworkers, it is suggested that Kj values of 0.6 and 0 ml/g be
considered, respectively, as conservative minimum Kd values for the pH ranges from 4 to <6 and
>6 to 10 for screening calculations of iodide migration in soils. These values represent primarily low
adsorptive, quartz-feldspar-rich soils with low contents of organic carbon matter. The available Kd
data are too unevenly distributed and limited as a function of pH to estimate a conservative
maximum Kd values for iodide. Values of Kd for iodide have been reported in the range from 1 to
10 ml/g for the pH range from 4 to 10. Most of the reported Kd values are however typically less
than 3 ml/g.
Studies suggest that the adsorption of iodide increases with increasing soil organic matter.
Unfortunately, the majority (>90%) of the reported Kd values for iodide adsorption on soil are
limited to mineral soils with organic matter contents of less than 0.2 wt%. Development of a look-
up table for iodide Kd values as a function of organic matter contents is precluded by this limited
range of data.
Sorption studies (see studies reviewed in sections below and references cited therein) indicate that
the adsorption of iodine anions does increase with decreasing pH. However, although the extent of
sorption is typically low, especially in systems containing little or no organic matter, I" and IO3 are
sorbed to a measurable extent by soils and some oxide and sulfide minerals at near neutral and
alkaline pH conditions. The adsorption behavior of IO3 also appears to be appreciably different
than that of I", in that IO3 sorbs much more strongly than I" to soil and mineral surfaces.
Mechanisms causing this sorption behavior of iodine at near neutral and alkaline pH conditions are
not completely understood. Some have proposed that this observed adsorption behavior in soils
5.33
-------
may be a result of the oxidation of I" and/or reduction of IO3 to the more reactive molecular 1° (aq)
and/or its hydrolysis products (Yu et al, 1996; Behrens, 1982; Whitehead, 1974). Such a reaction
process would help explain why some adsorption measurements do not reach steady-state
conditions even after many weeks or months of contact time. Some investigators believe that this
iodine oxidation/reduction process may be affected by the organic matter and/or microbial
processes in the soils. Some researchers have found positive correlations between the sorption of
iodine and organic matter in soils, and some studies indicate that the iodine sorption is decreased
when soils are treated with fungicide, bactericide, irradiation or heat. Others have proposed that
iodine sorption was primarily caused by physical processes associated with the surfaces and
entrapment in the micropores and structural cavities in the organic matter.
5.5.5.2 General Adsorption Studies
Kaplan et al. (2000b) investigated the sorption of 125r on calcite, chlorite, goethite, montmorillonite,
quartz, vermiculite, and illite. These minerals were components of the sediments used for iodine
sorption studies by Kaplan et al. (1996). Kaplan et al. (2000b) measured little or no sorption of I" on
calcite, chlorite, goethite, montmorillonite, quartz, and vermiculite. However, a significant amount
of T sorption (Kd = 15.14+2.84 ml/g) was measured for the illite sample, and was also determined
to be strongly dependent on pH. The Kd values for I" on illite decreased from 46 to 22 ml/g over
the pH range 3.6 to 9.4, respectively. Kaplan et al. (2000a) were able to desorb a significant fraction
of the T by treating the 125I~sorbed illite samples with dissolved fluoride (43 percent desorbed),
dissolved chloride (45 percent desorbed), dissolved bromide (52 percent desorbed), and dissolved
stable T (83 percent desorbed). Based on these results, Kaplan et al. (2000a) proposed that I" was
sorbed to the illite by reversible physical adsorption to pH-dependent edge sites.
Radlinger and Heumann (2000) studied the conversion of inorganic I" into humic substance/iodine
species in aquatic systems, and whether microorganisms affected such reactions. For natural
samples, I" was exclusively fixed by humic substance fractions where natural humic substance/iodine
species have also been observed in the original samples. Their results indicated that the
transformation was strongly affected by microbiological activity of the sample and occurred by a
complicated transfer mechanism by the formation of different intermediate humic substance/iodine
species.
Yu et al. (1996) completed a series of experiments to compare the reaction of I" with a series of
volcanic-ash soils to those for noncrystalline materials that constitute much of the inorganic fraction
of these soils. The measured uptake of I" by volcanic soils was different than the I" on the
noncrystalline phases that make up much of the clay-size fraction of some of these soils. Sorption
of T onto imogolite and ferrihydrite was rapid (<30 min), but was not particularly great. Iodide
sorption on imogolite was 3-to-4 times greater than that on ferrihydrite on a mass basis. In contrast,
rates of I" retention by volcanic-ash soils were slow and did not attain a steady-state after 300 h. Yu
et al. (1996) determined that this reaction was and its extent could be attenuated, but not be
suppressed, by sterilization. The I" sorbed by the soils could only be completely desorbed by boiling
the soils in 2 molar (M) sodium hydroxide solution. The results indicated that the amount of I"
sorbed by soils was inversely correlated with pH. However, no relationship with organic matter
concentrations, surface area, or imogolite and ferrihydrite concentrations was found. Yu et al. (1996)
determined that the reaction of I" with volcanic-ash soils was consistent with a rapid initial uptake by
soil mineral surfaces, followed by a slower reaction of soil organic matter with oxidized forms of I".
5.34
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They proposed that the sorption of I" to soils is primarily controlled by the slow oxidation of I" to
1° (aq) and that 1° (aq) or its hydrolysis product HOI° (aq) then reacts with soil organic matter.
Sheppard and Hawkins (1995) studied bog water and peat from an iodine-rich bog to investigate the
role of microorganisms in the retention and accumulation of iodine in a temperate wetland. Their
results indicated that sorption of iodine on fresh peat was slightly slower and more limited under
anoxic conditions. Autoclaving the peat and the addition of sucrose to the peat both decreased
iodine sorption. In addition, the reinoculation of the treated peat with a fresh microbial population
was not effective in increasing iodine sorption. This suggested that microbes may only play a minor
and indirect role in iodine sorption through the decomposition of organic matter.
Sheppard and Thibault (1992) investigated the behavior of iodine in organic and mineral soils. Their
results showed that microbes played an indirect role in sorption of iodine on soils, and their effect
was very short term. Desorption experiments indicate that the released iodine was generally anionic,
with some neutral or positively charge species also present. They found no firm evidence that
retention of iodine in soils was due to any specific chemical reaction. Sheppard and Thibault (1992)
proposed that iodine sorption was primarily through physical processes associated with the surface
and entrapment in the micropores and structural cavities of the organic matter.
Bors et al. (1991) conducted batch adsorption experiments of 125I on two types of soils, chernozem1
and podzol.2 Generally higher Kd values were observed for the chernozem soil, which is
characterized by a higher amount of organic matter and soil biomass. Incubation of soil samples
under varied conditions suggested that soil microflora participated in iodine immobilization. Bors et
al. (1988) used batch experiments to study the relationship of iodine sorption to soil organic C-
contents of different horizons of a podzolic soil profile. The measurements indicated that the
sorption of iodine tended to increase with increasing organic C-content of different soil horizons.
Analyses of the organic substances suggested that a considerable part of the radioiodine was fixed in
humic and fulvic acids.
Koch et al. (1989) investigated the role of microbial processes in the mobility of 129I for 12 samples
of nine organic soils, which varied widely in degree of humification and parent vegetation. The soils
samples were collected mainly on the Precambrian Shield of Ontario, Canada. Their experiments
were conducted using glucose, thymol, gamma irradiation (60Co) to stimulate or suppress microbial
activity in the soils. The presence of glucose generally increased I" removal from solution, whereas
thymol depressed removal. Gamma irradiation of the soils decreased I" removal from solution in all
samples. The iodine content of soils was directly related to both biologically- and nonbiologically-
mediated processes of I" removal from solution. Koch et al. (1989) concluded that microorganisms
play an important role in the processes of I" removal from solution in organic soils.
Chernozem and podzol are great soil groups of the 1938 USDA classification system. It is a group of zonal soils
whose surface horizon is dark and highly organic, below which is a lighter-colored horizon and accumulation of lime.
The soil group is developed under conditions of temperate to cool subhumid climate (USDA, 1938).
Podzol is a group of zonal soils having an organic mat and a very thin organic-mineral layer overlying gray, leached
A2 horizon and a dark brown, illuvial B horizon enriched in iron oxide, alumina, and organic matter. It develops under
coniferous or mixed forests or under heath, in a cool to temperate moist climate (USDA, 1938).
5.35
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Behrens (1982) found that iodine in both surface and soil water was to a large extent chemically
converted to an inert form by microbial processes. He presented evidence of involvement of
microorganisms in that autoclaving of the soils greatly reduced iodine sorption by the soil. Behrens
(1982) suggested that the reactions involve extracellular, enzymatic oxidation of I" to molecular
1° (aq) which then reacts with organic matter.
Benes et al. (1976) studied the interaction between humus and iodine in lake water using
centrifugation, ultrafiltration, and filtration through ion exchange membranes. Iodine was retained
to a considerable degree by the ultrafilter. This observation suggested that iodine was present not
only as simple halide ion, which should not be retained by the filter or simple adsorption reactions.
Benes et al. (1976) suggested that some kind of reaction between the halides and humus might be
responsible for the observed uptake of iodine.
Whitehead (1973,1974,1978) conducted several studies on iodine adsorption on geologic materials
and organic matter. Whitehead (1973) determined the iodine contents of soils from sites in the
United Kingdom. His results showed that soil organic matter and iron and aluminum oxides were
largely responsible for the retention of iodine in soils. The total concentrations of iodine in these
soils were positively correlated with aluminum oxide extracted by Tamm's reagent, with ferric oxide
extracted by citrate-dithionite, and with soil organic matter. Whitehead (1974) studied the uptake of
T, element iodine, and IO3, when added to sandy loam and to mixtures of the soil with organic
matter (composited grass roots), chalk, ferric oxide, and aluminum oxide (sesquioxides). The
behavior of elemental iodine was similar to I" in response to the various soil treatments. lodate,
however, differed considerably from the other two forms of iodine. With soil alone and with the
soil/chalk mixture, the decrease in the concentrations of dissolved IO3 with time was relatively slow,
although after 160 days its solubility was similar to that of I" and elemental iodine. The
incorporation of composted grass roots cause a rapid reduction in IO3 solubility, suggesting that the
organic matter accelerated the reduction of IO3 to elemental iodine or I". The results obtained when
the soil was treated with composted grass roots suggest that, in the presence of readily
decomposable organic matter, IO3 may be quickly reduced to elemental iodine or I". Whitehead
(1978) measured the iodine content in successive 10-cm horizons of 18 soil profiles from England
and Wales. The iodine content was correlated with the contents of "free" aluminum and iron oxides
(oxalate-soluble) and organic matter. In all 154 samples, the iodine content was closely correlated
with oxalate-soluble aluminum, but not with oxalate-soluble iron or organic matter. In the 5 most
acidic soils with pH values less than 4.8, however, the iodine content was more closely correlated
with iron than with aluminum.
5.5.5.3 Kj Studies of Iodine on Soil Materials
Kaplan and coworkers have conducted a series of laboratory studies on the adsorption of I" on soils
and mineral phases. Kaplan et al. (2000a) measured Kd values for I" as a function of pH on two
sediments from the Savannah River Site at Aiken, South Carolina. The Kd values ranged from 3.96
to 0.05 ml/g depending on sediment type and especially pH. The Kd values were consistently
greater for a wetland versus an upland sediment. Kaplan et al. (2000a) attributed this result to the
greater organic matter content of the wetland sediment. Their measurements also indicated that the
Kd values increased with decreasing pH as would be expected for dissolved anionic contaminants.
For the wetland sediment, the Kd values ranged from 3.96 to 1.17 ml/g. For the upland sediment,
the Kd values ranged from 0.90 to 0.05 ml/g.
5.36
-------
Kaplan et al. [2000b; also see Kaplan et al. (1996)] used the batch technique to measure the Kj values
for iodine, as I", on three sediment samples from the Hanford Site in southeastern Washington. The
measurements were conducted with a groundwater sample (pH 8.46) taken from a well located in an
uncontaminated area of the Hanford Site. The sediment samples included a Touchet bed sand
(sample TBS-1), a silty loam (sample Trench AE-3), and a very coarse sand (sample Trench-94).
The average measured Kd values at 330 days were 9.83, 6.83, and 4.72 ml/g for the very coarse sand,
Touchet bed sand, and the silty loam sediment samples, respectively. Kaplan et al. (2000b) also
noted that the Kd values increased as a function of time. For example, the Kd values measured for
the very coarse sand increased from 0.19 to 9.83 over 330 days of contact time.
Kaplan et al. (1998a) used the batch equilibration technique to measure the Kd values for I" under
oxic conditions on 20 sediment samples taken from a borehole in the Hanford formation at the
Hanford Site in southeastern Washington. Each sediment sample was equilibrated with
uncontaminated groundwater from the Hanford Site, which was subsequently spiked with 125I prior
to the Kd measurements. The groundwater solution has a low ionic strength and a pH of 8.4 [see
Table 1 in Kaplan et al. (1998a)]. The mean Kd values of three replicates measured for I" on each
Hanford sediment sample are listed in Table 5.11, and range from -0.03 to 0.23 ml/g. The
measurements indicated essentially no sorption of I" on Hanford sediments under these conditions.
Kaplan et al. (1998b) studied the effects of high pH on the sorption of iodine under oxic conditions
on sediments from the Hanford Site in Washington. Batch sorption experiments were completed
using the <2-mm size fraction of sediment collected from Trench AE-3 in the 200 Area of the
Hanford Site. The sediment was characterized as a silty loam with a cation exchange capacity (CEC)
of 6.4 meq/100 g. The carbonate content of the sediment consisted of primarily of calcite. Based
on analyses reported by Kaplan et al. (1996) for a sediment from a nearby location, the sediment was
expected to contain approximately 0.2-0.5 wt. percent amorphous Fe2O3 (Kaplan et al., 1998b). The
<2-mm size fraction contained 41 percent sand, 50 percent silt, and 9 percent clay. The clay fraction
of sediment contained primarily smectite (57 percent), illite (19 percent), and vermiculite
(14 percent). The groundwater selected for the sorption studies was an uncontaminated
groundwater from the Hanford Site with a low ionic strength and pH of 8.4 [see Table 2 in Kaplan
et al. (1998b)]. Based on the experimental conditions and geochemical modeling calculations, the
dissolved iodine was assumed to present as dissolved I". The mean Kd values of three replicates
measured for iodine in the high pH experiments are listed in Table 5.11.
Yoshida et al. (1998) used the batch equilibration method and 125I to measure the Kd values for I" and
IO3 on 68 soils. The soils were collected from upland fields, paddies, forests and open areas
throughout Japan. The soils were characterized with respect to soil pH, cation exchange capacity
(CEC), anion exchange capacity (AEC), active aluminum (extractable by oxalic acid and ammonium
oxalate), and total organic carbon (TOC). Two sets of experiments were conducted. One set of
experiments consisted of equilibrating wet (fresh) samples of each soil with deionized water. In the
second set of experiments, samples of each soil were autoclaved prior to equilibration with
deionized water. Table 5.12 lists the median Kd values for the 9 soil types. The measured Kd values
range widely for the wet soils from 1.1 to 10,200 ml/g for I", and from 2.1 to 8,210 ml/g for IO3.
These values were almost 2 orders of magnitude higher than Kd values previously reported for air-
dried Japanese soils. The Kd values for the sandy soil samples were low. Yoshida et al. (1998) found
relatively good correlations between Kd for wet soils and TOC and total nitrogen, indicating the
contribution of organic materials on the sorption of iodine. The Kd values decreased substantially
after autoclaving (Table 5.12). The effects of autoclaving on the sorption of I" were greater than
5.37
-------
those measured for IO3 sorption. Yoshida et al. (1998) estimated that 86 percent of I" sorption and
50 percent of IO3 sorption were attributable to microbial activities and/or soil fractions that were
affected by autoclaving.
Table 5.11. Kd values (ml/g) for the adsorption of iodine to
sediments from the Hanford Site in southeastern Washington
state (Kaplan ef a/., 1998a,1998b).(milliequivalent-meq)
pH
8.54
8.80
8.77
8.73
8.75
8.77
8.52
8.50
8.52
8.56
8.94
8.82
8.81
8.89
Kd
(ml/g)
0.00
±0.01
-0.01
±0.02
0.00
±0.03
-0.03
±0.02
-0.03
±0.03
0.06
±0.17
0.00
±0.02
-0.01
±0.01
-0.01
±0.02
-0.01
±0.01
0.12
±0.04
0.13
±0.06
0.02
±0.03
0.09
±0.05
CEC
(meq/100 g)
5.07
4.73
4.60
4.62
4.11
2.32
4.98
4.72
4.67
4.56
7.33
8.41
9.03
6.63
Solution
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Soil
Identification
Hanford sediment
B8500-07A
Hanford sediment
B8500-10A
Hanford sediment
B8500-12A
Hanford sediment
B8500-14A
Hanford sediment
B8500-15A
Hanford sediment
B8500-16A
Hanford sediment
B8500-17A
Hanford sediment
B8500-19A
Hanford sediment
B8500-20A
Hanford sediment
B8500-21A
Hanford sediment
B8500-22A
Hanford sediment
B8500-23A
Hanford sediment
B8500-24A
Hanford sediment
B8500-25A
Reference
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
5.38
-------
Continuation of Table 5.11
pH
8.88
8.84
8.56
8.93
8.92
8.89
8.1
9.9
10.2
11.0
11.9
Kd
(ml/g)
0.06
±0.05
0.23
±0.06
0.04
±0.07
-0.01
±0.01
0.08
±0.08
0.01
±0.02
0.22
±0.01
0.01
±0.01
-0.02
±0.02
-0.04
±0.02
0.01
±0.01
CEC
(meq/100 g)
8.36
7.77
10.98
8.39
6.21
6.65
6.4
6.4
6.4
6.4
6.4
Solution
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater and NaOH
Groundwater and NaOH
Groundwater and NaOH
Groundwater and NaOH
Groundwater and NaOH
Soil
Identification
Hanford sediment
B8500-27A
Hanford sediment
B8500-29A
Hanford sediment
B8500-31A
Hanford sediment
B8500-32A
Hanford sediment
B8500-34A
Hanford sediment
B8500-35A
Hanford Trench
AE-3 Sediment
Hanford Trench
AE-3 Sediment
Hanford Trench
AE-3 Sediment
Hanford Trench
AE-3 Sediment
Hanford Trench
AE-3 Sediment
Reference
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998a)
Kaplan et al.
(1998b)
Kaplan et al.
(1998b)
Kaplan et al.
(1998b)
Kaplan et al.
(1998b)
Kaplan et al.
(1998b)
Bird and Schwartz (1996) studied effect of sediment-to-water ratios, oxic and anoxic conditions,
alkalinity, microbial activity, and time on the partitioning of iodine, as 125I, on representative lake
sediments from the Canadian Shield. The Kd values were determined using the shaking, batch
equilibration technique with sediment-to-solution ratios of 1, 10, and 50 percent for organic
sediment, peaty sediment, clayey silt/sand sediment, and sand sediment. Bird and Schwartz (1996)
found that the alkalinity of water, which ranged from < 1.0 to 250 mg/1 as CaCO3, had only a minor
effect on Kd with values changing less than a factor of two. Treating organic soil with a fungicide,
bactericide, irradiation or heat decreased Kd values by a factor of 1.1, 2.3, 7.5, and 22, respectively.
These results suggested that microbial processes were important on the sorption of 125I to sediment,
although chemical and/or physical changes, as results of radiation or heat treatment, may also be
important. Under anoxic conditions, Kd values were low (generally less than 1 ml/g. Over a 48-day
period, Kd values increased from about 690 ml/g on day 6 to 3,840 ml/g on day 48 under oxic
conditions, and from about 0.5 to 16 ml/g under anoxic conditions. Most of the iodine adsorbed to
sediment under oxic conditions was released back into the water under anoxic conditions.
Geometric mean Kd values for iodine measured in untreated river water under oxic conditions, after
shaking for 48 hours, are listed in the Table 5.13. Lower Kd values were observed for clayey
5.39
-------
silt/sand sediment and sand sediment than for the high organic sediment and peaty sediment. The
measured Kd values tended to decrease with increasing sediment-to-water ratio, except for the sand
sediment. Under anoxic conditions, the Kd for iodine is much lower (approximately 1 ml/g). The
Kd values for sand sediment were low under both oxic and anoxic conditions.
Table 5.12. Median Kd values (ml/g) for I" and I0g measured
in deionized water (Yoshida et a/., 1998). [Median total organic
carbon (TOC) and median Kd values listed for volcanogenous
regosol and peat soils are mean and single-sample values,
respectively. Characterization information for soil types are
listed in Table 2 in (Yoshida etal., 1998).]
Soil Type
Andosols
Gray Lowland Soils
(Dystric Fluvisols)
Red and Yellow
Soils
(Orthic Acrisols)
Gley Soils
(Dystric Gleysols)
Brown Forrest Soils
(Hurnic Cambisols)
Sand-Dune
Regosols
(Arensols)
Volcanogenous
Regosols
(Vitric Andosols)
Peat Soils
(Dystric Histosols)
Others
n
20
14
10
5
3
3
2
1
10
Median
TOC
(g/kg)
63.4
15.7
10.8
21.0
54.0
0.6
9.9
213
16.1
Median Kd for
Iodide (I") (ml/g)
Wet Soil
1,610
413
209
1,540
1,290
8.5
3,180
5,590
230
Soil That
Was
Autoclaved
4.3
2.0
1.6
22
4.8
0.7
17.1
11.4
1.3
Median Kd for
lodate (IO3) (ml/g)
Wet Soil
1,480
412
217
1,190
1,120
7.9
1,040
5,650
133
Soil That
Was
Autoclaved
19.5
7.2
19.7
9.0
48.9
3.0
383
13.4
10.7
5.40
-------
Table 5.13. Geometric mean Kd values (ml/g) for iodine
measured in untreated river water under oxic conditions
(Bird and Schwartz, 1996)
Sediment Type
High Organic
Peaty
Clayey Silt/ Sand
Sand
Sedinient-to-Water Ratio (Percent)
1
247
104
32
0.1
10
109
68
15
0.5
50
21
14
2.9
0.2
Fukui et al. (1996) used batch equilibration experiments to study the sorption of I" and IO3 on a
sandy loam soil dried at 100°C and mixtures of that soil with various types of organic matter. The
experiments were conducted with three pH 6.5 groundwater solutions, a pH 8.6, 0.5xlO~3 mol/1
CaCO3 solution, and a pH 7.1 pond water. The groundwater solutions had essentially the same
composition except that one (GW1) was used as sampled, the second (GW2) was the same
groundwater solution after having N2 bubbled through for continuously for one month, and the
third (GW3) was the same groundwater solution after having been allowed to stand in air for
one month without N2 bubbling through it. Fukui et al. (1996) proposed multiple sorption
mechanisms for I", and a simple anion exchange process for IO3. To study the effect of pH on the
sorption of I" and IO3, batch experiments were conducted with a 125I-spiked CaCO3 solution; the
initial pH of the solution was adjusted over the range of 3.2 to 10.6 using 1 mol/1 HC1 or NaOH
solutions. After the seven-day equilibration period, the pH ranged from 6.1 to 7.1 due to the buffer
capacity of the soil. In the initial pH range of 3.2 to 10.6, the uptake of both iodine species
decreased with pH, especially in the weakly acidic region. Fukui et al. (1996) indicated that for the
experiments at the initial pH 3.2, approximately 60 percent of the I" species were converted to
molecular I2 and approximately 30 percent of the IO3 species to I". The Kd values for both iodine
species on insoluble humic substances were one order of magnitude larger than those for soil. The
average Kd values for I" after 14 days of contact time in the five solutions were 3.0, 1.9, 1.4, 3.3, and
6.4 ml/g, respectively, for the GW1, GW2, GW3, the CaCO3 solution, and the pond water. The
average Kd values for IO3 for the same solutions were 11, 8.3, 10, 12, and 10 ml/g, respectively.
Serne et al. (1993) used the batch adsorption technique to measure the Kd values for I" on
three characterized sediments from the Hanford Site in southeastern Washington. The sediment
samples included two loamy sands (samples TBS-1 and Trench-8) and one sand (sample CGS-1).
The measurements were conducted using an uncontaminated groundwater (pH 8.14) sample from
the Hanford site. After 35 days of contact time, the pH values for the CGS-1, TBS-1, and Trench-8
sediment suspensions were 7.9-8.4., 8.0-8.4, and 8.22, respectively. The Kd values measured for I"
were 1.2 ml/g for CGS-1 sediment at 35 days, 2.6 ml/g for TBS-1 sediment at 35 days, and 1.5 ml/g
for Trench-8 sediment at 44 days.
5.41
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Muramatsu et al. (1990) used the batch equilibration method to study the effects of heating and
gamma-irradiation of soil relative to the sorption of I" and IO3 on soil and selected soil components
in deionized water. The Kd values were measured for I" and IO3 on soils identified as a field soil
(andosol),1 rice paddy soil, and sandy soil, and on bentonite, kaolinite, quartz sand, Fe2O3, A12O3,
and humic acid. The organic carbon contents of the dry soils were 4.4, 2.4, and 1.4 percent for the
field, rice paddy, and sandy soils, respectively. The soil pH values were 5.4, 5.7, and 5.5, for the
field, rice paddy, and sandy soils, respectively. The sorption measurements were completed on the
untreated samples of each soil type as well as samples of each soil that were dried at several
temperatures (in air at ambient conditions and at 100, 150, 200, and 300°C) and that were gamma-
irradiated. Desorption experiments were also conducted at 25 and 95°C using the field soil and
several leachates, such as deionized water, NaOH, K2SO3, HC1, and others. The Kd values measured
for the untreated soils are listed in Table 5.14. The reader is referred to Table III in Muramatsu et al.
(1990) for the Kd values for the soils with different drying treatments. They noted that the measured
Kd values probably represented a mixture of the sorption of I" and IO3, because part of the initial
concentration of IO3 probably reduced to I" during the course of the measurements. Sorption of I"
and IO3 on the untreated field soil was considerably greater than their sorption on the sandy soil.
These results indicated that the sorbed iodine was associated with the organic matter in the soil
samples (Muramatsu et al., 1990). The Kd values decreased with increasing temperature of the soil
drying treatments. Iodine did not sorb significantly on quartz sand, bentonite, kaolinite, or A12O3
under these conditions. Significant sorption was measured for both I" and IO3 on Fe2O3 and humic
acid reagent used in these experiments.
Sheppard and Thibault (1988) studied the vertical migration of iodine in two types of mires typical
of the Precambrian Shield in Canada. Sheppard and Thibault (1988) derived in situ Kd values from
analyses of the dried peat and pore water. The Kd values determined for iodine ranged from 0.2 to
64 ml/g. The investigators proposed that the mobility of iodine was related to the aeration status
and botanical origin of the peat.
Gee and Campbell (1980) used batch equilibration methods to measure Kd values for 125I on a sandy
sediment from the Hanford Site in southeastern Washington. The iodine sorption experiments were
conducted with six synthetic groundwater solutions as well as with low and high nitrate salt
solutions. The 125I Kd values for the sand sediment with the groundwater solutions ranged from 4 to
15 ml/g. The Kd values for the salt solutions were 18 and 4 ml/g for the high and low salt
solutions, respectively. The addition of ethylenediamine tetraacetic acid (EDTA) to the 125I-spiked
solution had the apparent effect of decreasing the Kd values for the measurements made with the
high and low salt solutions.
Wildung et al. (1974) used the batch equilibration method to measure the Kd values for I" and methyl
iodide on 22 types of soils collected in Oregon, Washington, and Minnesota. The Kd values for I"
and methyl iodide range from 0.8 to 52.6 ml/g, and 0.1 to 3.1 ml/g, respectively. The sorption of I"
was determined to be correlated only to the content of silt-size particles in these soils. Wildung et al.
Andosol (or Andept) is defined as a black or dark brown soil that is formed from volcanic material. It is
characterized as being high in organic material, low in exchangeable bases, and high in exchangeable aluminum (Bates
and Jackson, 1980).
5.42
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(1974) found that the sorption of methyl iodide on these soils was positively correlated to clay-size
soil particles, organic carbon, and cation exchange capacity, and negatively correlated to pH.
Table 5.14. Values of Kd (ml/g) measured by Muramatsu
et al. (1990) for the sorption of I" and I0g on soil and
selected soil components of an andosol.
Soil
Identification
Field Soil
Rice Paddy Soil
Sandy Soil
Field Soil
Rice Paddy Soil
Sandy Soil
Kd
(ml/g)
7,500 for I
560 for I
35 for T
7,000 for IO3
430 for IO3
32 for IO3
CEC
(meq/100 g)
20
10
3.9
20
10
3.9
Solution
Deionized Water
Deionized Water
Deionized Water
Deionized Water
Deionized Water
Deionized Water
5.5.5.4- Published Compilations Containing Kj Values for Iodine
Because the references in this section are often cited or used for comparison in other publications,
the following summaries are provided for completeness. It is recommended that the reader review
the original reference and the references cited therein to understand the procedures and sources of
the Kd values used for each compilation. The compilations do not distinguish between oxidation
states for those contaminants that are redox sensitive or consider other important factors that
contribute to variability in sorption, such as pH. Moreover, in cases where very large Kd values are
listed, there is a risk that the original Kd measurement may have included precipitated components.
Looney et al. (1987) tabulated estimates for geochemical parameters needed for environmental
assessments of waste sites at DOE's Savannah River Plant in South Carolina. Looney et al. list Kd
values for several metal and radionuclide contaminants based on values that they found in 1-5
published sources. For iodine, Looney et al. list a "recommended" Kd of 0.2 ml/g, and a range from
0.001 to 1 ml/g. Looney et al. note that their recommended values are specific to the Savannah
River Plant site, and they must be carefully reviewed and evaluated prior to use in assessments at
other sites.
Thibault et al. (1990) (also see Sheppard and Thibault, 1990) present a compilation of soil Kd values
prepared to support radionuclide migration assessments for a Canadian geologic repository for spent
nuclear fuel in Precambrian Shield plutonic rock. Thibault et al. collected Kd values from other
compilations, journal articles, and government laboratory reports for important elements, such as
iodine, that would be present in the nuclear fuel waste inventory. The iodine K^ values listed in
Thibault et al. (1990) are included in Table 5.15. Thibault et al. (1990) describe the statistical methods
used for analysis of the compiled Kd values. The range for the Kd values used to calculate the
5.43
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"geometric mean" cover several orders of magnitude. Readers are cautioned against using
"geometric mean" values or any other form of averaged Kd values as "default" Kd values, as such
values are usually calculated from data compiled from different investigators for different soil
systems. These mean or average values do not represent any particular environmental system and
geochemical conditions. As discussed in Volume I (EPA, 1999b), the variation observed in the
literature for Kd values for a contaminant is due to differences in sorption mechanisms, geochemical
conditions, soil materials, and methods used for the measurements.
Table 5.15. Iodine Kd values (ml/g) listed in Thibault et a/.
(1990, Tables 4 to 8).
Soil Type
Sand
Silt
Clay
Organic
Kd Values (ml/g)
Geometric
Mean
1
5
1
25
Number of
Observations
22
33
8
9
Range
0.04 - 81
0.1 -43
0.2 - 29
1.4 - 368
McKinley and Scholtis (1993) compare radionuclide Kd sorption databases used by different
international organizations for performance assessments of repositories for radioactive wastes. The
iodine Kd values listed in McKinley and Scholtis (1993, Tables 1, 2, and 4) are listed in Table 5.16.
The reader should refer to sources cited in McKinley and Scholtis (1993) for details regarding their
source, derivation, and measurement. Radionuclide Kd values listed for cementitious environments
in McKinley and Scholtis (1993, Table 3) are not included in Table 5.16. The organizations listed in
the tables include: AECL, GSF, IAEA, KBS, NAGRA, NIREX, NRG, NRPB, PAGIS (CEC), PSE,
RIVM, SKI, TVO, and UK DoE (acronyms defined in Section A. 1.0 in Appendix A).
5.5.5.5 Kj Studies of Iodine on Pure Mineral, Oxide, and Crushed Rock Materials
Numerous adsorption studies have been conducted of iodine on pure minerals, oxide phases, and
other geologic-related materials. The Kd values listed in these studies are not necessarily relevant to
the mobility and sorption of iodine in soils. However, they are listed in Appendix F for
completeness. The references cited in Appendix F are listed in the main reference list in Chapter 6.
The potential value of the references that they cite and the sorption processes that they discuss is left
to the reader. Many of these studies were conducted because of extensive research interest in
developing getters (adsorbents) that could be added to waste streams and tailored barriers for
removal and/or immobilization of dissolved iodine. The studies of iodine sorption on crushed rock
were conducted typically as part of national research programs to investigate the feasibility of
geological disposal of high-level radioactive waste (HLW).
5.44
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Table 5.16. Iodine Kd values (ml/g) listed by McKinley and
Scholtis (1993, Tables 1, 2, and 4) from sorption databases
used by different international organizations for performance
assessments of repositories for radioactive wastes.
Organization
AECL
GSF
IAEA
KBS-3
NAGRA
NIREX
NRC
NRPB
PAGIS
PSE
RIVM
SKI
TVO
UK DoE
Argillaceous (Clay)
Sorbing
Material
Bentonite-Sand
Sediment
Pelagic Clay
Bentonite
Bentonite
Clay
Clay Muds tone
Clay, Soil Shale
Clay
Sub seabed
Sediment
Sandy Clay
Bentonite
Baltic Sea
Sediment
Lake Sediment
Clay
Coastal Marine
Water
Kd
(ml/g)
0.0085
0.5
200
13
5
5
0
0
0
100
0
1
1
100
100
10
20
Crystalline Rock
Sorbing
Material
Granite
Granite
Granite
Granite
Basalt
Tuff
Granite
Crystalline
Rock, Reducing
Crystalline Rock
Kd
(ml/g)
0
0
1
0
0
0
0
0.5
0.8
Soil/Surface Sediment
Sorbing
Material
Soil/ Sediment
Soil/ Sediment
Soil/Sediment
Soil/Sediment
Soil/ Sediment
Soil/ Sediment
Kd
(ml/g)
10
1
10
0
100
0.8
5.6 Neptunium Geochemistry and Kd Values
5.6.1 Overview: Important Aqueous- and Solid-Phase Parameters Controlling
Retardation
Neptunium is a transuranic element that may exist in several valence states. Neptunium(V) (Np(V))
and Np(IV) are the most important valence states in natural waters. Over the pH range of most
natural waters, Np(V) is present primarily as the cation NpOj. Neptunium(V) is considered
relatively mobile, and Np(V) solids are quite soluble. If the concentrations of dissolved Np(V) are
5.45
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sufficiently high, the solubility of Np(V) may be controlled by solids such as NpO2OH,
NaNpO2CO3, Na3NpO2(CO3)2, and KNpO2CO3. Neptunium (IV) exists in reducing (i.e., low Eh)
groundwater systems, and is not considered very mobile, because it forms sparingly soluble oxide
and hydroxide solids. Reduction of Np(V) to Np(IV) can occur by a variety of abiotic and biotic
processes.
Neptunium(V) aqueous species sorb to some extent to iron oxides and clays, but do not sorb to a
major degree on most common minerals. Because NpO2 does not compete favorably with
dissolved Ca2+ and other divalent ions for adsorption sites on soils, the Kd values for Np(V) are
relatively low. The adsorption of Np(V) has a strong dependence on pH. Typically, the sorption of
Np(V) on minerals is negligible at pH values less than 5, and increases rapidly at pH values between
5 to 7. This pH-dependency is expected for ions present in solution primarily as cations. In
carbonate-containing solutions, the adsorption of Np(V) on iron oxides has been observed to
decrease at pH values greater than 7 to 9 in response to the formation of aqueous Np(VI) carbonate
complexes.
5.6.2 General Geochemistry
Neptunium [Np, atomic number (Z) = 93] has 19 isotopes (Tuli, 2000). The atomic masses of these
isotopes range from 225 to 244. Neptunium-237 (237Np) has a half live (t,/2) of 2.14xl06 years and is
the most important neptunium isotope from an environmental perspective Potential sources of
237Np include fallout from nuclear weapons, effluent cooling water from fission reactors, industrial
processing of 237Np produced in fission reactors, and 237Np present as a component of high-level
nuclear wastes (Thompson, 1982). Although spent nuclear fuel contains a relatively small initial
concentration of 237Np, its concentration increases with time from the radioactive decay of Am-241
(t,/2 = 432.2 years). Because of its long half-life, 237Np will be a major contributor to the radiation
inventory of nuclear waste stored in geologic repositories after approximately 100,000 years
(Kaszuba and Runde, 1999).
Neptunium is a transuranic (actinide) element. Neptunium may exist in the +3, +4, +5, +6, and +7
valence states, but only the +4, +5, and possibly +6 states are relevant to natural environments.
Neptunium(VI) is stable only in highly oxidizing solutions and is therefore not important under
most environmental conditions. As noted in Volume II (EPA, 1999c), aqueous speciation and
solubility reactions, and sorption onto minerals and soils usually differ considerably for redox-
sensitive elements contaminants in their different valence states. Neptunium(V) is considered
relatively mobile, because Np(V) solids are quite soluble and Np(V) aqueous species do not readily
sorb on common minerals. Over the pH range of most natural waters, Np(V) will be present
primarily as the neptunyl ion, NpO2. Neptunium(IV) exists in reducing (i.e., low Eh) groundwater
systems, and like U(IV) and Pu(IV), may form sparingly soluble oxide and hydroxide solids that limit
the mobility of Np(IV) under reducing conditions. Redox-sensitive elements, such as neptunium,
can be immobilized by surface-mediated, heterogeneous reduction/sorption reactions on
Fe(II)-containing oxide and silicate minerals. Many of these minerals exist as coatings on soil
particles and/or primary constituents of soils. The heterogeneous electrochemical reactions occur
by electron transfer reactions by which the Fe(II) is oxidized to Fe(III), and the redox-sensitive
contaminant, is reduced to a lower valence state, such as reduction of Np(V) to Np(IV). These
reactions have recently been the subject of considerable interest because they may have a significant
effect on the mobility of redox-sensitive elements in the vadose zone, aquatic sediments, and
5.46
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groundwater. They are the basis for certain remediation technologies, such as permeable barriers
composed of zero-valent iron particles (i.e., as metallic iron) or sodium-dithionite reduced soils,
which are currently being tested for immobilization of groundwater contaminants.
Surface-mediated, heterogeneous reduction/sorption reactions are reviewed in detail by White
(1990), and studied experimentally by White and Yee (1985), White et al. (1994), and many others.
Experimental studies include investigations of Fe(II)-containing silicate minerals and/or rocks in the
reduction of: Np(V) (Hakanen and Lmdberg, 1991; Susak et al., 1983; Meyer et al., 1984); Fe(III)
(White and Peterson, 1996; White and Yee, 1985); Cr(VI) (White and Peterson, 1996; Eary and Rai,
1989, Ilton and Veblen, 1994); Cu(II) (White and Peterson, 1996); nitrate (Postma, 1990); and V(V)
(White and Peterson, 1996). These surface-mediated, heterogeneous reduction/sorption reactions
have also been observed in sorption experiments conducted with crushed rock materials. For
example, Bondietti and Francis (1979) showed that Np(V) could be reduced to less soluble Np(IV)
by Fe(II) in minerals in igneous rocks.
The reduction of Np(V) to Np(IV) by biotic processes has also been demonstrated. Lloyd
(2000a) conducted a biochemical study of the reduction of Np(V). Their experiments demonstrated
that Shewanella putrefadens reduced Np(V) to a lower valence state, possibly Np(IV). However, this
reduction process was not sufficient to remove neptunium from solution. Lloyd et al. (2000a) were
able to remove 237Np and its daughter product of protactinium-233 (233Pa) from solution by
bioprecipitation using a combination of the two organisms S. putrefadens and Citrobacter sp. The
bioprecipitation resulted from bioreduction to Np(IV) by S. putrefadens in concert with phosphate
liberation by the Citrobacter sp from the glycerol 2-phosphate solution.
The environmental chemistry and mobility of neptunium in surface water, groundwater, and
geologic environments has been reviewed extensively by others, such as Silva and Nitsche (1995),
Tanaka et al. (1992), Lieser and Muhlenweg (1988); Coughtrey et al. (1984), Thompson (1982),
Onishi et al. (1981), and Ames and Rai (1978). Although somewhat dated, the extensive review by
Coughtrey et al. (1984) is particularly noteworthy.
5.6.3 Aqueous Speciation
Lemire et al. (2001) have published an extensive, detailed review of the chemical thermodynamics of
neptunium. The thermodynamic data for neptunium aqueous species and solids are limited and not
well known relative to other radionuclides. Figure 5.7 is an Eh-pH diagram that shows the
dominant aqueous species of neptunium as a function of pH and Eh (redox potential) at 25°C with
respect to the thermodynamic stability of water.
Possible Np(V) aqueous species include those listed in Table 5.17.1 The distribution of aqueous
species for Np(V) (see Figure 5.8) as a function of pH was calculated for an oxidizing environment
containing the water composition listed in Table 5.1. The speciation calculations indicate that the
neptunyl ion, NpO2+, is the dominant aqueous species for most environmentally-relevant pH
conditions. The aqueous speciation of neptunium is dominated by NpO2 over a wide range of
Lemire et al. (2001) was published after the writing of this report had been completed. Therefore, the Np(V) species
and associated thermodynamic constants from Lemire et aL (2001) were not included in the speciation calculations listed
in this section.
5.47
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environmental pH and Eh conditions. At pH conditions greater than about 8.5, aqueous Np(V)
complexes, such as NpO2(OH)° (aq) and NpO2(CO3)~ may be more dominate.
Figure 5.7. Eh-pH stability diagram for the dominant neptunium
aqueous species at 25°C. [Diagram based on a total concentration
of 10"8 mol/l dissolved neptunium and 10"3 mol/l dissolved carbonate.]
1
.5
£T
g
_c
0
-.5
C
""•*.^
• ">">'»- •
Np02+ ^\^NP02(C03)23-
^^\Np02(C03)35'
\ NP Np02C03 \\\
& j^OH3+ "^s^ \ \ "-
. \ ^ Np(OH)22+ ^ ^ ^ .
-M Xr /Np(OH)3+
X i J^r J-.^
"•- k. N "^~
^VN
Np3""^--. Np(OH)4°(aq) \ -
^~^^ \
~^~ >•
**^^
**^^
) 2 4 6 8 10 12 1
PH
4
Several recent studies have been published regarding the aqueous speciation of Np(V). For
example, the hydrolysis and complexation of Np(V) has been studied by Runde eta!. (1996) and
Fanghanel eta!. (1995). Increasing carbonate at high pH increases the importance of carbonate
complexes such as NpO2(CO3)3~. Neck eta!. (1994, 1997) has studied the complexation reactions of
Np(V) in carbonate solutions. In strongly alkaline NaOH)/Na2CO3 solutions, Neck eta!. (1997)
found that the mixed Np(V) hydroxo-carbonate complexes NpO2(OH)(CO3)2~ and
NpO2(OH)2(CO3)3~ may form. Morgenstern and Kim (1996) used absorption spectroscopy to
confirm the formation of the Np(V) phosphate complexes NpO2PO^ and NpO2HPO4.
Ultratfiltration and centrifugation studies conducted with colloidal iron by Itagaki eta!. (1991)
suggest that Np(V) pseudocolloids can be formed in the pH conditions where the aqueous species
NpO2OH° (aq) is dominant. Under reducing conditions, precipitation of amorphous NpO2-xH2O
and its colloids were found to be important to the environmental behavior of neptunium (Itagaki et
a!., 1991).
5.48
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Table 5.17. Np(V) aqueous species.
Aqueous Species
NpO+, NpO2OH° (aq), NpO2(OH)2
NpO2CO3, NpO2(CO3)^, NpO2(CO3)35-,
NpO2(OH)(CO3)^, NpO2(OH)2(CO3)3-.
NpO2H2PO°4 (aq), NpO2HPO4, NpO2SO4
NpO2Cl° (aq), NpO2F° (aq)
Figure 5.8. Calculated aqueous speciation for Np(V) as
a function of pH. [Neptunium(V) aqueous speciation was
calculated based on a total concentration of dissolved neptunium
of 1x10"8 mol/l, the water com position in Table 5.1, and
thermodynamic constants from Lemire et al. (1984).]
100
•2 80
-t 60
Q
-£ 40
0
I 20
0
Np02(C03)
Np02(OH)° (aq)
Np02(C03)
The aqueous speciation of neptunium in the presence of humic substances in natural waters
conditions has been reviewed by Moulin et al. (1992). The complexation and reduction of Np(V)
with humic substances has also been investigated by various laboratories with different experimental
methods (Artinger et al., 2000; Sakamoto et al., 2000; Zeh et al., 1999; Marquardt and Kim, 1998,
1996; Rao and Choppm,1995; Choppm, 1992; Kim and Sekme, 1991). Zeh et al. (1999) used
absorption spectrometry and ultrafiltration in the neutral pH range to investigate the reduction of
Np(V) to Np(IV) in a groundwater containing humic substances. The data suggest that Np(IV) may
form a humic colloid-born species that can be separated from water by ultrafiltration. The original
dissolved Np(V) was present as NpO2, NpO2CO3, and/or Np(V)-humate depending on pH. If the
reduction of Np(V) to Np(IV) takes place by the formation of humic colloid-born species of
5.49
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Np(IV), the colloids may remain stable in groundwater and mobile in porous aquifer systems (Zeh et
al., 1999).
Marquardt and Kim (1998) investigated the humate complexation of NpO2 over the pH range 7 to
9 with purified humic acid extracted from a groundwater from Gorleben in northern Germany.
Their results indicated that the only reaction competitive with Np(V) -humate complexation was the
formation of Np(V) complexes with dissolved carbonate. Their calculations suggested that humate
complexation of Np(V) may play an important role in groundwater with relatively high humic acid
concentrations. However, Marquardt and Kim (1998) noted that the reduction of Np(V) to Np(IV)
by interaction with humic acid was another important reaction to be considered with respect to the
environmental behavior of neptunium. Marquardt et al. (1996) studied the influence of neptunium
and humic acid concentrations, pH value, ionic strength, temperature, and competing cations on the
complexation of Np(V) with humic acids. The stability constant for the complex was found to
increase with increasing concentration of Np(V). No effect of ionic strength between 0.001 and 0.1
M was observed, and neptunium-humate complexation decreased with increasing temperature.
Calcium and magnesium competed with neptunium for the available humic ligands. Under
anaerobic conditions and at very low metal concentrations, humic acid reduces Np(V) to Np(IV),
and the rate of reduction was dependent on pH. Rao and Choppin (1995) used near-IR (infrared)
absorption and two sources of humic acid to determine the binding constants for Np(V)-humate
complexes in the pH range 4.5 to 7.5. Their results indicated that Np(V) forms a single type of
complex with humate, probably a 1:1 Np(V)-carboxylate complex.
5.6.4 Dissolution/Precipitation/Coprecipitation
The solubility of Np(V) has been studied extensively for the purpose of estimating the maximum
solubility concentrations of dissolved neptunium that might be released from a geologic repository
for high level waste (HLW) with subsequent migration in groundwater systems (e.g., Novak and
Roberts, 1995; Neck et al., 1994; Lemire, 1984; and others). In carbonate-free aqueous solutions
with OH" as the only complexing ligand, the maximum concentration of dissolved Np(V) is likely
determined by the solubility product of solids, such as Np2O5'xH2O (Efurd et al., 1998) or solid
NpO2OH (Al Mahamid et al., 1998; Roberts et al., 1996). In carbonate-rich solutions, a variety of
solids, such as hydrated NaNpO2CO3 (Neck et al., 1994; Lemire et al., 1993), Na3NpO2(CO3)2
(Al Mahamid et al., 1998; Neck et al., 1994; Lemire et al., 1993), and KNpO2CO3 (Al Mahamid et al.,
1998; Lemire et al., 1993), have been studied as possible solubility controls for the maximum
concentrations of dissolved Np(V) under oxidizing conditions. Neck et al. (1994) studied the
solid/liquid reactions of Np(V) in carbonate solutions at different ionic strength. At the conditions
of his experiments, he found that NaNpO2CO3 and Na3NpO2(CO3)2 formed at pH values less than
8.5 and greater than 9, respectively.
Solids such as Np(IV) hydrous oxide (Nakayama et al., 1996; Rai and Ryan, 1985); amorphous
NpO2-xH2O (Rai et al., 1987), and amorphous NpO2 (Rai et al., 1999) have been studied as possible
solubility controls for Np(IV).
5.50
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5.6.5 Adsorption/Desorption
5.6.5.1 Guidance for Screening Calculations ofA-dsorbtion
Neptunium(V) aqueous species sorb to some extent to iron oxides and clays, but do not sorb to a
major degree on most common minerals. Therefore, dissolved Np(V) is considered to be relatively
mobile in soil systems. Because NpOj does not compete favorably with dissolved Ca2+ and other
divalent ions for adsorption sites on soils, the Kd values for Np(V) are relatively low (Kaplan and
Serne, 2000). Coughtrey et al. (1984) review sorption studies published prior to 1984 for neptunium
to soils and crushed rock materials. Some of these studies and more recent neptunium adsorption
studies are described in the sections below.
The limited number of Kd studies identified during this review for the adsorption of Np(V) on soil
precluded derivation of Kd look-up tables of conservative minimum and maximum Kd values for
Np(V). Of the studies discussed below, only Routson et al. (1975, 1977), Serne et al. (1993), and
Kaplan et al. (1996) report Kd values and corresponding pH values for Np(V) adsorption on soil.
The majority of their values are limited to the pH range from 6.2 to 8.5, and only Routson et al.
(1975, 1977) reports Np(V) Kd values for pH values less than 8.5. For the pH range from 4 to 10, it
is suggested that a Kd of 0.2 ml/g be used as a minimum Kd value for screening calculations of
americium migration in soils. This value was reported for pH 7.8 by Routson et al. (1975, 1977) and
is the lowest Kd value that they gave for experiments conducted with very to moderately dilute
calcium and sodium electrolyte solutions. The other Kd values reported by Routson et al. (1975,
1977) for these solution concentrations ranged from 0.36 ml/g at pH 7.4 to 3.51 at pH 4.1. It is
likely that the conservative minimum Kd for Np(V) in the pH range from 4 to 10 is greater than this
limiting value, but no additional Np(V) Kd studies substantiate this assumption. As noted
previously, it is recommended that for site-specific calculations, partition coefficient values measured
at site-specific conditions are absolutely essential.
Given the limited availability of Kd studies for the adsorption of Np(V) on soil, readers may want to
consider using geochemical models [see Section 5 in Volume I (EPA, 1999b)] to estimate the mass
of adsorbed Np(V). Kohler et al. (1999) and Girvin et al. (1991) have determined the neptunium
surface complexation reactions and associated intrinsic constants for modeling Np(V) adsorption on
hematite and amorphous iron oxyhydroxide, respectively. If soil characterization studies indicate
that either of these iron oxides is an important constituent for soils at the site under study,
geochemical models can be used to calculate the mass of Np(V) adsorbed on these oxides, and
based on the iron oxide concentration in the site soils, estimate a Kd for Np(V).
5.6.5.2 General .Adsorption Studies
Experimental studies indicate that the adsorption of Np(V) has a strong dependence on pH,
especially for iron oxides (Kohler eta!., 1999; Girvin eta!., 1991; Allard, 1984). Typically, the
sorption of Np(V) on minerals is negligible at pH values less than 5, and increases rapidly when pH
is between 5 to 7. This pH-dependency is expected for ions that are present in solution primarily as
cations, such as NpOj (EPA, 1999b). In carbonate-containing solutions, the adsorption of Np(V)
on iron oxides has been observed to decrease at pH values greater than 7 to 9 in response to the
formation of aqueous Np(VI) carbonate complexes (Kohler eta!., 1999). This latter behavior is
5.51
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analogous to that observed for the adsorption of U(VI) in carbonate-solutions at alkaline pH values
[see Section 5.11 in Volume II (EPA, 1999c)].
Table 5.18. Neptunium(V) h^ values (ml/g) measured for
three sediments by Kaplan et al. (1996).
Sediment Sample
Touchet Bed Sand (sample TBS-1)
Silty Loam (Sample Trench AE-3)
Very Coarse Sand (Sample Trench-94)
Kd Values (ml/g)
7 days
2.17
2.67
14.17
77 days
3.62
13.48
19.86
5.6.5.3 Kj Studies for Neptunium on Soil Materials
Serne et aL (1993) measured Kd values for Np(V) in groundwater contacting the <2-mm size fraction
of Trench-8 sediment from the DOE's Hanford Site in eastern Washington. The pH of the
Hanford groundwater was 8.3. The particle size distribution for the in .rrf//Trench-8 sediment
contained 9.7 percent gravel, 78.6 percent sand, 6.3 percent silt, and 5.4 percent clay. The <2-mm
size fraction of this sediment contained 87 percent sand, 7 percent silt, and 6 percent clay. The Kd
values measured at 5 and 44 days were 13.5+3 and 29.1+3.6 ml/g, respectively.
Kaplan et al. (1996) used the batch technique to measure the Kd values for Np(V) on three sediment
samples from the Hanford Site in southeastern Washington. The measurements were conducted
with a groundwater (pH 8.46) taken from a well located in an uncontaminated area of the Hanford
Site. The sediment samples included a Touchet bed sand (sample TBS-1), a silty loam (sample
Trench AE-3), and a very coarse sand (sample Trench-94). Kaplan et al. (1996) observed that the Kd
values increased as a function of time. The Kd values measured at 7 and 77 days for the
three sediments are listed in Table 5.18.
Sheppard and Thibault (1988) studied the vertical migration of neptunium in two types of mires
typical of the Precambrian Shield in Canada. Sheppard and Thibault (1988) derived in situ Kd values
from analyses of the dried peat and pore water. The Kd values determined for neptunium ranged
from 31 to 2,600 ml/g. Neptunium was quickly immobilized in the reducing environment of the
mire, which was the cause for the large in situ neptunium Kd values.
Nishita et al. (1981) studied the extractability of 237Np(V) several types of soils as a function of pH.
The extractability of neptunium was considered to parallel the solubilization of insoluble aluminum,
iron, and/or manganese hydrous oxides, which are important adsorbents for dissolved contaminants
in soil systems (Nishita et al., 1981). The Np(V) Kd values and soil characteristics determined for
these soils by Nishita et al. (1981) are listed in Table 5.19. Nishita et al. (1981) concluded that Np(V)
will likely be more mobile than americium, curium, and plutonium in geologic systems.
5.52
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Table 5.19. Measured Np(V) Kd values (ml/I) and soil
properties for soils studied by Nishita et al. (1981).
Soil
Type
Silt Clay Loam
Sharpsburg
Sandy Loam
Malbis
Sandy Loam
Lyman
Silty Clay
Holts ville
Loam
Aiken
Silt Loam
Yolo
Muck
Egbert
pH1
5.9
5.3
5.0
7.8
6.0
6.7
7.2
OM2
Fraction
(%)
2.8
2.4
5.7
0.6
8.4
2.5
40.8
CEC
(meq/100
g)
20
15
15
30
15
25
60
Free Fe
Oxides
(%)
1.29
1.65
1.52
1.20
5.29
2.41
1.57
Mn3
Fraction
(%)
0.06
0.05
0.04
0.04
0.10
0.08
0.10
Extract
pH
5.83
6.85
4.08
5.57
4.42
6.06
7.29
8.28
5.56
6.57
6.13
6.83
6.24
7.25
Kd
(ml/g)
35
95
3
18
32
41
117
26
108
52
81
786
929
1 Saturated paste.
2 Organic matter.
3 In 4 M HNO3 extract.
Sheppard et al. (1979) examined the extent to which neptunium was sorbed by colloidal-size soil
particles which are potentially diffusible in soil/water systems. The batch equilibration experiments
were conducted with distilled water and 14 soils from Muscatine, Illinois; Hanford, Washington;
Barnwell, South Carolina; Idaho Falls, Idaho; and Paradise and Placerville, California.
Centrifugation measurements indicated that some 237Np was retained by the colloidal-size soil
particles. The sorption of 237Np on the soil particles was not complete after four to six months, and
the proportion of radionuclide retained by the colloidal-size fraction decreased with time. Sheppard
et al. (1979) found it difficult to correlate the results with the chemical and physical properties of the
soil samples. This was attributed to the lack of precise distribution ratios, competition with cation
exchange reactions, and complexation with humic and fulvic acid materials. Based on these results,
Sheppard et al. (1979) suggest that colloids of clay and humic acids are potentially important
processes for the transport of actinides in soil/water systems.
Dahlman et al. (1976) determined the Kd values for 237Np(V) in suspensions of <2 |im-clay particles
from a silt loam soil. At pH 6.5 in a 0.005 mol/1 solution of dissolved calcium, the
measured Kd value was 320 ml/g.
5.53
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Table 5.20. Properties of soils used in Kd measurements
by Routson etal. (1975, 1977).
Soil
Washington
South Carolina
CaCO3
(mg/g)
0.8
<0.2
Silt
(%)
10.1
3.6
Clay
(%)
0.5
37.2
CEC
(meq/100 g
)
4.9
2.5
pH
7.0
5.1
Table 5.21. Neptunium Kd values (ml/g) measured for
Washington and South Carolina soil samples in Ca(N03)2 and
NaN03 solutions by Routson etal. (1975, 1977).
Concentration of
Electrolyte (mol/1)
Kd Values (ml/g)
Washington Soil
South Carolina Soil
Ca(NO3)2 Solution
0.002
0.02
0.05
0.10
0.20
2.37
0.93
0.78
0.62
0.36
0.25
Not Determined
Not Determined
Not Determined
0.16
NaNO3 Solution
0.015
0.030
0.30
0.75
3.00
3.90
3.51
3.28
3.28
3.19
0.66
0.57
0.51
0.45
0.43
Routson eta!. (1975, 1977) used batch equilibration experiments to measure Kj values for 237Np on
two soils as a function of the concentrations of dissolved calcium and sodium. The soil samples
were selected to represent a range of weathering intensities. For arid conditions in the western
United States, sandy (coarse-textured), low-exchange capacity soil was selected from a low rainfall
area in eastern Washington. For humid conditions in the southeastern United States, a moderate-
exchange capacity soil was selected from South Carolina. Properties of the soils used for these
measurements are listed in Table 5.20. The Kd values were measured in 0.002, 0.02, 0.05, 0.10, and
5.54
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0.20 mol/1 Ca(NO3)2 solutions, and 0.015, 0.030, 0.30, 0.75, and 3.0 mol/1 NaNO3 solutions. The
pH values for selected samples of the 237Np solutions in the calcium and sodium systems were 7.4
and 4.1 for the Washington soil, and 7.8 and 6.2 for the South Carolina soil. The Kd values (Table
5.21) decreased with increasing concentrations of dissolved calcium and sodium.
For the solution concentrations used in these experiments, the Kd values for 237Np on the
Washington soil ranged from 2.37 to 0.36 ml/g as a function of dissolved calcium, and 3.90 to
3.19 ml/g as a function of dissolved sodium. For the South Carolina soil, the Kd values for 237Np
ranged from 0.25 to 0.16 ml/g as a function of dissolved calcium, and 0.66 to 0.43 ml/g as a
function of dissolved sodium.
5.6.5.4 Published Compilations Containing Kj Values for Neptunium
Because the references in this section are often cited or used for comparison in other publications,
the following summaries are provided for completeness. It is recommended that the reader review
the original reference and the references cited therein to understand the procedures and sources of
the Kd values used for each compilation. The compilations do not distinguish between oxidation
states of contaminants that are redox sensitive (e.g., neptunium) or consider other important factors
that contribute to variability in sorption, such as pH. Moreover, in cases where very large Kd values
are listed, there is a risk that the original Kd measurement may have included precipitated
components.
Baes and Sharp (1983) present a simple model developed for order-of-magnitude estimates for
leaching constants for solutes in agricultural soils. As part of this model development, they reviewed
and determined generic default values for input parameters, such as Kd. A literature review was
completed to evaluate appropriate distributions for Kd values for various solutes, including
neptunium. Because Baes and Sharp (1983) are cited frequently as a source of Kd values in other
published Kd reviews (e.g., Looney eta!., 1987; Sheppard and Thibault, 1990), the neptunium Kd
values listed by Baes and Sharp are reported here for completeness. Based on the distribution that
Baes and Sharp determined for the Kd values for cesium and strontium, they assumed a lognormal
distribution for the Kd values for all other elements in their compilation. Baes and Sharp listed an
estimated default Kd of 11 ml/g for neptunium based on 44 Kd values that ranged from 0.16
to 929 ml/g for agricultural soils and clays over the pH range 4.5 to 9.0. Their compiled Kd values
represent a diversity of soils, pure clays (other Kd values for pure minerals were excluded), extracting
solutions, measurement techniques, and experimental error.
Looney eta!. (1987) tabulated estimates for geochemical parameters needed for environmental
assessments of waste sites at DOE's Savannah River Plant in South Carolina. Looney eta!, list Kd
values for several metal and radionuclide contaminants based on values that they found in 1-5
published sources. For neptunium, Looney eta!, list a "recommended" Kd of 10 ml/g, and a range
from 0.1 to 1,000 ml/g. Looney eta!, note that their recommended values are specific to the
Savannah River Plant site, and they must be carefully reviewed and evaluated prior to use in
assessments at other sites.
5.55
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Table 5.22. Neptunium Kd values (ml/g) listed in
Thibault etal. (1990, Tables 4 to 8).
Soil Type
Sand
Silt
Clay
Organic
Kd Values (ml/g)
Geometric
Mean
5
25
55
1,200
Number of
Observations
16
11
4
3
Range
0.5 - 390
1.3-79
0.4 - 2,575
857 - 1,900
Thibault eta!. (1990) (also see Sheppard and Thibault, 1990) present a compilation of soil Kd values
prepared to support radionuclide migration assessments for a Canadian geologic repository for spent
nuclear fuel in Precambrian Shield plutonic rock. Thibault et al. collected Kd values from other
compilations, journal articles, and government laboratory reports for important elements, such as
neptunium, that would be present in the nuclear fuel waste inventory. The neptunium Kd values in
Thibault et al. (1990) are included in Table 5.22. Thibault et al. (1990) describe the statistical methods
used for analysis of the compiled Kd values. The range for the Kd values used to calculate the
"geometric mean" cover several orders of magnitude. Readers are cautioned against using
"geometric mean" values or any other form of averaged Kd values as "default" Kd values, as such
values are usually calculated from data compiled from different investigators for different soil
systems. These mean or average values do not represent any particular environmental system and
geochemical conditions. As discussed in Volume I (EPA, 1999b), the variation observed in the
literature for Kd values for a contaminant is due to differences in sorption mechanisms, geochemical
conditions, soil materials, and methods used for the measurements.
McKinley and Scholtis (1993) compare radionuclide Kd sorption databases used by different
international organizations for performance assessments of repositories for radioactive wastes. The
neptunium Kd values listed in McKinley and Scholtis (1993, Tables 1, 2, and 4) are listed in
Table 5.23. The reader should refer to sources cited in McKinley and Scholtis (1993) for details
regarding their source, derivation, and measurement. Radionuclide Kd values listed for cementitious
environments in McKinley and Scholtis (1993, Table 3) are not included in Table 5.23. The
organizations listed in the tables include: AECL, GSF, IAEA, KBS, NAGRA, NIREX, NRC,
NRPB, PAGIS (CEC), PSE, RIVM, SKI, TVO, and UK DoE (acronyms defined in Section A.1.0 in
Appendix A).
5.56
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Table 5.23. Neptunium Kd values (ml/g) listed by McKinley and
Scholtis (1993, Tables 1, 2, and 4) from sorption databases used
by different international organizations for performance
assessments of repositories for radioactive wastes.
Organization
AECL
GSF
IAEA
KBS-3
NAGRA
NIREX
NRC
NRPB
PAGIS
PAGIS SAFIR
PSE
RIVM
SKI
TVO
UK DoE
Argillaceous (Clay)
Sorbing
Material
Bentonite-Sand
Sediment
Pelagic Clay
Bentonite
Bentonite
Clay
Clay Muds tone
Clay, Soil Shale
Clay
Bentonite
Subseabed
Clay
Sediment
Sandy Clay
Bentonite
Bentonite
Baltic Sea
Sediment
Ocean Sediment
Lake Sediment
Clay
Coastal Marine
Water
Kd
(ml/g)
30
30
5,000
600
1,000
5,000
100
10
30
150
2,000
200
5
10
1,000
1,000
50,000
50,000
50,000
50
5,000
Crystalline Rock
Sorbing
Material
Granite
Granite
Granite
Granite
Basalt
Tuff
Granite
Crystalline
Rock, Reducing
Kd
(ml/g)
80
5,000
1,000
10
10
10
5,000
200
Soil/Surface Sediment
Sorbing
Material
Soil/ Sediment
Soil/ Sediment
Soil/ Sediment
Soil/ Sediment
Soil/ Sediment
Soil/ Sediment
Soil/ Sediment
Kd
(ml/g)
10
10
1,000
30
500
1,000
10
5.57
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5.6.5.5 Kj Studies of Neptunium on Pure Mineral, Oxide, and Crushed Rock Materials
Numerous adsorption studies have been conducted of neptunium on pure minerals, oxide phases,
and other geologic-related materials. The Kd values listed in these studies are not necessarily
relevant to the mobility and sorption of neptunium in soils. However, they are listed in Appendix G
for completeness. The references cited in Appendix G are listed in the main reference list in
Chapter 6. The potential value of the references that they cite and the sorption processes that they
discuss is left to the reader. Surface complexation modeling techniques have been used to
understand the mechanisms of neptunium adsorption (Kohler et al., 1999; Girvin et al., 1991; Fujita et
al., 1995). Kohler et al. (1999) and Girvin et al. (1991) derived the neptunium surface complexation
reactions and associated intrinsic constants to model Np(V) adsorption of hematite and amorphous
iron oxyhydroxide, respectively. The studies of neptunium sorption on crushed rock were
conducted typically as part of national research programs to investigate the feasibility of geological
disposal of high-level radioactive waste (HLW).
5.7 Radium Geochemistry and Kd Values
5.7.1 Overview: Important Aqueous- and Solid-Phase Parameters Controlling
Retardation
Radium (Ra) is an alkaline earth element, and can exist in nature only in the +2 oxidation state. In
the pH range of 3 to 10, the uncomplexed ion Ra2+ is the dominant aqueous species for dissolved
radium in natural waters. In sulfate-containing waters, precipitation and redissolution of calcium
(Ca), strontium (Sr), and barium (Ba) sulfates, rather than adsorption/desorption, could control the
concentrations of dissolved radium in the soil environment. Precipitation of radium is readily
possible as the solid-solution solids (Ba,Ra)SO4 and (metal, Ra)CO3 in waters where the
concentrations of dissolved sulfate and carbonate, respectively, are sufficiently high. This reaction,
as noted by some investigators, may also be the cause for some very high adsorption values reported
for radium in the literature. Very limited sorption data, especially Kd values, exist for radium on soils
and sediments. The adsorption behavior of radium will be similar to that of strontium. Relative to
other alkaline earth elements, radium is the most strongly sorbed by ion exchange on clay minerals.
The adsorption of radium is strongly dependent on ionic strength and concentrations of other
competing ions in that adsorption of radium decreases with increasing ionic strength. Radium is
also strongly adsorbed to mineral oxides present in soils, especially at near neutral and alkaline pH
conditions. The results of some studies also suggest that radium may be strongly adsorbed by
organic material in soils.
5.7.2 General Geochemistry
Radium [Ra, atomic number (Z) = 88] is an alkaline earth element, and can exist in nature only in the
+2 oxidation state. Radium and barium are adjacent to each other in the alkaline earth group. Due
to the similarity of their ionic radii (Table 5.24), the chemical behavior of radium is very similar to
that of barium.
Of the reported 34 radium isotopes, all are radioactive (Tuli, 2000). Only four radium isotopes are
found naturally. These include 226Ra [half life fe) = 1,600 y], 228Ra (t,A = 5.75 y), 224Ra (t% = 3.66 d),
and 223Ra (t% - 11.435 d). The radium isotopes 226Ra and 223Ra are intermediate radioactive decay
5.58
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products of the naturally occurring 238U and 235U decay series, respectively. 228Ra and 224Ra are decay
products of the naturally occurring 232Th decay series. The isotope 226Ra is generally assumed the
most important occurring radium isotopes in nature due to its long half life and the natural
abundance of 238U. However, King et al. (1982) found a relatively high proportion of 228Ra in a study
of groundwater geochemistry of 228Ra, 226Ra, and 222Rn in 150 groundwater samples from a wide
range of aquifer lithologies in South Carolina. The isotope 225Ra (tVi = 14.9 d) of the 237Np decay
chain has not been detected in nature (Molinari and Snodgrass, 1990). All of the other radium
isotopes have shorter half lives. The chemistry and radiochemistry of radium and the other elements
of the uranium and thorium natural decay series are reviewed by Molinari and Snodgrass (1990).
From the standpoint of health risk, 223Ra, 224Ra, and 226Ra are also important in that they decay to
produce radioactive isotopes of the noble gas radon that in turn decay by alpha particle emission [see
a discussion of the geochemical and sorption behavior of radon in Section 5.7 of EPA (1999c)].
Table 5.24. Ionic radii (A) for alkaline earth elements
(Molinari and Snodgrass, 1990).
Alkaline Earth
Element
Mg2+
Ca2+
Sr2+
Ba2+
Ra2+
Ionic Radii (A)
Crystal
0.65
0.99
1.13
1.35
1.52
Hydrated
4.28
4.12
4.12
4.04
3.98
The fate and mobility of radium in surface water, groundwater, and geologic environments have
been reviewed extensively. For example, see Benes (1990), Frissel and Koster (1990), Dickson
(1990), Onishi et al. (1981), Ames and Rai (1978), and others. Readers should also be aware of the
detailed reviews published in The Environmental Behavior of Radium (IAEA, 1990) and Radon, Radium
and Uranium in Drinking Water (Cothern and Rebers, 1990). Chapters pertinent to this Kd review, for
example, include the following reviews:
• Chemistry and radiochemistry of radium and the other elements of the uranium and thorium
natural decay series (Molinari and Snodgrass, 1990)
• Relationship of radium and radon in geologic formations (Michel, 1990)
• Removal of radium from drinking water (Hanslik and Mansfeld, 1990)
• Removal of radium from uranium mining effluents and leaching from sludges (Huck and
Anderson, 1990)
• Behavior of radium in soil (Frissel and Koster, 1990)
• Behavior of radium in continental surface water (Benes, 1990)
• Behavior of radium in groundwater (Dickson, 1990).
5.59
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5.7.3 Aqueous Speciation
The thermodynamic properties of radium aqueous species and solids are reviewed by Langmuir and
Riese (1985). Radium exists in nature only in the +2 oxidation state. In the pH range of 3 to 10, the
uncomplexed ion Ra2+ is expected to be the dominant aqueous species for dissolved radium. The
species Ra2+ is the only aqueous species for radium in the thermodynamic database for the
MINTEQA2 geochemical code (Allison eta!., 1991). Dissolved radium shows little tendency to
form aqueous complexes, although the aqueous complexes RaOH+, RaCl+, RaCOj (aq), and
RaSO4 (aq) are known. The thermodynamic constants for such aqueous complexes however are not
well established (Langmuir and Riese, 1985; Benes eta!., 1982).
Benes eta!. (1982) studied the speciation of radium using centrifugation and free-liquid
electrophoresis. The electrophoresis measurements conducted in 0.01 mol/1 chloride solutions in
the pH range 2 to 7 indicated that dissolved radium was present primarily as Ra2+. The results of
measurements conducted with solutions containing dissolved sulfate, carbonate, and bicarbonate
indicated that significant concentrations of aqueous radium complexes were present.
5.7.4 Dissolution/Precipitation/Coprecipitation
In moderate to high sulfate waters, precipitation and redissolution of calcium, strontium, and barium
sulfates, rather than adsorption/desorption, could control the concentrations of dissolved radium in
the soil environment. Although radium is moderately soluble in natural waters and radium salts are
less soluble than barium salts, the solubility of radium is unlikely controlled by pure RaSO4 in natural
waters. Precipitation of radium is possible as the solid-solution solids (Ba,Ra)SO41 and
(metal,Ra)CO3 in waters where the concentrations of dissolved sulfate and carbonate, respectively,
are sufficiently high.
The (Ba,Ra)SO4 coprecipitation process is well known as a preferred means for the removal of
dissolved 226Ra from effluents from uranium mining and milling operations (Huck and Anderson,
1990; Clifford, 1990). The addition of BaCl2 reacts with dissolved sulfate present in the effluent to
cause the almost instantaneous precipitation of (Ba,Ra)SO4 solid. The (Ba,Ra)SO4 coprecipitation
process is so efficient that microcrystals of BaSO4 have been incorporated into specific adsorbents
for removal of radium from drinking water and other radium-contaminated solutions (Clifford,
1990).
The (Ba,Ra)SO4 coprecipitation process has also been shown to be important process in controlling
the solubility of radium in natural waters (Baraniak et ai., 1999; Martin and Akber, 1999; Pardue and
Guo, 1998; Landa and Gray, 1995; Dickson, 1990; Benes and Strejc, 1986; Langmuir and Melchior,
1985; Benes et at., 1983; Church, 1979). Pardue and Guo (1998) studied the solubility of 226Ra in
contaminated sediments used integrated microcosm, geochemical modeling, and field-sampling
techniques. Their results indicated that the solubility of radium was controlled by the
coprecipitation of radium with barite. This conclusion was confirmed using selective extractions,
solution saturation measurements, theoretical solubility calculations, and x-ray diffraction. Landa
The barium end-member composition of this solid solution corresponds to the mineral known as barite (BaSO4).
5.60
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and Gray (1995) used selective extraction studies and leaching studies of components from uranium
mill tailings to show that alkaline earth sulfate was an important sorption phase for 226Ra.
Langmuir and Melchior (1985) investigated the geochemical controls for radium in some deep brines
in north Texas. Analyses of the geochemistry of calcium, strontium, barium, and radium sulfates in
these brines indicated that the concentrations of dissolved radium were likely controlled by
coprecipitation of trace radium in sulfate minerals such as celestite (SrSO4) and barite (BaSO4).
Langmuir and Melchior (1985) dismissed solid RaSO4 as the solubility control for dissolved radium,
because the brine compositions calculated to be 5 to 6 orders of magnitude undersaturated with
respect to RaSO4. Benes et at. (1983) investigated the particulate forms of radium and barium in a
system consisting of uranium mine waste waters purified by coprecipitation with barium sulfate and
of adjacent river waters. Results of selective dissolution analyses identified the following four
particulate forms of both elements: "loosely bound," "acid soluble," (Ba,Ra)SO4, and "in crystalline
detritus." Benes et al. (1983) determined that the main form of radium in the system was
(Ba,Ra)SO4. Radium was present mainly as "acid soluble" or "in crystalline detritus" in the river
water upstream of the sulfate-containing mine water discharge. In marine environments, radium is
also found to be efficiently scavenged and exchanged during the diagenetic1 formation of barite
(Church, 1979).
It should be noted that under certain reducing conditions, the (Ba,Ra)SO4 coprecipitate is not
thermodynamically stable. Sulfate-reducing bacteria can produce rapid dissolution of (Ba,Ra)SO4
sludge under suitable reducing conditions and appropriate carbon sources (Huck and Anderson,
1990). For example, Pardue and Guo (1998) observed that 226Ra in contaminated sediment was
remobilized under anaerobic, sulfate-reducing conditions.
5.7.5 Adsorption/Desorption
5.7.5.1 Guidance for Screening Calculations of Adsorption
Compared to most other contaminants, very limited sorption data, especially Kd values, exist for
radium on soils and sediments. Moreover, the reader is cautioned that any data that indicates very
high adsorption of radium on geological materials should be suspect due to the possibility that
(Ba,Ra)SO4 coprecipitation may have occurred during the measurements. Development of K^
look-up tables for radium is not possible given the minimal number of adsorption studies.
However, as an alkaline earth element, the adsorption behavior of radium will be similar and
somewhat greater to that of strontium for which extensive studies and data exist [see Section 5.8 in
Volume II (EPA 1999c)]. For screening calculations of radium migration in soils, the Kd look-up
table listed for strontium in Volume II can be used as general guidance for radium. Given the
absence of definitive maximum and minimum Kd values for radium as a function of the key
geochemical parameters, such as pH, EPA suggests that Kd values measured for radium at site-
specific conditions are thus essential for site-specific contaminant transport calculations and
conceptual models.
Diagenetic - "caused by the chemical, physical, and biological changes undergone by a sediment after its initial
deposition, and during and after its consolidation into a coherent solid rock, exclusive of surficial weathering and
metamorphism" (Bates and Jackson, 1980).
5.61
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Radium migrates as a cation competing with other alkaline earth cations for sorption sites in soils
systems. Relative to other alkaline earth elements, the relative affinity this group of elements for ion
exchange on clay minerals has been described as follows (Sposito, 1989):
Ra2+ > Ba2+ > Sr2+ > Ca2+ > Mg2+
The adsorption of radium has been shown to be strongly dependent on ionic strength and
concentrations of other competing ions in that adsorption of radium decreases with increasing ionic
strength.
The studies reviewed below indicate that radium is readily adsorbed to clays and mineral oxides
present in soils, especially at near neutral and alkaline pH conditions. For the pH conditions of
most natural waters, dissolved radium will be present primarily as the uncomplexed Ra2+ cation.
Sorption studies discussed below confirm the adsorption behavior expected for Ra2+ as a function of
pH. Radium adsorption on mineral phases is negligible at acidic pH values, and increases with
increasing pH. Because adsorption of cations is coupled with a release of H+ ions, cation adsorption
is greatest at high pH and decreases with decreasing pH. For iron oxides, the increase in adsorption
starts typically at pH values from 6 to 8 and is at a maximum by pH ~10 or less. As discussed in
Volume I (EPA, 1999b), the pH range at which adsorption of cations begins to increase on mineral
surfaces depends on the values of the zero point charge (PZC) (often used as a reference for how
the surface charge of solids varies with pH) for each type of mineral. In general, at pH values less
than PZC, the mineral surface serves as a strong adsorbent for anions. At pH values greater than
the PZC, the surface strongly adsorbs dissolved cationic constituents.
Only a few studies have been conducted with respect to the adsorption of radium on organic matter.
The results of these studies also suggest that radium may be strongly adsorbed by organic material in
soils (Greeman eta!., 1999; Nathwani and Phillips, 1979a; 1979b). As noted previously, one of the
major recommendations of this report is that for site-specific calculations, partition coefficient
values measured at site-specific conditions are absolutely essential.
5.7.5.2 General .Adsorption Studies
Sturchio et al. (2001) studied the processes that affect the mobility of radium isotopes in continuous
Paleozoic carbonate aquifers in Missouri, Kansas, and Oklahoma. The concentrations of dissolved
radium were correlated to the salinity and concentrations of other alkaline earth elements. The
behavior of radium in the aquifers was explained by salinity-dependent sorption/desorption
processes on surface mineral coatings.
Greeman et al. (1999) measured the abundance, chemical, and mineralogical form of 226Ra, 238U, and
232Th in soil samples from 12 sites in the eastern United States. Selective chemical extraction and
size fraction techniques were used to determine the abundance and radiometric equilibrium
condition for 226Ra, 238U, and 232Th in the following soil fractions: exchangeable cation, organic
matter, "free iron oxide," sand, silt, and clay size. Their results indicated that radium was enriched in
the exchangeable cation and organic (humic) matter fractions in these soils.
Landa and Gray (1995) conducted selective extraction studies and studies of radionuclide sorption
by and leaching from components of uranium mill tailings. Their results indicated that alkaline earth
5.62
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sulfate and hydrous ferric oxide solids are important sorption phases for 226Ra.
Willet and Bond (1995) used batch equilibration measurements to investigate the sorption of 226Ra
on highly weathered and sandy soils from the area near the Ranger Uranium Mine in Australia. The
measurements were conducted using a background electrolyte solution of 0.0025 M MgSO4, and
initial concentrations of dissolved 226Ra of 100 Bq/1 (1.2xl041 mol/1), 150 Bq/1 (l.SxlO41 mol/1),
and 200 Bq/1 (2.4x10"11 mol/1). The pH values of the soil/electrolyte suspensions ranged between
4.6 and 7. The sorption data for 226Ra were fitted to the Freundlich isotherm. Willet and Bond
(1995) found no clear relationship between 226Ra sorption and the concentrations of organic matter
and clay in the soils. Sorption of 226Ra on all soils increased with increasing soil pH. For some soils,
the sorption of 226Ra was nearly complete. Because trace concentrations of dissolved 226Ra were
used for their measurements, Willet and Bond (1995) concluded that precipitation of radium salts
was not likely in their experiments. The results also indicated that the 226Ra was strongly sorbed to
these soils. Less than 1 percent of the sorbed 226Ra was remobilized by resuspension of the soils in
the background electrolyte solution.
Berry et al. (1994) conducted diffusion experiments of the mobility of radium through sandstone.
Their results indicated that the sorption of radium was affected by competitive ion effects. Radium
sorption in the high ionic strength groundwater experiment was less than 50 percent of the sorption
measured in the lower ionic strength groundwater solution.
Nathwani and Phillips (1979a, 1979b) used batch equilibration experiments to study 226Ra adsorption
by soil with different physical-chemical characteristics. The measured 226Ra adsorption followed
Freundlich and Langmuir adsorption isotherms over a large range of 226Ra concentrations. Organic
matter and clay were determined to be the dominant phases contributing to the adsorption of 226Ra
on these soils. Nathwani and Phillips (1979a) suggested that adsorption affinity of the organic
matter and clays was primarily due to their cation exchange capacity (CEC). Their results also
indicated that organic matter adsorbed approximately 10 times more 226Ra than did the clays. The
results of Nathwani and Phillips (1979b) show that the addition of competing alkaline earth cations
to the system can greatly affect radium sorption on the clay minerals.
5.7.5.3 Kj Studies for Radium on Soil Materials
Compared to other contaminants and in particular to 90Sr, which is another alkaline earth isotope of
environmental concern, a very limited number of published Kd studies was identified for the
adsorption of radium on soils. Meier et al. (1994) completed parametric studies with site-specific
waters and crushed sedimentary rocks from strata that overly the Gorleben salt dome in Germany.
In the pH range from approximately 4 to 9, the adsorption and desorption of radium increased with
increasing pH. The Kd values measured for the adsorption of radium on a sandy sediment in
groundwater were 6.7, 12,6, 26.3, and 26.3 ml/g at pH values of 6, 7, 8, and 9, respectively. For the
same system, the desorption Kj values were 10.9, 31, 38, and 29 ml/g at pH values of 6, 7, 8, and 9,
respectively. As commonly reported for the sorption of strontium (EPA, 1999c), these adsorption
and desorption Kd values indicate that radium is essentially completely reversibly adsorbed.
Nathwani and Phillips (1979b) measured the adsorption of 226Ra on soils as a function of calcium
concentrations. All experiments were conducted with a constant total concentration of 226Ra of
10 pCi/ml (4.5xlO~8 mol/1). The dissolved concentrations of calcium varied from 0 to 0.5 mol/1.
Values for Kd for 226Ra estimated from those plotted in Nathwani and Phillips (1979b) are listed in
Table 5.25. The characteristics of the soil samples for which Kd values are reported in Nathwani and
5.63
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Phillips (1979b) are given in Table 5.26, as reported in the companion paper by Nathwani and
Phillips (1979a). Nathwani and Phillips (1979b) also plot Kd values for each soil for calcium
concentrations of 0.02 and 0.04 mol/1. These values, however, are not significantly different from
the Kd values determined for the respective soils at calcium concentrations of 0.01 and 0.05 mol/1,
and show the same trend as a function of calcium concentrations. Sorption of 226Ra by all soils
decreased with increasing concentrations of dissolved calcium. The Kd values reported by Nathwani
and Phillips (1979b) are unusually large, and orders of magnitude greater than those reported by
most researchers. This suggests that precipitation of radium may have occurred during the course of
these measurements.
Serne (1974) measured Kd values for radium on four sandy, arid soil samples from Utah using a
simulated river water solution. The final pH values of the soil/river water suspensions ranged from
7.6 to 8.0. The soil consisted primarily of quartz and feldspar with 2-5 percent calcite and minor
amounts of muscovite and smectite. The Kd values ranged from 214 to 467 ml/g for the four soil
samples (Table 5.27). Serne (1974) was able to correlate the Kd values to the cation exchange
capacity (CEC) values for these soils.
Table 5.25. Radium Kd values (ml/g) as function of calcium
concentration [Nathwani and Phillips, 1979b)]. [^ values
estimated from Figure 3 in Nathwani and Phillips (1979b).]
Soil Series
Wendover
Grimsby
St. Thomas
Kd (ml/g) as Function of
Concentration of Calcium (mol/1)
0
9.5xl05
1.2xl05
3.8xl04
0.005
3.1xl05
4.0xl04
l.lxlO4
0.01
2.1xl05
3.2x1 04
7.1xl03
0.05
l.lxlO5
1.9xl04
4.3xl03
5.7.5.4- Published Compilations Containing Kj Values for Radium
Because the references in this section are often cited or used for comparison in other publications,
the following summaries are provided for completeness. As noted previously, the reader is
cautioned that any very high adsorption results, including Kd values, reported for radium should be
suspect due to the possibility that (Ba,Ra)SO4 coprecipitation occurred during the measurements. It
is recommended that the reader review the original reference and the references cited therein to
understand the procedures and sources of the Kd values used for each compilation. The
compilations do not typically consider important factors that contribute to variability in sorption,
such as pH. Moreover, in cases where very large Kd values are listed, there is a risk that the original
Kd measurement may have included precipitated components.
5.64
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Table 5.26. Properties of soil samples for which Kd values are given
in Nathwani and Phillips (1979b). [The soil characteristics were taken
from Nathwani and Phillips (1979a).]
Soil
Series
Wendover
Grimsby
St.
Thomas
Texture
Silty Clay
Silt Loam
Sand
pH
5.4
4.3
5.2
Organic
Matter
%
16.2
1.0
3.1
Sand
%
6.7
43.7
91.1
Silt
%
47.9
48.9
6.8
Clay
%
45.4
7.4
1.3
CEC
(meq/100 g)
34.7
10.4
10.9
Table 5.27. Radium Kd values (ml/g) measured by Serne
(1974) for sandy, arid soil samples from Utah.
Soil
Sample
Soil I
Soil II
Soil III
Soil IV
Final pH
7.9-8.0
7.6-7.7
7.8-7.9
7.6-7.8
Kd
(ml/g)
354 ± 15
289 ± 7
467 ± 15
214 ± 15
Looney et al. (1987) tabulated estimates for geochemical parameters needed for environmental
assessments of waste sites at DOE's Savannah River Plant in South Carolina. Looney et al. list Kd
values for several metal and radionuclide contaminants based on values that they found in 1-5
published sources. For radium, Looney et al. list a "recommended" Kd of 100 ml/g, and a range
from 10 to 1,000,000 ml/g. Looney et al. note that their recommended values are specific to the
Savannah River Plant site, and they must be carefully reviewed and evaluated prior to use in
assessments at other sites.
Thibault et al. (1990) (also see Sheppard and Thibault, 1990) present a compilation of soil Kd values
prepared to support radionuclide migration assessments for a Canadian geologic repository for spent
nuclear fuel in Precambrian Shield plutonic rock. Thibault et al. collected Kd values from other
compilations, journal articles, and government laboratory reports for important elements, such as
radium, that would be present in the nuclear fuel waste inventory. The radium Kd values listed in
Thibault etal. (1990) are included in Table 5.28. Thibault et al. (1990) describe the statistical methods
used for analysis of the compiled Kd values. The range for the Kd values used to calculate the
"geometric mean" cover several orders of magnitude. Readers are cautioned against using
"geometric mean" values or any other form of averaged Kd values as "default" Kd values, as such
values are usually calculated from data compiled from different investigators for different soil
5.65
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systems. These mean or average values do not represent any particular environmental system and
geochemical conditions. As discussed in Volume I (EPA, 1999b), the variation observed in the
literature for Kd values for a contaminant is due to differences in sorption mechanisms, geochemical
conditions, soil materials, and methods used for the measurements.
5.7.5.5 Kj Studies of Radium on Pure Mineral, Oxide, and Crushed Rock Materials
Numerous adsorption studies have been conducted of radium on pure minerals, oxide phases, and
other geologic-related materials. The Kd values listed in these studies are not necessarily relevant to
the mobility and sorption of radium in soils. However, they are listed in Appendix H for
completeness. The references cited in Appendix H are listed in the main reference list in Chapter 6.
The potential value of the references that they cite and the sorption processes that they discuss is left
to the reader. The results of the studies conducted with minerals and oxide phases demonstrate that
the adsorption of radium on is dependent on pH, as expected for cations, and decreases with
increasing concentrations of competing ions. The studies of radium sorption on crushed rock were
conducted typically as part of national research programs to investigate the feasibility of geological
disposal of high-level radioactive waste (HLW).
Table 5.28. Radium ^ values (ml/g) listed in
Thibault et a/. (1990, Tables 4 to 8).
Soil Type
Sand
Silt
Clay
Organic
Kd Values (ml/g)
Geometric
Mean
500
36,000
9,100
2,400
Number of
Observations
3
3
8
1
Range
57 - 21,000
1,262 - 530,000
696 - 56,000
None Listed
5.8 Technetium Geochemistry and Kd Values
5.8.1 Overview: Important Aqueous- and Solid-Phase Parameters Controlling
Retardation
In natural environments, the most stable valence states of technetium (Tc) are +7 and +4 under
oxidizing and reducing conditions, respectively. Technetium(VII) in oxic environmental systems is
highly mobile (i.e., Kd values are ~0 ml/g). The dominant aqueous Tc(VII) species in oxic waters is
the oxyanion TcO4~, which is highly soluble and essentially nonadsorptive. However, under reducing
conditions in soil and geologic systems, Tc(IV) is expected to dominate because of biotic and abiotic
reactive processes, such as surface-mediated reduction of Tc(VII) by iron (Fe)(II). In the absence of
aqueous complexing agents other than OH", Tc(IV) is considered to be essentially immobile,
because it readily precipitates as sparingly soluble hydrous oxide and forms strong surface complexes
with surface sites on iron and aluminum oxides and clays.
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5.8.2 General Geochemistry
Technetium [Tc, atomic number (Z) = 43] exists in valence states from +7 to -1. In natural
environments, the most stable valence states of technetium are +7 and +4 under oxidizing and
reducing conditions, respectively. Other valence states are encountered chiefly in complex
compounds. The chemical behavior of technetium in these two oxidation states differs drastically.
In the +7 valence state, dissolved technetium exists as pertechnetate anion, TcO4, over the complete
pH range of natural waters. Because the pertechnetate anion is highly soluble and is not strongly
sorbed, it is highly mobile in most oxidizing systems. In the +4 valence state, technetium exists as
the tetravalent cation and is relatively immobile in the absence of strongly complexing ligands.
Tc(IV) is highly sorbed, and forms the sparingly soluble TcO2-nH2O solid. There are 34 reported
isotopes of technetium (Tuli, 2000). Of the few technetium isotopes having long half lives, "Tc (t1/2
= 2.11 x 105 yr) is a long lived fission product generated during the irradiation of uranium-
containing nuclear fuels and is the primary isotope of environmental interest. Most of the other
technetium isotopes have half lives of hours or less.
The behavior of technetium in environmental systems has been reviewed extensively by others.
Reviews include Lieser (1993), Gu and Schulz (1991), Sparks and Long (1987), Meyer et oL (1985a),
Beasley and Lorz (1984), Coughtrey et al. (1983), Omshi et aL (1981), Wildung et aL (1979), Ames and
Rai (1978), and others. Huges and Rossotti (1987) review in detail the solution chemistry of
technetium.
5.8.3 Aqueous Speciation
Rard eta!. (1999) have published an extensive, detailed review of the chemical thermodynamics of
technetium. Figure 5.9 is an Eh-pH diagram that shows the dominant aqueous species of
technetium as a function of pH and Eh (redox potential) at 25°C with respect to the thermodynamic
stability of water. Technetium is present in +7 and +4 valence states under oxidizing and reducing
conditions, respectively.. The predominate Tc(VII) aqueous species is the pertechnetate oxyanion
TcO4. The TcO4 ion is stable over the complete pH range of natural waters, and its not known to
form any strong aqueous complexes. Although the thermodynamic stability of TcO4 is well
established, thermodynamic data for technetium aqueous species and solids in its various valence
states is extremely limited.
5.8.4 Dissolution/Precipitation/Coprecipitation
Technetium(VII), TcO4, is highly soluble, and does not form solubility-controlling phases in soil
systems. Technetium(VII) can be reduced to Tc(IV) by biotic and abiotic processes. This reduction
usually results in the immobilization of technetium under reducing conditions via the formation of
the sparingly soluble solid TcO2-nH2O. The solubility of Tc(IV) has been studied experimentally by
Eriksen eta!. (1992), Meyer eta!. (1991b), Guppy eta!. (1989), and others.
The reduction of Tc(VII) to Tc(IV) by surface-mediated processes has been the subject of extensive
studies [e.g., Wharton et aL (2000); Byegard et al. (1992); Eriksen and Cm (1991); Hames et al. (1987);
Bondietti and Francis (1979]. Such reactions have recently been the subject of considerable interest
because they could have a significant affect on the mobility of technetium in waste streams, vadose
5.67
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zones, aquatic sediments, and groundwater. They are the basis for certain remediation technologies,
such as permeable barriers composed of zero-valent iron particles (i.e., as metallic iron) or sodium-
dithionite reduced soils, which are currently being tested for immobilization of groundwater
contaminants.
Figure 5.9. Eh-pH stability diagram for the dominant technetium
aqueous species at 25°C. [Diagram based on a total
concentration of 10"8 mol/l dissolved technetium.]
1
^ .5
-i— •
0
^
t~
LU
0
-.5
(
x /c02+ ~"^~^
-S£f^ TcOOhT ^--^ -
^^^ Tc04
^x^.
\ xxx^
v "***x
^v^ ^^^
Tc3+^^^_ TcO(OH)2° (aqr\
"^--.
"^-~,
) 2 4 6 8 10 12 1
PH
4
Cui and Eriksen (1996b) studied the surface-mediated reduction of TcO4 by Fe(II)-containing
mineral material collected from granite fractures and by hornblende and magnetite at neutral to
alkaline pH conditions. Batch equilibration experiments were conducted using anoxic synthetic
groundwater and perchlorate solutions. The Eh values for the various mineral/water systems
studied ranged from +60 to -150 mV. Cui and Eriksen (1996b) determined that magnetite was at
least 1 order of magnitude more effective than hornblende in the reduction of Tc(VII).
Hornblende, on the other hand, was slightly less effective than crushed granite and the fracture
filling materials with respect to Tc(VII) reduction. Desorption experiments conducted by Cui and
Eriksen (1996b) with anoxic and oxic groundwater solutions over 6 and 21 days, respectively,
indicated that remobilization of the reduced/sorbed technetium was a slow process. A desorption
experiment conducted with the addition of a high concentration of H2O2 resulted in a sudden
5.68
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increase of remobilized dissolved technetium. Cui and Eriksen suggested that the addition of H2O2
consumed the reductive capacity of the solids which permitted the reoxidation and dissolution of the
sorbed hydrous Tc(IV) oxide precipitate.
Byegard et al. (1992) studied the sorption of technetium on magnetite in groundwater solutions
under oxic and anoxic conditions. The solutions used in the laboratory experiments included a
synthetic groundwater (pH 8.2) and a natural groundwater solutions (pH 8.8). For some
experiments, 6 ppm dissolved Fe(II) was added to the synthetic groundwater solution. Under
anoxic conditions, 95 percent of the initial concentration of dissolved technetium was sorbed within
48 hours on the magnetite reacted with the Fe (II)-augmented synthetic groundwater. However, only
16 and 15 percent of the technetium was sorbed in the magnetite systems reacted with Fe(II)-free
synthetic and natural groundwater, respectively. Under oxic conditions at 48 hours, the amounts of
sorbed technetium were 9, 4, and 4 percent, respectively, for magnetite reacted with
Fe(II)-augmented synthetic, Fe(II)-free synthetic, and natural groundwater. Byegard et al. (1992) also
observed a significant decrease in the concentration of dissolved technetium in experiments
conducted with Fe(II)-augmented synthetic groundwater under anoxic conditions in the absence of
magnetite.
Haines et al. (1987) used Fourier transform infrared (FTIR) spectroscopy to study the sorption of
Tc(VII) at room temperature on synthetic magnetite and hematite particles that were contacted with
dissolved TcO4 and nitrogen-purged pH 6 water. Their experiments indicated that, under these
geochemical conditions, sorption of Tc(VII) on magnetite occurred by a surface-mediated
reduction/precipitation mechanism. Based on pHZPC values reported by others for hematite and
magnetite of 8.5 + 0.5 and 6.5 + 0.2, respectively, hematite and, to a lesser degree, magnetite, were
expected to have a positive surface charge and thus capable of anion adsorption in these pH 6
solutions. Haines et al. (1987) proposed that the sorption of Tc(VII) at pH values less than pHzpc is
initiated by electrostatic attraction of the TcO4 anion by the magnetite surface. The sorbed Tc(VII)
is then reduced to Tc(IV) by Fe(II) centers on the magnetite surface to simultaneously both
precipitate sparingly soluble Tc(IV) oxide and oxidize Fe(II) to Fe(III) to form an Fe(III) oxide on
the magnetite surface.
Microbial reduction of Tc(VII) has been suggested as a potential mechanism for removal of
technetium from contaminated groundwater and waste streams [e.g., Lovley (1993, 1995); Lloyd et al.
(1997, 2000b)]. Dissimilatory metal reducing bacteria (Lloyd and Macaskie, 1996; Wildung et al.,
2000) and the sulfate reducing bacterium Desulfovibrio desulfuricans [Lloyd et al., 1998, 1999] are
capable of coupling the oxidation of organic carbon or hydrogen to the reduction of Tc(VII).
Wildung et al. (2000) examined the influence of electron donor and the presence of dissolved
bicarbonate on the rate and extent of enzymatic reduction of TcO4~ by the subsurface metal-
reducing bacterium Shewanellaputrefadens CN32 under anoxic conditions. Using a variety of analytical
methods and geochemical modeling techniques, these studies indicated that the dominant solid-
phase reduction product in both saline (0.85 percent NaCl) and bicarbonate (30mM NaHCO3)
systems was amorphous Tc(IV) hydrous oxide. However, the reduced technetium in saline systems
was associated principally with the cell surface, whereas Tc(IV) in bicarbonate systems was present
primarily in extracellular particulates.
5.69
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5.8.5 Adsorption/Desorption
5.8.5.1 Guidance for Screening Calculations ofA-dsorbtion
The dominant aqueous Tc(VII) species under oxidizing conditions is the oxyanion TcO4~, which is
highly soluble and essentially nonadsorptive. For soils with low contents of organic material, the
reported Kd values range from 0 to approximately 0.5 ml/g, although most values are less than
0.1 ml/g (see studies reviewed below). The very low Kd values measured for Tc(VII) and the limited
availability of Kd values as a function of key geochemical parameters, such as pH, preclude
derivation of meaningful look-up tables for Tc(VII) Kd values. For screening calculations of off-site
migration of Tc(VII), a Kd of 0 ml/g si suggested as a conservative minimum value for low organic
soils under oxidizing conditions at pH values greater than 5. As an anion, the adsorption of TcO4~ is
expected to increase with decreasing pH at pH values less than 5. Values of Kd measured by Kaplan
et al. (2000a) for a wetland sediment ranged from approximately 0 ml/g at pH 4.6 to 0.29 at pH 3.2.
The maximum Kd value that they determined for an upland sediment was 0.11 ml/g at pH 3. As
noted previously, however, one of the major recommendations of this report is that for site-specific
calculations, partition coefficient values measured at site-specific conditions are absolutely essential.
The sorption of TcO4~ has been found to be positively correlated to the organic carbon content of
soils (Wildung et al., 1974; 1984). However, studies of the effect that organic material has on the
sorption of Tc(VII) in soils are limited. As an extreme example, Sheppard and Thibault (1988)
reported Kd values of greater than 2 ml/g based in situ Kd values derived from analyses of the dried
peat and pore water from the Precambrian Shield in Canada. Measurable adsorption of Tc(VII)
observed in experiments conducted with organic material as well as with crushed rock and Fe(II)-
containing minerals has been attributed to the reduction of Tc(VII) to Tc(IV). Technetium(IV) is
essentially immobile, because it readily precipitates as sparingly soluble hydrous oxides and forms
strong complexes with surface sites on iron and aluminum oxides and clays. Studies by Landa et al.
(1977) however suggest that anaerobic conditions may not be a prerequisite to technetium sorption
by soils, and that the living and nonliving organic fraction of soil may have a role in technetium
sorption.
5.8.5.2 General Adsorption Studies
Winkler et al. (1988) studied the sorption of technetium on several sands and single mineral phases
using batch-equilibration, recirculation column and flow-through column experiments. The studies
were conducted under anoxic conditions using simulated calcium-bicarbonate groundwater
containing 13.6 meq/1 salt. No significant sorption of technetium was observed in the experiments
conducted with the sands (Kd values less than 0.2 ml/g). Little or no sorption was measured in
experiments conducted with pure quartz or smectite. Significant sorption, however, was observed in
the pyrite experiments, and appeared to be a function of time even at low concentrations (i.e., 1
wt. percent) of pyrite. In a flow-through column experiment containing a 1:99 mixture (by
wt. percent) of pyrite to quartz, rapid and complete sorption of technetium was observed. The
sorbed technetium was localized in the first 50 mm of the 200-mm long columns. The pH and Eh
values for this flow-through experiment were, respectively, 7 and +400 mV.
Lieser and Bauscher (1987) conducted batch equilibration experiments to study the sorption and
desorption of technetium under oxic and anoxic conditions on five sediments and
5.70
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five corresponding groundwater samples from Gorleben, Federal Republic of Germany. The
sediments consisted mainly of quartz and minor concentrations of calcite, dolomite, anhydrite,
kaolinite, illite, and montmorillonite. Low sorption of technetium (i.e., Kd values of 0.1-0.3 ml/g)
was observed in those experiments conducted under oxic conditions. Technetium sorption
decreased with increasing salinity, and was found to be reversible. Under anoxic conditions, high
sorption values were measured for technetium on these water/sediment systems. Technetium
sorption did not depend on salinity and was not reversible under these anoxic conditions. However,
at higher salinity values, steady-state sorption of technetium was attained more quickly. When the
Eh was increased, Lieser and Bauscher (1987) determined that the sorption behavior changed
abruptly from high to low adsorption at Eh of 170+60 mV and pH of 7.0+0.5. In all experiments,
analysis of the dissolved technetium indicated the presence of only Tc(VII) and no colloidal size
technetium was detected by ultrafiltration.
5.8.5.3 Kj Studies for Technetium on Soil Materials
Kaplan and coworkers have conducted a series of laboratory studies on the adsorption of TcO4 on
soils and mineral phases. Kaplan et al. (2000a) measured Kd values for TcO4 as a function of pH on
two sediments from the Savannah River Site at Aiken, South Carolina. The Kd values ranged
from 0.29 to -0.13 ml/g. Their measurements also indicate that the Kd values increased with
decreasing pH as would be expected for dissolved anionic contaminants. Kaplan et al. (1998b)
suggested that the negative Kd values measured for technetium were due to anion exclusion effects.
Kaplan et al. (1998a) used the batch equilibration technique to measure the Kd values for technetium
under oxic conditions on 20 sediment samples from the Hanford formation at the Hanford Site in
southeastern Washington. The sediment samples were taken from 1 borehole, and included material
from 3 layers of the Hanford formation. Each sediment sample was equilibrated with
uncontaminated groundwater from the Hanford Site, which was spiked with "Tc prior to the Kd
measurements. The groundwater solution had a low ionic strength and a pH of 8.4 [see Table 1 in
Kaplan et al. (1998a)]. The mean Kd values of 3 replicates measured for technetium on each
Hanford sediment sample are listed in Table 5.29, and ranged from -0.04 to 0.01 ml/g. The
measurements indicated essentially no sorption of technetium on Hanford sediments under these
conditions.
Kaplan et al. (1998b) studied the effects of ionic strength and high pH on the sorption of technetium
under oxic conditions on sediments from the Hanford Site in Washington. Batch sorption
experiments were completed using the <2-mm size fraction of sediment collected from
Trench AE-3 in the 200 Area of the Hanford Site. The sediment was characterized as a silty loam
with cation exchange capacity (CEC) of 6.4 meq/100 g. The carbonate content of the sediment is
primarily calcite. Based on analyses reported by Kaplan et al. (1996) for a sediment from a nearby
location, the Trench AE-3 sample was expected to contain approximately 0.2-0.5 wt. percent
amorphous Fe2O3 (Kaplan et al. (1998b). The <2-mm size fraction contained 41 percent sand,
50 percent silt, and 9 percent clay. The clay fraction of the sediment contained primarily smectite
(57 percent), illite (19 percent), and vermiculite (14 percent). The groundwater selected for the
sorption studies was an uncontaminated groundwater from the Hanford Site with a low ionic
strength and pH of 8.4 [see Table 2 in Kaplan et al. (1998b)]. The effect of ionic strength on the
sorption of technetium was investigated by adding 0.05, 0.10, 0.50, and 1.00 M NaClO4 solutions to
the groundwater. To study the effect of high pH on the sorption of technetium, NaOH was added
5.71
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to separate samples of groundwater to adjust the pH to 8.1, 9.9, 10.2, 11.0, and 11.9. The sediment
was equilibrated with these amended groundwater solutions. Based on the experimental conditions
and geochemical modeling calculations, the dissolved technetium was assumed to be TcO4. The
mean Kd values of three replicates measured for technetium in the ionic strength and pH
experiments are listed in Table 5.29. Kaplan et al. (1998a) noted that negative Kd values had usually
been attributed to experiment error associated with measuring the concentration of a nonadsorbing
solute.
However, based on the results of Kaplan and Serne (1998), Kaplan et al. (1998b) suggested that the
negative Kd values measured for technetium were due to anion exclusion effects. The reason for the
high Kd value (3.94 ml/g) measured for technetium in the 1.00 M NaClO4 solution was not known.
Kaplan et al. (1998b) speculated that the higher ionic strength allowed greater interaction of the
technetium with the mineral surfaces by decreasing the double layer around the sediment particles.
The Kd values determined for technetium (i.e., 1.04 to 1.07 ml/g) at pH values greater than 8.8 in the
NaOH-amended solutions were greater than expected. The reason for these greater-than-expected
values was not known.
Kaplan and Serne (1998) [also see Kaplan et al. (1996)] used the batch technique to measure the Kd
values for technetium, as TcO4, on three sediment samples from the Hanford Site in southeastern
Washington. The measurements were conducted with a groundwater sample (pH 8.3) taken from a
well located in an uncontaminated area of the Hanford Site. The sediment samples included a loamy
sand (sample TSB-1), a silty loam (sample Trench AE-3), and a very coarse sand (sample
Trench-94). The Kd values at 266 days of contact time were very low for all sediments, and ranged
from 0.11 to-0.18mg/l.
Kaplan et al. (1996) measured the Kd values for TcO4 on sediment samples from the 200 Areas at the
Hanford Site. The sorption measurements were conducted under oxidizing conditions using
uncontaminated groundwater (pH=8.46) from Hanford Site Well S3-25. The sorption values for
TcO4 were determined for three sediment samples, Trench 94, Trench AE-3, and TSB-1, using the
batch equilibration method. As with previous sorption studies of Hanford Site sediments at
oxidizing conditions, no significant sorption of TcO4 was observed. The average Kd values
measured after 30 days of contact for Trench 94 and Trench AE-3 sediments were -0.02 ml/g and
-0.05 ml/g, respectively. The effect of contact time (7-398 days) on the sorption of TcO4 to these
sediments was also studied by Kaplan et al. (1996). The Kd values for all three sediments varied from
slightly negative at short contact times to slightly positive at contact times greater than 300 days.
The values ranged from -0.18 to 0.11 ml/g.
5.72
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Table 5.29. Technetium Kd values (ml/g) measured for
Hanford sediments under oxidizing conditions (Kaplan
etal. (1998a, 1998b).
pH
8.54
8.80
8.77
8.73
8.75
8.77
8.52
8.50
8.52
8.56
8.94
8.82
8.81
8.89
8.88
Kd
(ml/g)
-0.01
±0.02
-0.02
±0.03
0.01
±0.00
-0.01
±0.03
0.00
±0.02
-0.01
±0.02
-0.04
±0.01
-0.02
±0.02
0.00
±0.01
0.00
±0.02
0.00
±0.02
-0.01
±0.03
0.00
±0.02
0.01
±0.01
0.00
±0.01
CEC
(meg/100 g)
5.07
4.73
4.60
4.62
4.11
2.32
4.98
4.72
4.67
4.56
7.33
8.41
9.03
6.63
8.36
Solution
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Soil
Identification
Hanford Sediment
B8500-07A
Hanford Sediment
B8500-10A
Hanford Sediment
B8500-12A
Hanford Sediment
B8500-14A
Hanford Sediment
B8500-15A
Hanford Sediment
B8500-16A
Hanford Sediment
B8500-17A
Hanford Sediment
B8500-19A
Hanford Sediment
B8500-20A
Hanford Sediment
B8500-21A
Hanford Sediment
B8500-22A
Hanford Sediment
B8500-23A
Hanford Sediment
B8500-24A
Hanford Sediment
B8500-25A
Hanford Sediment
B8500-27A
Reference and
Comments
Kaplan etal. (1998a)
Kaplan etal. (1998a)
Kaplan etal. (1998a)
Kaplan etal. (1998a)
Kaplan etal. (1998a)
Kaplan etal. (1998a)
Kaplan etal. (1998a)
Kaplan etal. (1998a)
Kaplan etal. (1998a)
Kaplan etal. (1998a)
Kaplan etal. (1998a)
Kaplan etal. (1998a)
Kaplan etal. (1998a)
Kaplan etal. (1998a)
Kaplan etal. (1998a)
5.73
-------
Continuation of Table 5.29
pH
8.84
8.56
8.93
8.92
8.89
7.74
7.76
7.73
7.70
8.1
9.9
10.2
11.0
11.9
Kd
(ml/g)
-0.01
±0.03
-0.03
±0.04
-0.01
±0.01
-0.02
±0.01
-0.02
±0.01
-0.16
±0.04
-0.13
±0.00
-0.28
±0.01
3.95
±0.99
-0.02
±0.01
1.04
±0.06
1.05
±0.02
1.07
±0.05
1.07
±0.03
CEC
(meg/100 g)
7.77
10.98
8.39
6.21
6.65
6.4
6.4
6.4
6.4
6.4
6.4
6.4
6.4
6.4
Solution
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
0.05 M NaClO4
+ Groundwater
0.1 M NaClO4 +
Groundwater
0.5 M NaClO4 +
Groundwater
1.0 M NaClO4 +
Groundwater
Groundwater +
NaOH
Groundwater +
NaOH
Groundwater +
NaOH
Groundwater +
NaOH
Groundwater +
NaOH
Soil
Identification
Hanford Sediment
B8500-29A
Hanford Sediment
B8500-31A
Hanford Sediment
B8500-32A
Hanford Sediment
B8500-34A
Hanford Sediment
B8500-35A
Hanford Trench
AE-3 Sediment
Hanford Trench
AE-3 Sediment
Hanford Trench
AE-3 Sediment
Hanford Trench
AE-3 Sediment
Hanford Trench
AE-3 Sediment
Hanford Trench
AE-3 Sediment
Hanford Trench
AE-3 Sediment
Hanford Trench
AE-3 Sediment
Hanford Trench
AE-3 Sediment
Reference and
Comments
Kaplan et aL (1998a)
Kaplan et al. (1998a)
Kaplan et aL (1998a)
Kaplan et aL (1998a)
Kaplan et aL (1998a)
Kaplan et aL (1998b)
Kaplan et aL (1998b)
Kaplan et aL (1998b)
Kaplan et aL (1998b)
Kaplan et aL (1998b)
Kaplan et aL (1998b)
Kaplan et aL (1998b)
Kaplan et aL (1998b)
Kaplan et aL (1998b)
Serne et al. (1993) measured the Kd values for TcO4 on three samples of sediment from the Hanford
Formation at the Hanford Site. The sediment samples included two loamy sands (samples TBS-1
and Trench-8) and one sand (sample CGS-1). The measurements were conducted using an
uncontaminated groundwater (pH 8.14) sample from the Hanford site. Characterization of the
sediment samples indicated that they contained very little amorphous or poorly crystalline hydrous
Al, Mn, and Fe oxides (determined by hydroxylamine-hydrochloride extraction) and very low
organic carbon contents. The sorption measurements were conducted using the batch-equilibration
5.74
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method under oxidizing conditions. After 35 days of contact, the pH values for the CGS-1, TBS-1,
and Trench-8 sediment suspensions were 7.9-8.4., 8.0-8.4, and 8.22, respectively. Under the
conditions studied, no significant sorption of Tc was measured. The Kd values for TcO4 were
0.1 ml/g for CGS-1 sediment at 35 days, 0.1 ml/g for TBS-1 sediment at 35 days, and 0.2 ml/g for
Trench-8 sediment at 44 days.
Sheppard and Thibault (1988) studied the vertical migration of technetium in two types of mires
typical of the Precambrian Shield in Canada. Sheppard and Thibault (1988) derived in situ Kd values
from analyses of the dried peat and pore water. The Kd values determined for technetium were
greater than 2 ml/g. Technetium was quickly immobilized in the reducing environment of the mire,
which was the cause for the observed in situ technetium Kd values.
Table 5.30. Average technetium Kd values (ml/g) (based on
three replicates) for technetium on Hanford sediments after
21.5 days from Gee and Campbell (1980).
Solution
Sample
Number
1
2
3
4
5
6
Silty Sediment
pH
8.26
8.07
8.52
7.78
7.90
7.75
Kd (ml/g)
-2.77
-1.13
-0.04
0.57
0.54
-0.51
Surface Sand
pH
8.14
8.00
8.25
7.63
7.57
7.44
Kd (ml/g)
0.07
-1.62
-0.31
0.06
0.52
0.38
Gee and Campbell (1980) conducted batch equilibration and unsaturated flow column studies of the
sorption behavior of technetium using two Hanford Site sediments and a series of synthetic
groundwater solutions. The sediments include a sample from the Ringold Formation (deep, old
sediment) and a surface soil sample at the Hanford Site. Except for the degree of crystallinity of the
clay minerals, the two sediments were similar with respect to their clay mineralogy. The
compositions of the synthetic groundwater solutions were developed to simulate the composition of
a solution as it changes in response to its initial contact and saturation of arid sediment and its
progressive percolation through the arid sediment. Gee and Campbell (1980) found low sorption of
technetium, and no relationship in the technetium Kd values with respect to soil type, solution
composition, pH, and contact time. The technetium Kd values for the two sediment samples ranged
from -2.34 to 1.27 ml/g at 8.5 days, and -2.77 to 0.57 ml/g at 21.5 days. The average batch Kj
values determined for the technetium-spiked sediments equilibrated with the six synthetic
groundwater solutions at 21.5 days are listed in Table 5.30. Although there was general agreement
5.75
-------
between the sorption measured by the batch and unsaturated column studies, Gee and Campbell
noted that the Kd values measured for technetium, using the unsaturated column studies were less
variable and better defined.
Wildung et al. (1974) used the batch equilibration method to measure the Kd values for TcO4 on
22 types of soils collected in Oregon, Washington, and Minnesota. The Kd values for TcO4 ranged
from 0.007 to 2.8 ml/g. Wildung et al. (1974) determined that the sorption of TcO4was positively
correlated to soil organic carbon, and negatively correlated to pH.
Routson et al. (1975, 1977) used batch equilibration experiments to measure Kd values for "Tc on
soil as a function of the concentrations of dissolved calcium bicarbonate. The soil sample was a
moderate-exchange capacity soil from South Carolina. The soil contained <0.2 mg/g CaCO3,
3.6 percent silt, 37.2 percent clay, and had a cation exchange capacity (CEC) of 2.5 meq/100 g and
saturated paste pH equal to 5.1. The measurements indicated that "Tc did not sorb to the South
Carolina soil sample. The measured Kd values for "Tc were +0.019, -0.052, -0.033, and +0.010 ml/g
for 0.002, 0.008, 0.020, and 0.200 M NaHCO3 solutions.
5.8.5.4- Published Compilations Containing JC, Values for Technetium
Because the references in this section are often cited or used for comparison in other publications,
the following summaries are provided for completeness. It is recommended that the reader review
the original reference and the references cited therein to understand the procedures and sources of
the Kd values used for each compilation. The compilations do not distinguish between oxidation
states for those contaminants that are redox sensitive or consider other important factors that
contribute to variability in sorption, such as pH. Moreover, in cases where very large Kd values are
listed, there is a risk that the original Kd measurement may have included precipitated components.
Baes and Sharp (1983) present a simple model developed for order-of-magnitude estimates for
leaching constants for solutes in agricultural soils. As part of this development, they reviewed and
determined generic default values for input parameters, such as Kd. A literature review was
completed to evaluate appropriate distributions for Kd values for various solutes, including
technetium. Because Baes and Sharp (1983) are cited frequently as a source of Kd values in other
published Kd reviews (e.g., Looney et al., 1987; Sheppard and Thibault, 1990), the technetium Kd
values listed by Baes and Sharp are reported here for completeness. Based on the distribution that
Baes and Sharp determined for the Kd values for cesium and strontium, they assumed a lognormal
distribution for the Kd values for all other elements in their compilation. Baes and Sharp listed an
estimated default Kd of 0.033 ml/g for technetium based on 24 Kd values that range from 0.0029
to 0.28 ml/g for agricultural soils and clays over the pH range 4.5 to 9.0. The 24 Kd values were
taken from Cast et al. (1979) which was not available for our review. Their compiled Kd values
represent a diversity of soils, pure clays (other Kd values for pure minerals were excluded), extracting
solutions, measurement techniques, and experimental error.
Looney et al. (1987) tabulated estimates for geochemical parameters needed for environmental
assessments of waste sites at DOE's Savannah River Plant in South Carolina. Looney et al. list Kd
values for several metal and radionuclide contaminants based on values that they found in 1-5
published sources. For technetium, Looney et al. list a "recommended" Kd of 0.001 ml/g, and a
range from 0.001 to 100 ml/g. Looney et al. note that their recommended values are specific to the
5.76
-------
Savannah River Plant site, and they must be carefully reviewed and evaluated prior to use in
assessments at other sites.
Thibault et al. (1990) (also see Sheppard and Thibault, 1990) present a compilation of soil Kd values
prepared to support radionuclide migration assessments for a Canadian geologic repository for spent
nuclear fuel in Precambrian Shield plutonic rock. Thibault et al. collected Kd values from other
compilations, journal articles, and government laboratory reports for important elements, such as
technetium, that would be present in the nuclear fuel waste inventory. The technetium Kj values
listed in Thibault et al. (1990) are included in Table 5.31. Thibault et al. (1990) describe the statistical
methods used for analysis of the compiled Kd values. The range for the Kd values used to calculate
the "geometric mean" cover several orders of magnitude. Readers are cautioned against using
"geometric mean" values or any other form of averaged Kd values as "default" Kd values, as such
values are usually calculated from data compiled from different investigators for different soil
systems. These mean or average values do not represent any particular environmental system and
geochemical conditions. As discussed in Volume I (EPA, 1999b), the variation observed in the
literature for Kd values for a contaminant is due to differences in sorption mechanisms, geochemical
conditions, soil materials, and methods used for the measurements.
Table 5.31. Technetium ^ values (ml/g) listed in
Thibault et al. (1990, Tables 4 to 8).
Soil Type
Sand
Silt
Clay
Organic
Kd Values (ml/g)
Geometric
Mean
0.1
0.1
1
1
Number of
Observations
19
10
4
24
Range
0.01 - 16
0.01 - 0.4
1.16-1.32
0.02 - 340
McKinley and Scholtis (1993) compare radionuclide Kd sorption databases used by different
international organizations for performance assessments of repositories for radioactive wastes. The
technetium Kd values listed in McKinley and Scholtis (1993, Tables 1, 2, and 4) are listed in
Table 5.32. The reader should refer to sources cited in McKinley and Scholtis (1993) for details
regarding their source, derivation, and measurement. Radionuclide Kd values listed for cementitious
environments in McKinley and Scholtis (1993, Table 3) are not included in Table 5.32. The
organizations listed in the tables include: AECL, GSF, IAEA, KBS, NAGRA, NIREX, NRG,
NRPB, PAGIS (CEC), PSE, RIVM, SKI, TVO, and UK DoE (acronyms defined in Section A.1.0 in
Appendix A).
5.77
-------
Table 5.32. Technetium Kd values (ml/g) listed by McKinley and
Scholtis (1993, Tables 1, 2, and 4) from sorption databases
used by different international organizations for performance
assessments of repositories for radioactive wastes.
Organization
AECL
GSF
IAEA
KBS-3
NAGRA
NIREX
NRC
NRPB
PAGIS
PAGIS SAFIR
PSE
RIVM
SKI
TVO
UK DoE
Argillaceous (Clay)
Sorbing
Material
Bentonite-Sand
Sediment
Pelagic Clay
Bentonite
Bentonite
Clay
Clay Mudstone
Clay, Soil Shale
Clay
Bentonite
Subseabed
Clay
Sediment
Sandy Clay
Bentonite
Lake Sediment
Clay
Coastal Marine
Water
Kd
(ml/g)
0.09
7
100
2
250
250
0
0.4
0
20
0
19
1
1
50
100
40
100
Crystalline Rock
Sorbing
Material
Granite
Granite
Granite
Granite
Basalt
Tuff
Granite
Crystalline
Rock, Reducing
Crystalline Rock
Kd
(ml/g)
26
50
250
0.4
0.4
0.4
5
50
0
Soil/Surface Sediment
Sorbing
Material
Soil/ Sediment
Soil/Sediment
Soil/ Sediment
Soil/Sediment
Soil/Sediment
Soil/ Sediment
Soil/Sediment
Soil/Sediment
Kd
(ml/g)
0.03
5
1
0
1.7
1,000
10
0.03
5.78
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5.8.5.5 Kj Studies of Technetium on Pure Mineral, Oxide, and Crushed Rock Materials
Numerous adsorption studies have been conducted of technetium on pure minerals, oxide phases,
and other geologic-related materials. The Kd values listed in these studies are not necessarily
relevant to the mobility and sorption of technetium in soils. However, they are listed in Appendix I
for completeness. The references cited in Appendix I are listed in the main reference list in
Chapter 6. The potential value of the references that they cite and the sorption processes that they
discuss is left to the reader. Like those for iodine, many of the studies that used on pure mineral and
oxide materials were conducted because of extensive research interest in developing getters
(adsorbents) that could be added to waste streams and tailored barriers for removal and/or
immobilization of dissolved iodine. The studies of technetium sorption on crushed rock were
conducted typically as part of national research programs to investigate the feasibility of geological
disposal of high-level radioactive waste (HLW).
5.9 Conclusions
One objective of this volume report is to provide a "thumb-nail sketch" of the geochemistry of
americium, arsenic, curium, iodine, neptunium, radium, and technetium. These contaminants
represent six nonexclusive contaminant categories: radionuclides, cations, anions, redox-sensitive,
non-retarded (non-attenuated) contaminants, and retarded (attenuated) contaminants (Table 5.33).
By categorizing the contaminants in this manner, general geochemical behaviors of one contaminant
may be extrapolated by analogy to other contaminants in the same category. For example, anions,
such as NO3~ and Cl", commonly adsorb to geological materials to a limited extent. This is also the
case observed for the sorption behavior of anionic iodide.
Important solution speciation, coprecipitation/dissolution, and adsorption reactions were discussed
for each contaminant. Where relevant, the distributions of aqueous species for certain contaminants
were calculated using the chemical equilibria code MINTEQA2 (Version 3.11, Allison et al., 1991)
for the water composition described in Table 5.1. The purpose of these calculations was to identify
the compositions of the dominant aqueous species that might exist in an oxidizing groundwater for
the contaminants reviewed in Volume III.. A summary of the results of these calculations is
presented in Table 5.34. The aqueous speciation of iodide, iodate, radium, and pertechnetate does
not change in the pH range of 3 and 10 under oxidizing conditions; they exist as I", IO3, Ra2+, and
TcO4, respectively. Americium(III), As(V), and Np(V) have two or three different dominant species
across this pH range. The aqueous speciation of arsenic, iodine, neptunium, and technetium is
especially complicated, because these contaminants may exist in multiple oxidation states over the
range of pH-Eh conditions of natural waters.
One general conclusion that can be made from the calculations in Table 5.34 is that, as the pH
increases, the aqueous complexes tend to become increasingly more negatively charged. For
example, americium and neptunium are cationic at pH 3. At pH values greater than 7, they exist
predominantly as either neutral or anionic species. Negatively charged complexes tend to adsorb
less to soils than their respective cationic species. This rule-of-thumb stems from the fact that most
minerals in soils have a net negative charge in the range of pH values of most natural waters.
Conversely, the solubility of several of these contaminants decreases dramatically as pH increases.
Therefore, the net contaminant concentration in solution does not necessarily increase as the
dominant aqueous species becomes more negatively charged.
5.79
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Table 5.33. Selected chemical and transport
properties of the contaminants.
Element
Am
As
Cm
I
Np
Ra
Tc
Radio-
nuclide1
X
X
X
X
X
X
Primary Species at pH 7
and Oxidizing Conditions
Cationic
X
X
X
X
Anionic
X
X
X
Neutral
Redox
Sensitive2
X
X
X
X
Transport Through
Soils at pH 7
Not
Retarded3
X
Retarded3
(Attenuated)
X
x (limited)
X
x (limited)
x (limited)
x
1 Contaminants that are primarily a health concern as a result of their radioactivity are identified in this
column.
2 The redox status column identifies contaminants (As, I, and Tc) that have variable oxidation
states within the pH and Eh limits commonly found in the environment.
3 Retarded or attenuated (nonconservative) transport means that the contaminant moves slower
than water through geologic material. Nonretarded or nonattenuated (conservative) transport means that
the contaminant moves at the same rate as water.
Another objective of this report is to identify the important chemical, physical, and mineralogical
characteristics controlling sorption of these contaminants. There are several aqueous- and solid-
phase characteristics that can influence contaminant sorption. These characteristics commonly
have an interactive effect on contaminant sorption, such that the effect of one parameter on
sorption varies as the magnitude of other parameters changes. A list of some of the more
important chemical, physical, and mineralogical characteristics affecting contaminant sorption are
listed in Table 5.35.
5.80
-------
Table 5.34. Distribution of dominant aqueous species for
each contaminant at three pH values for an oxidizing
water described in Table 5.11.
Element
AmfTII)
As(V)
Cm (III)2
I(-I)
I(V)
Np(V)
Ra
Tc(VH)
pH3
Species
Am3+
AmF2+
AmSO4
H2AsO4
H3AsO° (aq)
Cm3+
r
I03
NpO2
Ra2+
Tc04
%
65
20
14
86
14
100
100
100
100
100
100
PH7
Species
AmCO+
Am(OH)2+
Am(CO3)2
Am3+
AmF2+
HAsO2-
H2AsO4
CmCO+
Cm3+
Cm(OH)2+
r
I03
NP0+
NpO2(CO3)-
NpO2(OH)° (aq)
Ra2+
Tc04
%
88
4
3
2
2
66
34
43
42
15
100
100
97
1
1
100
100
pHIO
Species
Am(OH)°3 (aq)
Am(CO3)2
AmCO3
HAsO2-
AsO3-
Cm(OH)+
CmCO+
Cm(OH)2+
Cm(CO3)2
r
I03
NpO2(OH)° (aq)
Np02(C03)-
NpO+
NP02(C03)3-
Ra2+
Tc04
%
56
42
2
97
3
86
8
4
2
100
100
51
41
4
4
100
100
1 Only species comprising 1 percent or more of the total contaminant distribution are presented. Hence,
the total of the percent distributions presented in table may not always equal 100 percent.
2 Thermodynamic constants for aqueous species and solids containing curium are extremely limited. By
analogy, the aqueous speciation of curium is expected to be very similar to Am(III). Aqueous speciation
calculated for Cm(III) is based on a limited set of constants.
The effect of pH on both adsorption and (co)precipitation of contaminants is pervasive. The pH
influences a number of aqueous and solid phase properties that directly affect the sorption to some
degree of each contaminant reviewed in Volumes II and III. These effects are summarized in
Volume I (EPA, 1999b), and discussed in greater detail for each contaminant in the individual
geochemistry review sections in Volumes II and III. As discussed above, pH has a significant effect
on the aqueous speciation of most contaminants (Table 5.34). The adsorption behavior of each
contaminant depends on the ionic charge and composition of its important aqueous species present
in the environmental system being studied. Additionally, pH affects the number of adsorption sites
on variable-charged minerals (aluminum- and iron-oxide minerals), partitioning of contaminants to
organic matter, CEC, formation of aqueous complexes, oxidation state of contaminants and
complexing/precipitating ligands, and H+/OPT-competition for anionic or cationic adsorption sites.
5.81
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Table 5.35 lists some of the more important aqueous- and solid phase parameters affecting
contaminant sorption.
Table 5.35. Some of the more important aqueous- and
solid-phase parameters affecting contaminant sorption.
Element
Am
As
Cm
I
Np
Ra
Tc
Important Aqueous- and Solid-Phase Parameters Influencing
Contaminant Sorption1
Clay Minerals, Iron-/Aluminum-Oxide Minerals, pH
Clay Minerals, Dissolved Phosphate, Iron-/Aluminum-Oxide Minerals,
pH, Redox
Clay Minerals, Iron-/Aluminum-Oxide Minerals, pH
Dissolved Halides, Organic Matter, Redox, Volatilization, pH
Clay Minerals, Iron-/Aluminum-Oxide Minerals, pH
BaSO4 Coprecipitation, Dissolved Alkaline Earth Elements, Cation
Exchange Capacity, Clay Minerals, Ionic Strength, Iron-/Aluminum-
Oxide Minerals, Organic Matter, pH
Organic Matter, Redox
1 Parameters listed in alphabetical order.
Unlike most ancillary parameters, the redox status of a system also dramatically influences the
sorption of several contaminants included in this study. If the bulk redox potential of a soil/water
system is above the potential of the specific element redox reaction, the oxidized form of the redox
sensitive element will exist. Below this critical value, the reduced form of the element will exist.
Such a change in redox state can alter Kd values by several orders of magnitude, and can have direct
and indirect effects on contaminant coprecipitation. The direct effect occurs with contaminants like
technetium where the oxidized species form more soluble solid phases than the reduced species.
The indirect effects occur when the contaminants adsorb to redox sensitive solid phases or
precipitate with redox sensitive ligands. An example of the former involves the reductive dissolution
of ferric oxide minerals, which can adsorb (complex) metals strongly. As the ferric oxide minerals
dissolve, the adsorption potential of the soil is decreased. Another indirect effect of redox on
contaminant sorption involves sulfur-ligand chemistry. Under reducing conditions, S(VI) (SO^,
sulfate) will convert into S(II) (S2~, sulfide) and then the S(II) may form sparingly soluble sulfide
minerals. Thus, these two redox sensitive reactions may have off-setting net effects on total
contaminant sorption (sulfide precipitates may sequester some of the contaminants previously
bound to ferric oxides).
The greatest limitation of using a Kd value to calculate a retardation term for contaminant transport
modeling is that it is only applicable to a single set of environmental conditions. The Kd values
reported in the literature for any given contaminant may vary by as much as six orders of magnitude.
One of the major recommendations of this three-volume report is that for site-specific calculations,
5.82
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Kd values measured at site-specific conditions are absolutely essential. To select appropriate Kd
value(s) for contaminant transport modeling, it is incumbent upon technical staff to understand (1)
the strengths, weaknesses, and underlying assumptions of the different Kd methods (see EPA,
1999b); and (2) the important geochemical processes and knowledge of the important ancillary
parameters affecting the sorption chemistry of each contaminant of interest. This volume, which is
an extension of Volume II (EPA, 1999c), provides "thumb-nail sketches" of the important aqueous
speciation, (co)precipitation/dissolution, and adsorption processes affecting the sorption of
americium, arsenic, curium, iodine, neptunium, radium, and technetium under oxidizing conditions.
5.83
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6.26
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6.27
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APPENDIX A
Acronyms, Abbreviations, Symbols, and Notation
-------
Appendix A
Acronyms, Abbreviations, Symbols, and Notation
A. 1 Acronyms and Abbreviations
AA
AEG
AECL
ASCII
ASTM
CCM
CDTA
CEAM
CEC
CEC
DDLM
DIRM
DLM
DOE
DTPA
EDTA
EDS
EDX
EPA
EXAFS
GSF
GTLM
HEDTA
HFO
HLW
IAEA
ICP
ICP/MS
IEP (or icp)
IR
KBS
LANL
LLNL
LLW
MCL
MEPAS
MS-DOS®
Atomic absorption
Anion exchange capacity
Atomic Energy of Canada Limited
American Standard Code for Information Interchange
American Society for Testing and Materials
Constant capacitance (adsorption) model
Trans-l,2-diaminocyclohexane tetra-acetic acid
Center for Exposure Assessment Modeling (at EPA's Environmental
Research Laboratory in Athens, Georgia)
Commission of European Communities
Cation exchange capacity
Diffuse double layer (adsorption) model
Dissimilatory iron-reducing bacteria
Diffuse (double) layer (adsorption) model
U.S. Department of Energy
Diethylenetriaminepentacetic acid
Ethylenediaminetriacetic acid
Energy dispersive spectroscopy
Energy dispersive x-ray analysis
U.S. Environmental Protection Agency
X-ray absorption fine structure spectroscopy
Gesellschaft fur Strahlen- und Umweltforschung m.b.H., Germany
Generalized two-layer surface complexation model
N-(2-hydroxyethyl) ethylenedinitrilotriacetic acid
Hydrous ferric oxide
High-level radioactive waste
International Atomic Energy Agency, Vienna, Austria
Inductively coupled plasma
Inductively coupled plasma/mass spectroscopy
Isoelectric point
Infrared
Swedish Nuclear Safety Board
Los Alamos National Laboratory
Lawrence Livermore National Laboratory
Low-level radioactive waste
Maximum Contaminant Level
Multimedia Environmental Pollutant Assessment System
Microsoft® disk operating system (Microsoft and MS-DOS are register
trademarks of Microsoft Corporation.)
A.2
-------
NAGRA
NEA
NBS
NIREX
NIST
NPL
NRG
NRPB
NWWA
OECD
PAGIS
PAGIS SAFIR
PC
PNL
PNNL
PSE
REE
RIVM
SCM
SDMP
SEM
SKI
TDS
TLM
TOG
TVO
UK
UK DoE
UNSCEAR
USDA
WHO
XANES
Nationale Genossenschaft fur die Lagerung radioaktiver AbfalleSwiss
(National Cooperation for Storage of Radioactive Waste), Switzerland
OECD Nuclear Energy Agency, France
U.S. National Bureau of Standards
United Kingdom Nirex Ltd.
U.S. National Institute of Standards and Technology
Superfund National Priorities List
U.S. Nuclear Regulatory Commission
National Radiological Protection Board, United Kingdom
National Water Well Association
Organisation for Economic Co-operation and Development, France
Performance Assessment of Geological Isolation Systems, Commission of
the European Communities (CEC), Belgium
PAGIS Safety Assessment and Feasibility Interim Report
Personal computers (operating under the MS-DOS® and Microsoft®
Windows operating systems, Microsoft® Windows is a trademark of
Microsoft Corporation.)
Pacific Northwest Laboratory. In 1995, DOE formally changed the name of
the Pacific Northwest Laboratory to the Pacific Northwest National
Laboratory.
Pacific Northwest National Laboratory
Projekt Sicherheitsstudien Entsorgung, Germany
Rare earth element
Rijksinstituut voor Volksgezondheid en Milieuhygience (National Institute of
Public Health and Environment Protection), Netherlands
Surface complexation model
NRC's Site Decommissioning Management Plan
Scanning electron microscopy
Statens Karnkraftinspektion (Swedish Nuclear Power Inspectorate)
Total dissolved solids
Triple-layer adsorption model
Total organic carbon
Teollisuuden Voima Oy (Industrial Power Company), Finland
United Kingdom (UK)
United Kingdom Department of the Environment
United Nations Scientific Committee on the Effects of Atomic Radiation
U.S. Department of Agriculture
World Health Organization
X-ray absorption near-edge structure spectroscopy
A.3
-------
A.2 List of Symbols for the Elements and Corresponding Names
Symbol
Ac
Ag
Al
Am
Ar
As
At
Au
B
Ba
Be
Bi
Bk
Br
C
Ca
Cd
Ce
Cf
Cl
Cm
Co
Cr
Cs
Cu
Dy
Er
Es
Eu
F
Fe
Fm
Fr
Ga
Gd
Element
Actinium
Silver
Aluminum
Americium
Argon
Arsenic
Astatine
Gold
Boron
Barium
Beryllium
Bismuth
Berkelium
Bromine
Carbon
Calcium
Cadmium
Cerium
Californium
Chlorine
Curium
Cobalt
Chromium
Cesium
Copper
Dysprosium
Erbium
Einsteinium
Europium
Fluorine
Iron
Fermium
Francium
Gallium
Gadolinium
Symbol
Ge
H
He
Hf
Hg
Ho
I
In
Ir
K
Kr
La
Li
Lu
Lr
Md
Mg
Mn
Mo
N
Na
Nb
Nd
Ne
Ni
No
Np
O
Os
P
Pa
Pb
Pd
Pm
Po
Element
Germanium
Hydrogen
Helium
Hafnium
Mercury
Holmium
Iodine
Indium
Iridium
Potassium
Krypton
Lanthanum
Lithium
Lutetium
Lawrencium
Mendelevium
Magnesium
Manganese
Molybdenum
Nitrogen
Sodium
Niobium
Neodymium
Neon
Nickel
Nobelium
Neptunium
Oxygen
Osmium
Phosphorus
Protactinium
Lead
Palladium
Promethium
Polonium
Symbol
Pr
Pt
Pu
Ra
Rb
Re
Rh
Rn
Ru
S
Sb
Sc
Se
Si
Sm
Sn
Sr
Ta
Tb
Tc
Te
Th
Ti
Tl
Tm
U
V
w
Xe
Y
Yb
Zn
Zr
Element
Praseodymium
Platinum
Plutonium
Radium
Rubidium
Rhenium
Rhodium
Radon
Ruthenium
Sulfur
Antimony
Scandium
Selenium
Silicon
Samarium
Tin
Strontium
Tantalum
Terbium
Technetium
Tellurium
Thorium
Titanium
Thallium
Thulium
Uranium
Vanadium
Tungsten
Xenon
Yttrium
Ytterbium
Zinc
Zirconium
A.4
-------
A.3 List of Symbols and Notation
pb Porous media bulk density (mass/length3)
A Angstrom, 1040 meters
ads Adsorption or adsorbed
AEC Anion exchange capacity
Aj Concentration of adsorbate (or species) i on the solid phase at equilibrium
am Amorphous
aq Aqueous
CEC Cation exchange capacity
Ci Curie
d Day
dpm Disintegrations per minute
e" Free electron
Eh Redox potential of an aqueous system relative to the standard hydrogen electrode
F Faraday constant, 23,060.9 cal/V'mol
g Gram
3H Tritium
h Hour
I Ionic strength
TAP Ion activity product
IEP Isoelectric point
Kd Concentration-based partition (or distribution) coefficient
Kj.^98 Equilibrium constant at 298 K
KrT Equilibrium constant at temperature T
1 Liter
M Molar
m Meter
mCi Millicurie, 10~3 Curies
meq Milliequivalent
mi Mile
ml Milliliter
mol Mole
mV Millivolt
N Constant in the Freundlich isotherm model
n Total porosity
ne Effective porosity
pCi Picocurie, 1042 Curies
pE Negative common logarithm of the free-electron activity
pH Negative logarithm of the hydrogen ion activity
pH pH for zero point of charge
ppm Parts per million
R Ideal gas constant, 1.9872 cal/mol
Rf Retardation factor
s Solid phase species
sec Second
A. 5
-------
SI Saturation index, as defined by log
T Absolute temperature, usually in Kelvin unless otherwise specified
t Time
t,/2 Half life
TDS Total dissolved solids
TU Tritium unit which is equivalent to 1 atom of 3H (tritium) per 1018 atoms
of *H (protium)
vc Velocity of contaminant through a control volume
vp Velocity of the water through a control volume
y Year
Z Valence state
z Charge of ion
{ } Activity
[ ] Concentration
A.6
-------
APPENDIX B
Definitions
-------
Appendix B
Definitions
Adsorption - partitioning of a dissolved species onto a solid surface.
Adsorption Edge - the pH range where solute adsorption sharply changes from ~10 percent to
~90 percent.
Actinon - name occasionally used, especially in older documents, to refer to 219Rn which forms from
the decay of actinium.
Activity - the effective concentration on an ion that determines its behavior to other ions with
which it might react. An activity of ion is equal to its concentration only in infinitely dilute
solutions. The activity of an ion is related to its analytical concentration by an activity
coefficient, Y-
Alkali Metals - elements in the 1A Group in the periodic chart. These elements include lithium,
sodium, potassium, rubidium, cesium, and francium.
Alpha Particle - particle emitted from nucleus of atom during 1 type of radioactive decay. Particle
is positively charged and has 2 protons and 2 neutrons. Particle is physically identical to the
nucleus of the 4He atom (Bates and Jackson 1980).
Alpha Recoil - displacement of an atom from its structural position, as in a mineral, resulting from
radioactive decay of the release an alpha particle from its parent isotope (e.g., alpha decay of 222Rn
from 226Ra).
Amphoteric Behavior - the ability of the aqueous complex or solid material to have a negative,
neutral, or positive charge.
Basis Species - see component species.
Cation Exchange - reversible adsorption reaction in which an aqueous species exchanges with an
adsorbed species. Cation exchange reactions are approximately stoichiometric and can be
written, for example, as
CaX(s) + ^Sr^Caq) = 90SrX(s) + Ca2+(aq)
where X designates an exchange surface site.
Cation Exchange Capacity (CEC) - the sum total of exchangeable cations per unit mass of
soil/sediment that a soil can adsorb.
B.2
-------
Chernozem - A great soil group of the 1938 USDA classification system. It is a group of zonal
soils whose surface horizon is dark and highly organic, below which is a lighter-colored horizon
and accumulation of lime. The soil group is developed under conditions of temperate to cool
subhumid climate (USDS, 1938).
Clay Content - particle size fraction of soil that is less than 2 |J,m (unless specified otherwise).
Code Verification - test of the accuracy with which the subroutines of the computer code perform
the numerical calculations.
Colloid - any fine-grained material, sometimes limited to the particle-size range of <0.00024 mm
(i.e., smaller than clay size), that can be easily suspended. In its original sense, the definition of a
colloid included any fine-grained material that does not occur in crystalline form.
Complexation (Complex Formation) - any combination of dissolved cations with molecules or
anions containing free pairs of electrons.
Component Species - "basis entities or building blocks from which all species in the system can be
built" (Allison et al. 1991). They are a set of linearly independent aqueous species in terms of
which all aqueous speciation, redox, mineral, and gaseous solubility reactions in the
MINTEQA2 thermodynamic database are written.
Detrital Mineral - "any mineral grain resulting from mechanical disintegration of parent rock"
(Bates and Jackson 1980).
Deuterium (D) - stable isotope 2H of hydrogen.
Diagenetic - "caused by the chemical, physical, and biological changes undergone by a sediment
after its initial deposition, and during and after its consolidation into a coherent solid rock,
exclusive of surficial weathering and metamorphism" (Bates and Jackson 1980).
Disproportionation - is a chemical reaction in which a single compound serves as both oxidizing
and reducing agent and is thereby converted into more oxidized and a more reduced derivatives
(Sax and Lewis 1987). For the reaction to occur, conditions in the system must be temporarily
changed to favor this reaction (specifically, the primary energy barrier to the reaction must be
lowered). This is accomplished by a number of ways, such as adding heat or microbes, or by
radiolysis occurring. Examples of plutonium disproportionation reactions are:
3Pu4+ + 2H2O = 2Pu3+ + PuO^+ +4H+
3PuC>2 + 4H+ = Pu3+ + 2PuO22+ +2H2O.
Electron Activity - unity for the standard hydrogen electrode.
Evaporite - A sedimentary rock composed principally of minerals precipitated from a saline
solution as a result of extensive or total evaporation of the solution.
B.3
-------
Far Field - the portion of a contaminant plume that is far from the point source and whose
chemical composition is not significantly different from that of the uncontaminated portion of
the aquifer.
Fulvic Acids - breakdown products of cellulose from vascular plants (also see humic acids). Fulvic
acids are the alkaline-soluble portion which remains in solution at low pH and is of lower
molecular weight (Gascoyne 1982).
Humic Acids - breakdown products of cellulose from vascular plants (also see fulvic acids). Humic
acids are defined as the alkaline-soluble portion of the organic material (humus) which
precipitates from solution at low pH and are generally of high molecular weight (Gascoyne
1982).
Hydrolysis - a chemical reaction in which a substance reacts with water to form two or more new
substances. For example, the first hydrolysis reaction of U4+ can be written as
U4+ + H2O = UOH3+ + FT.
Hydrolytic Species - an aqueous species formed from a hydrolysis reaction.
Ionic Potential - ratio (z/r) of the formal charge (z) to the ionic radius (r) of an ion.
Isoelectric Point (iep) - pH at which a mineral's surface has a net surface charge of zero. More
precisely, it is the pH at which the particle is electrokinetically uncharged.
Lignite - a coal that is intermediate in coalification between peat and subbituminous coal.
Marl - an earthy substance containing 35 to 65 percent clay and 65 to 5 percent carbonate formed
under marine or freshwater conditions
Mass Transfer - transfer of mass between 2 or more phases that includes an aqueous solution, such
as the mass change resulting from the precipitation of a mineral or adsorption of a metal on a
mineral surface.
Mass Transport - time-dependent movement of 1 or more solutes during fluid flow.
Methylation - a chemical process for introducing a methyl group (CH3—) into a species
Mire - a small piece of marshy, swampy, or boggy ground.
Model Validation - integrated test of the accuracy with which a geochemical model and its
thermodynamic database simulate actual chemical processes.
Monomeric Species - an aqueous species containing only one center cation (as compared to a
polymeric species).
B.4
-------
Near Field - the portion of a contaminant plume that is near the point source and whose chemical
composition is significantly different from that of the uncontaminated portion of the aquifer.
Peat - an unconsolidated deposit of semicarbonized plant remains in a water saturated environment.
pHzpc - As the pH increases, mineral surfaces become increasingly more negatively charged. The
pH where the surface is uncharged (i.e., zero net charge) is referred to as the pH of zero-point-
of-charge, pH2pc.
Polynuclear Species - an aqueous species containing more than one central cation moiety, e.g.,
(UO2)2CO3(OH)3 and P
Podzol - A great soil group of the 1938 USDA classification system. It is a group of zonal soils
having an organic mat and a very thin organic-mineral layer overlying gray, leached A2 horizon
and a dark brown, illuvial B horizon enriched in iron oxide, alumina, and organic matter. It
develops under coniferous or mixed forests or under heath, in a cool to temperate moist climate
(USDA, 1938).
Protium (H) - stable isotope ^ of hydrogen.
Retrograde Solubility - solubility that decreases with increasing temperature, such as those of
calcite (CaCO3) and radon. The solubility of most compounds (e.g., salt, NaCl) increases with
increasing temperature.
Species - actual form in which a dissolved molecule or ion is present in solution.
Specific Adsorption - surface complexation via a strong bond to a mineral surface. For example,
several transition metals and actinides are specifically adsorbed to aluminum- and iron-oxide
minerals.
Sol - a homogeneous suspension or dispersion of colloidal matter in a fluid.
Solid Solution - a solid material in which a minor element is substituted for a major element in a
mineral structure.
Thoron - name occasionally used, especially in older documents, to refer to 220Rn which forms from
the decay of thorium.
Tritium (T) - radioactive isotope 3H of hydrogen.
Tritium Units - units sometimes used to report tritium concentrations. A tritium unit (TU) is
equivalent to 1 atom of 3H (tritium) per 1018 atoms of1!! (protium). In natural water that
produces 7.2 x 1CT3 disintegrations per minute per milliliter (dpm/ml) of tritium, 1 TU is
approximately equal to 3.2 picocuries /milliliter (pCi/ml).
zpc (zero-point-of-charge) - see pH .
B.5
-------
APPENDIX C
Americium Adsorption Studies
-------
Appendix C
Americium Adsorption Studies
The Kd values listed in the following pure mineral/oxide and crushed rock studies are not
necessarily relevant to the mobility and sorption of americium in soils. However, the studies and
associated references are listed below for completeness. The potential value of these references and
the sorption processes that were investigated is left to the reader. The references cited in this
appendix are listed in the main reference list in Chapter 6.
C.1 Adsorption Studies of Americium on Single Mineral Phases
Several studies have been conducted for the adsorption of americium on pure mineral and oxide
phases. Some of these studies are listed in Tables C.la and C.lb.
Degueldre eta!. (1994) investigated the role of dissolved carbonate concentrations on the sorption
and desorption of americium on montmorillonite, illite, and hematite colloids. Strong adsorption of
americium on the mineral colloids was observed. They concluded that in addition to the
uncomplexed aqueous ion Am3+, americium hydroxyl, carbonate, and carbonate-hydroxyl
complexes also sorbed to the minerals. A small decrease in americium adsorption was observed at
pH 8 in experiments conducted with greater than 0.02 M total dissolved carbonate. Under these
pH/carbonate conditions, the anion Am(CO3)f dominates the speciation of dissolved americium.
Degueldre eta!. (1994) concluded that the formation of americium carbonate and hydroxyl-
carbonate aqueous complexes compete with the formation of americium surface complexes on the
mineral colloids. At pH 10, this decreasing trend ends due to the formation and adsorption of
strong aqueous americium hydroxyl complexes.
Samadfam eta!. (2000) studied the sorption of Am(III) on kaolinite as a function of pH from 3.5
to 10 and humic acid concentrations over the range 0 to 20 ppm in 0.1 M NaClO4 solutions. In the
absence of humic acid, the adsorption of Am(III) increased with pH from 3.5 to 10. The presence
of humic acid increased the adsorption of Am(III) to kaolinite at pH values less than 5 and
decreased adsorption at higher pH values.
Beall eta!. (1978) determined the sorption behavior of Am(III) on attapulgite, montmorillonite, and
kaolinite in NaCl brines buffered at pH 5. Their results indicated that the Am(III) adsorption on
these minerals decreased with increasing ionic strength.
C.2
-------
Table C.1 Americium adsorption studies on pure
mineral and oxide phases.
Mineral/Oxide
Albite
Almandine
Anhydrite
Anorthite
Apatite
Attapulgite
Augite
Bentonite
Beryl
Biotite
Bytownite
Calcite
Chalcopyrite
Chalcocite
Clay (abyssal red clay)
Clinoptolite
Corundum
Dolomite
Epidote
Fluorite
Galena
Gibbsite
Goethite
References
Allard (1984)
Allard (1984)
Allard (1984)
Allard (1984)
Allard (1984), Andersson et al. (1982)
Allard (1984), Andersson et al. (1982), Beall et al. (1978)
Allard (1984)
Nagasaki et al. (1994)
Allard (1984)
Ticknor et al. (1996), Allard (1984), Andersson et al. (1982)
Allard (1984)
Allard (1984), Andersson et al. (1982)
Allard (1984)
Allard (1984)
Enckson (1980)
Triay et al. (1991)
Allard (1984)
Brady et al. (1999), Allard (1984)
Allard (1984)
Allard (1984), Andersson et al. (1982)
Allard (1984)
Allard (1984)
Ticknor et al. (1996)
C.3
-------
Continuation of Table C.1
Mineral/Oxide
Halloysite
Hematite
Hornblende
Illite
Kaolinite
Limonite
Magnetite
Microcline
Molybdenite
Monazite
Montmorillonite
Muscovite
Olivine
Pyrite
Quartz
Romanechite
Serpentine
Sphene
Zircon
References
Allard (1984)
Degueldre et al. (1994), Allard (1984)
Allard (1984)
Degueldre et al. (1994)
Samadfam et al. (2000), Allard (1984), Beall et al. (1978)
Allard (1984)
Allard (1984), Andersson et al. (1982)
Allard (1984)
Allard (1984)
Allard (1984)
Ticknor et al. (1996), Degueldre et al. (1994), Allard (1984), Andersson
etal. (1982), Beall et al. (1978)
Allard (1984)
Allard (1984), Andersson etal. (1982)
Allard (1984)
Ticknor etal. (1996), Allard (1984), Andersson etal. (1982)
Tnay etal. (1991)
Allard (1984)
Allard (1984)
Allard (1984)
Degueldre et al. (1994) investigated the role of dissolved carbonate concentrations on the sorption
and desorption of americium on montmorillonite, illite, and hematite colloids. Strong adsorption of
americium on the mineral colloids was observed. Degueldre et al. (1994) concluded that in addition
to the uncomplexed aqueous ion Am3+, americium hydroxyl, carbonate, and carbonate-hydroxyl
complexes also sorbed to the minerals. A small decrease in americium adsorption was observed at
pH 8 in experiments conducted with greater than 0.02 M total dissolved carbonate. Under these
pH/carbonate conditions, the anion Am(CO3)f dominates the speciation of dissolved americium.
Degueldre et al. (1994) concluded that the formation of americium carbonate and hydroxyl-
carbonate aqueous complexes compete with the formation of americium surface complexes on the
C.4
-------
mineral colloids. At pH 10, this decreasing trend ends due to the formation and adsorption of
strong aqueous americium hydroxyl complexes.
Samadfam et al. (2000) studied the sorption of Am(III) on kaolinite as a function of pH from 3.5
to 10 and humic acid concentrations over the range 0 to 20 ppm in 0.1 M NaClO4 solutions. In the
absence of humic acid, the adsorption of Am(III) increased with pH from 3.5 to 10. The presence
of humic acid increased the adsorption of Am(III) to kaolinite at pH values less than 5 and
decreased adsorption at higher pH values.
Beall et al. (1978) determined the sorption behavior of Am(III) on attapulgite, montmorillonite, and
kaolinite in NaCl brines buffered at pH 5. Their results indicated that the Am(III) adsorption on
these minerals decreased with increasing ionic strength.
C.2 Kd Studies of Americium on Crushed Rock Materials
Some studies have been conducted for the sorption of americium on crushed rock as part of
national research programs to investigate the feasibility of geological disposal of high-level
radioactive waste (HLW). Some of these studies are listed in Table C.2.
Table C.2 Americium adsorption studies on
crushed rock and related materials.
Geologic Material
Basalt
Diabase
Gneiss
Granite
Sandstone
Shale
Tuff
References
Allard (1984), Salter et al., 1981a), Ames and McGarrah, 1980)), Silva et al.,
1979),
Andersson et al. (1982)
Andersson et al. (1982)
Ticknor et al. (1996), Allard (1984), Andersson et al. (1982), Silva et al.
(1979)
Barney (1982)
Silva etal (1979)
Tnay etal. (1991), Barney (1982)
C.5
-------
APPENDIX D
Arsenic Adsorption Studies
-------
Appendix D
Arsenic Adsorption Studies
The Kd values listed in the following pure mineral/oxide and crushed rock studies are not
necessarily relevant to the mobility and sorption of arsenic in soils. However, the studies and
associated references are listed below for completeness. The potential value of these references and
the sorption processes that were investigated is left to the reader. The references cited in this
appendix are listed in the main reference list in Chapter 6.
Several studies have been conducted on the adsorption to arsenic on pure mineral and oxide phases.
Some of these studies are listed in Table D.la and Dl.b.
Table D.1 Arsenic adsorption studies on
pure mineral and oxide phases.
Mineral/Oxide
Alumina (activated)
Aluminum hydroxide
(amorphous)
Bauxite (activated)
Biotite (Hydroxy
Al-coated)
Corundum
Ettringite
Ferrihydrite
Gibbsite
Goethite
Hydrous iron oxide
Illite
References
Gupta and Chen (1978)
Manning and Goldberg [1997a, As(III)], Davis (1978), Anderson et al.
(1976)
Gupta and Chen (1978)
Huang (1975)
Halter and Pfeider [2001, As(V)]
Mynemefal. [1997, As (V)]
Jam etal. [1999, As(V) and As (III)], Raven et al. [1998, As(V) and
As(III)], Fuller etal. [1993, As (III)], Waychunas etal. [1992, As(V)],
Davis etal. [1989, As (V)]
Ladeirae^/. (2001)
O'Reilly etal. [2001, As(V)], Manning etal. [1998, As(III)], Sun and
Doner [1998, As (III)], Fendorf etal. [1997, As(V)]
Wilkie and Henng [1996, As(V) and As(III)], Hsi etal. (1994)
Manning and Goldberg [1997a, As (III)]
D.2
-------
Continuation of Table D.1.
Mineral/Oxide
Iron oxyhydroxide
(amorphous)
Kaolinite
Montmorillonite
Muscovite (Hydroxy
Al-coated)
Silica (amorphous)
References
Belzile and Tessier (1990), Benjamin and Bloom (1981), Pierce and
Moore [1982, As(V) and As(III)], Leckie et aL (1980)
Lin and Puls [2000, As(V) and As(III)], Takahashi et d. [1999, As(V)],
Manning and Goldberg [1997a, As(III)], Frost and Griffin [1977,
As (V) and As (III)]
Manning and Goldberg [1997a, As (III)], Frost and Griffin [1977,
As (V) and As (III)]
Huang (1975)
Takahashi et al [1999, As(V)]
D.3
-------
APPENDIX E
Curium Adsorption Studies
-------
Appendix E
Curium Adsorption Studies
The Kd values listed in the following pure mineral/oxide and crushed rock studies are not
necessarily relevant to the mobility and sorption of curium in soils. However, the studies and
associated references are listed below for completeness. The potential value of these references and
the sorption processes that were investigated is left to the reader. The references cited in this
appendix are listed in the main reference list in Chapter 6.
E.1 Adsorption Studies of Curium on Single Mineral Phases
Few studies have been conducted on the adsorption of curium on pure mineral, oxide phases, and
other geologic-related materials. Some of these studies are listed in Table E.I.
Samadfam eta!. (2000) studied the sorption of Cm(III) on kaolinite as a function of pH over the
range 3.5 to 10 and humic acid concentrations over the range 0 to 20 ppm in 0.1 M NaClO4
solutions. In the absence of humic acid, the adsorption of Cm(III) increased with pH over the range
3.5 to 10. The presence of humic acid increased the adsorption of Cm(III) to kaolinite at pH values
less than 5, and decreased adsorption at high pH values.
Beall et al. (1978) determined the sorption behavior of Cm(III) on attapulgite, montmorillonite, and
kaolinite in NaCl brines buffered at pH 5. Their results indicated that the Cm(III) adsorption on
these minerals decreased with increasing ionic strength.
Table E.1 Curium adsorption studies on
pure mineral and oxide phases.
Mineral/Oxide
Attapulgite
Bentonite
Clay (abyssal red clay)
Goethite
Montmorillonite
Kaolinite
Silica gel
References
Beall etal. (1978)
Baston^rf/. (1997)
Erickson (1980)
Alberts etal. (1986)
Beall etal (1978)
Samadfam etal. (2000), Alberts etal. (1986),
Beall etal (1978)
Alberts etal (1986)
E.2
-------
E.2 Kd Studies of Curium on Crushed Rock Materials
Few studies have been conducted for the sorption of curium on crushed rocks. Studies identified
were conducted as part of national research programs to investigate the feasibility of geological
disposal of high-level radioactive waste (HLW). Some of these studies are listed in Table E.2.
Table E.2 Curium adsorption studies on
crushed rock and related materials.
Geologic Material
Basalt
Granite
Granodiorite
Shale
Tuff
References
Silvaef al. (1979)
Silvaef al. (1979)
Bastonetal. (1997)
Silvaef al. (1979)
Bastonetal. (1997)
E.:
-------
APPENDIX F
Iodine Adsorption Studies
-------
Appendix F
Iodine Adsorption Studies
The Kd values listed in the following pure mineral/oxide and crushed rock studies are not
necessarily relevant to the mobility and sorption of iodine in soils. However, the studies and
associated references are listed below for completeness. The potential value of these references and
the sorption processes that were investigated is left to the reader. The references cited in this
appendix are listed in the main reference list in Chapter 6.
F.1 Adsorption Studies of Iodine on Single Mineral Phases
Numerous studies have been conducted of the adsorption of iodine on pure mineral, oxide phases,
and other geologic-related materials. Many of these studies were conducted because of extensive
research interest in developing getters (adsorbents) that could be added to waste streams and tailored
barriers for removal and/or immobilization of dissolved iodine. Some of these studies are listed in
Tables F.la, F.lb, and F.lc.
Balsley eta!. (1996) used potentiometric titrations to investigate the sorption of I" on cinnabar (HgS)
and chalcocite (Cu2S) in 1, 0.1, 0.01, and 0.001 mol/1 NaCl solutions. The surfaces of cinnabar and
chalcocite are negatively charged at pH values greater than 3. Despite the anionic nature of the
surfaces of these 2 metal sulfides, I" sorbs strongly to both minerals. Measured Kd values far exceed
those reported for all other minerals with maximal Kd values of 1,375 and 3,080 ml/g, respectively,
for chalcocite and cinnabar between pH values of 4 and 5. Adsorption was substantial at all pH
values from 4 to 10. Balsley eta!. (1996) suggest that I" sorption apparently occurs by exchange of
hydroxyls attached to H and Cu sites.
Ticknor eta!. (1996) studied the effects of fulvic acid on the sorption of 125I (oxidation state not
specified) on crushed granite, biotite, goethite, and montmorillonite. Batch sorption experiments
with 28-day reaction times were conducted under aerobic conditions using in synthetic groundwater
solutions with pH values ranging from 7.6 to 7.8. Low sorption was measured for iodine on
montmorillonite in the presence of dissolved organic material. The Kd values for iodine sorbed to
montmorillonite ranged from 1.9 to 4.0 ml/g. No sorption (Kd values less than 1) was measured for
iodine on the other solids. Ticknor eta!. (1996) noted that if the sorbed iodine was associated with
the organic directly, then sorption of iodine on goethite should have been observed because the
goethite sorbed more fulvic acid than did the montmorillonite.
F.2
-------
Table F.1 Iodine adsorption studies on
pure minerals and oxide phases.
Mineral/Oxide
AgCl
A1(OH)3
A12O3
Allophane
Anhydrite
Apatite
Argentite
Attapulgite
Bauxite
Bentonite
Biotite
Bornite
Bournonite
Calcite
Cerusite
Chalcopyrite
Chalcocite
Chalcopyrite
Chlorite
Chysocolla
Cinnabar
References
Andersson et al. (1982), Allard et al. (1980)
Andersson et aL (1982), Allard et aL (1980)
Hakem et al. (1996, colloidal-size), Muramatsu et al. (1990), Ticknor and
Cho (1990, T and IO3)
Sazarashi et al. (1994,1995)
Strickert «-«/. (1980)
Strickert^/. (1980)
Kaplan et al. (2000a)
Sazarashi et d. (1994,1995), Rancon (1988), Allard et aL (1980)
Rancon (1988)
Muramatsu et al. (1990), Rancon (1988)
Ticknor et al. (1996)
Strickert^/. (1980)
Strickert^/. (1980)
Kaplan et al. (2000b), Kaplan et al. (1999), Ticknor and Cho (1990, I
and IO3)
Rancon (1988)
Sazarashi et al. (1995,1994), Rancon (1988), Andersson et al. (1982),
Strickert et al. (1980), Allard et al. (1980)
Kaplan et al. (2000a), Balsley et al. (1996,1997)
Kaplan et al. (2000a), Huie et al. (1988)
Kaplan et al. (2000b), Kaplan et al. (1999), Ticknor and Cho (1990, I
and IO3)
Rancon (1988), Strickert et al. (1980)
Kaplan et al. (2000a), Balsley et aL (1996, 1997), Sazarashi et aL
(1994,1995), Ikeda et al. (1994), Andersson et al. (1982), Allard et al.
(1980)
F.3
-------
Continuation of Table F.1.
Mineral/Oxide
Clinoptilolite
Coals (sub-bituminous)
Covellite
Cu (metal)
Cu2O
CuO
Dolomite
Enargite
Epidote
Fe2O3
Ferric hydroxide
Galena
Goethite
Gypsum
Halloysite
Hematite
Illite
Imogolite-bearing soil
Kaolinite
Laterite
References
Rancon (1988)
Balsleye^/. (1997)
Balsleye^/. (1997)
UnqefaL (1980)
ttzqetai (1980)
Haq^/. (1980)
Stnckert ef al. (1980)
Strickerte^/. (1980)
Ticknor and Cho (1990, I and IO3)
Muramatsu et al. (1990)
Andersson et al. (1982), Allard et al. (1980)
Kaplan et al. (2000a), Huie et al. (1988), Rancon (1988), Stnckert et al.
(1980), Allard et al. (1980)
Kaplan et al. (2000b), Kaplan et al. (1999), Ticknor et al. (1996), Ticknor
and Cho (1990, I and IO3), Rancon (1988)
Ticknor and Cho (1990, 1 and IO3)
Allard et al. (1980)
Ticknor and Cho (1990, 1 and IO3), Huie et al. (1988, rare-earth
hematite), Rancon (1988)
Kaplan et al. (2000b), Kaplan et al. (1999), Rancon (1988)
Balsleye^/. (1997)
Muramatsu et al. (1990), Ticknor and Cho (1990, 1 and IO3), Rancon
(1988)
Rancon (1988)
F.4
-------
Continuation of Table F.1.
Mineral/Oxide
Lignite
Limonite
Molybdenite
Montmorillonite
Muscovite
Olivine
PbOH
Pyrite
Quartz
Sepiolite
Serpentine
Siderite
SiO2 (colloidal size)
Stibnite
Tennantite
Tetrahedrite
Tiemannite
TiO2 (colloidal size)
Vermiculite
Witherite
Zinc hydrotalcite
Zeolite
References
Balsley^/. (1997)
Rancon (1988), Allard et aL (1980)
Huie etaL (1988)
Kaplan a aL (2000b), Kaplan et aL (1999), Ticknor et aL (1996), Sazarashi
etaL (1994,1995), Allard et aL (1980)
Ticknor and Cho (1990, I and IO3)
Andersson etaL (1982), Allard et al. (1980)
Andersson etaL (1982). Allard etaL (1980)
Huie etaL (1988), Stnckert etaL (1980)
Kaplan et al. (2000b), Kaplan et al. (1999), Ticknor et al. (1996),
Muramatsu et al. (1990), Ticknor and Cho (1990, I and IO3), Rancon
(1988), Allard etaL (1980)
Rancon (1988)
Andersson etaL (1982), Allard etaL (1980)
Rancon (1988)
Hakem^a/. (1996)
Kaplan etaL (2000a), Huie etaL (1988)
Stnckert et al. (1980)
Stnckert etaL (1980)
Zhuang^/. (1995)
Hakem^a/. (1996)
Kaplan etaL (2000b), Kaplan etaL (1999), Rancon (1988)
Stnckert et al. (1980)
Balsleyetal. (1997)
Rancon (1988)
F.5
-------
Ticknor and Cho (1990) studied the extent of sorption of I" and IO3 on bentonite and minerals
found in fractures in granite. The batch equilibration experiments were conducted using synthetic
groundwater (pH 7.7) and samples of crushed calcite, chlorite, epidote, goethite, gypsum, hematite,
kaolinite, muscovite, and quartz. Ticknor and Cho found no measurable sorption of I" or IO3 on
bentonite. No detectable sorption was measured for I" on any of the minerals selected for study.
Sorption of IO3 was observed with chlorite and hematite over the range of conditions used for their
study, and with kaolinite, epidote, goethite, and quartz at some of the experimental conditions.
There was no detectable sorption of IO3 on calcite, gypsum, and muscovite. Inorganic I" was not
sorbed by any of the minerals under any of the conditions used in their study.
Couture and Seitz (1983) measured the sorption of T, IO3, and IO4 (periodate) on hematite and
kaolinite, and the sorption of IO3 by pelagic red clay in buffer solutions and sea water. lodate was
strongly sorbed by hematite at pH values up to 9. Hematite is more effective than kaolinite as a
sorbent for IO3, and IO3 is more strongly sorbed than I" by kaolinite and hematite. Periodate (IO4)
was more strongly sorbed by hematite than IO3, and I" appears to be less strongly sorbed. At pH 4,
pure kaolinite (free of hematite and goethite) sorbs IO3 only slightly and I" not all. Couture and Seitz
(1983) determined that the sorption of IO3 was reversible upon change in pH. Because IO3 and IO4
were sorbed to a greater extent than I" by a number of minerals, especially iron oxides, the apparent
sorption of I" may be due to the sorption of oxidized iodine species. Based on their measurements,
Couture and Seitz (1983) suggest that results from published iodine sorption experiments should be
questioned unless the oxidation state of the iodine absorbate and the mineralogical purity of the
adsorbents were established.
F.2 Kd Studies of Iodine on Crushed Rock Materials
Numerous studies have been conducted of the sorption of iodine on crushed rock as part of
national research programs to investigate the feasibility of geological disposal of high-level
radioactive waste (HLW). Some of these studies are listed in Table F.2.
F.6
-------
Table F.2. Iodine adsorption studies on
crushed rock and related materials.
Geologic
Basalt
Basalt secondary
mineral assemblage
from vesicles, vugs,
and fractures
Chalk
Granite
Limestone
Shale
Siltstone
Tuff
Materials References
Strickert et al. (1980), Salter et al. (1981a), Ames and McGarrah (1980)
Salter et al. (1981b) [assemblage consists of approximately 98 percent
smectite clay and trace amounts of amorphous iron oxide, calcite, quartz,
opal (or a-cristobalite), and zeolites]
Strickert et aL (1980)
Strickert et aL (1980)
Strickert et aL (1980)
Strickert et aL (1980)
Strickert^/. (1980)
Strickert et aL (1980)
F.7
-------
APPENDIX G
Neptunium Adsorption Studies
-------
Appendix G
Neptunium Adsorption Studies
The Kd values listed in the following pure mineral/oxide and crushed rock studies are not
necessarily relevant to the mobility and sorption of neptunium in soils. However, the studies and
associated references are listed below for completeness. The potential value of these references and
the sorption processes that were investigated is left to the reader. The references cited in this
appendix are listed in the main reference list in Chapter 6.
G.1 Adsorption Studies of Neptunium on Single Mineral Phases
Several studies have been conducted for the adsorption of Np(V) on pure mineral and oxide phases.
Some of these studies are listed in Tables G.la and G.lb.
Table G.1 Neptunium adsorption studies on
pure mineral and oxide phases.
Mineral/Oxide
Albite
Anorthite
Apatite
Aragonite
Attapulgite
Bentonite
Biotite
Bohmite
Bytownite
Calcite
Chlorite
Clay
Corundum
Dolomite
Refer ences(s)
Torstenfelt et al. (1988), Beall et al. (1980)
Torstenfelt^/. (1988)
Andersson et al. (1982)
Keeney-Kennicutt et al. (1984)
Andersson et al. (1982)
Torstenfelt et al. (1988)
Nakayama and Sakamoto (1991), Andersson et al. (1982), Beall et al. (1980)
Kung and Triay (1994), Nakayama and Sakamoto (1991)
Beall etaL (1980)
Ticknor (1993). Keeney-Kennicutt et al. (1984), Andersson et al.
(1982)
Ticknor (1993)
Hart etal. (1994), Keeney-Kennicutt et al. (1984)
Nakayama and Sakamoto (1991)
Brady etal (1999)
G.2
-------
Continuation of Table G.1.
Mineral/Oxide
Epidote
Fluorite
Gibbsite
Goethite
Gypsum
Hematite
Hornblende
Hydrargillite
Illite
Iron oxyhydroxide
(amorphous)
Kaolinite
Lepidocrocite
Magnetite
Microcline
MnO2
Montmorillonite
Muscovite
Olivine
Quartz
SiO2 (amorphous)
Smectite
References(s)
Ticknor (1993)
Andersson et al. (1982)
Del Nero et al (1997)
Fujita etaL (1995), Tochiyama et al. (1995), Kung and Tnay (1994),
Ticknor (1993), Combes etaL (1992), Nakayama and Sakamoto (1991),
Keeney-Kennicutt et al. (1984)
Ticknor (1993)
Nakata et al. (2000), Kohler et al. (1999), Tochiyama et al. (1995), Ticknor
(1993), Nakayama and Sakamoto (1991), Allard (1984)
Torstenfelt^/. (1988)
Del Nero et al. (1997)
Nagasaki et al. (1998), Ticknor (1993), Torstenfelt et al. (1988)
Garnet al. (1991)
Ticknor (1993), Relyea and Martin (1982)
Nakayama and Sakamoto (1991)
Nakata et al. (2000), Fujita et al. (1995), Tochiyama et al. (1995), Nakayama
and Sakamoto (1991), Allard (1984), Andersson et al. (1982)
Torstenfelt et al. (1988), Beall et al. (1980)
Keeney-Kennicutt etaL (1984)
Nagaski and Tanaka (2000), Andersson et al. (1982)
Ticknor (1993)
Andersson et al. (1982)
Kohler et al. (1999), Ticknor (1993), Pratopo et al. (1991), Nakayama et al.
(1988), Andersson et al. (1982), Beall et al. (1980)
Kung and Triay (1994)
Koz^etal. (1996)
G.3
-------
Kung and Triay (1994) used the batch sorption technique to investigate the effect of naturally
occurring organic materials on the sorption of Np(V) on synthetic goethite, bohmite [also known as
boehmite, AIO(OH)], amorphous silica oxides, and crushed tuff1 material. The presence of the
model amino acid and fulvic acids did not affect the sorption of neptunium on tuff material or on
iron or aluminum oxides. The interaction between organic material and mineral surfaces in the
natural environment is important to mineral surface geochemistry. Sorption of organic material
onto mineral surfaces will affect not only the solubility and charge of organics in solution but also
mineral surface properties. These changes may affect the reactivity of the mineral to other metal
ions. Sorption of neptunium on iron oxides was 10 times greater than that on aluminum oxides.
Sorption of neptunium on crushed tuff material was much lower than that on the oxide surfaces.
Presence of model organic materials did not influence the sorption of neptunium on tuff or on iron
or aluminum oxides, suggesting that the dissolved model organics did not readily complex
neptunium or change the surface properties of the adsorbents.
Table G.2 Measured Np(V) Kd values (ml/1) for fracture-coating
minerals as a function of total dissolved solids (IDS) and normal
versus low oxygen conditions (Ticknor, 1993).
Mineral
Calcite
Chlorite
Epidote
Goethite
Gypsum
Hematite
Illite
Kaolinite
Muscovite
Quartz
Normal Oxygen Conditions
1,850 TDS
>49l
>57
2 2
>55
1.2
>69
>57
2.6
1.6
0
2,830 TDS
>48
>38
1.1
>39
1.4
>68
>39
1.4
0.5
0
12,580 TDS
6.5
12
0.6
16
0.6
>68
7.8
2.0
0.9
0
Low Oxygen Conditions
1,850 TDS
>66
>58
13
>62
5.1
>84
>79
4.1
6.3
1.3
2,830 TDS
>64
13
9.1
>47
5.1
>80
>76
2.1
5.3
0.8
12,580 TDS
15
9.5
3.6
20
3.0
>72
25
2.6
3.0
3.7
1 Kd values listed in italics are based on minimum analytical detection limit values for concentrations of dissolved
neptunium.
Ticknor (1993) completed batch equilibration experiments to study the sorption of 237Np in
three synthetic groundwater solutions on minerals present as fracture coatings in granitic rock. The
minerals included calcite, chlorite, epidote, goethite, gypsum, hematite, illite, kaolinite, muscovite,
Tuff is a general term for a rock formed of consolidated clastic (principally broken fragments) material that
originated from volcanic explosion or aerial expulsion from a volcanic vent (Bates and Jackson 1980).
G.4
-------
and quartz. Under the experimental conditions, Np(V) was the most probable oxidation state. The
pH of each solution was adjusted to 7.5 using NaOH. The three saline solutions differed in their
content of total dissolved solids (TDS). At the end of the contact period (27 to 30 days), the
measured pH under normal oxygen conditions was 8.0 to 8.5 for all mineral/solution combination
except for the kaolinite experiments which had a final pH between 6.5 to 7.0. The pH of the
mineral/solution experiments conducted under low oxygen conditions was 9.5. Ticknor (1993)
suggested the competition for sorption sites by increases in TDS may account for the decrease of
neptunium sorption on some mineral surfaces (Table G.2).
Combes et al. (1992) studied the sorption of dissolved Np(V) on goethite (crystalline a-FeOOH)
using synchrotron-based, extended X-ray absorption fine structure (EXAFS) and X-ray absorption
near-edge structure (XANES) spectroscopies. The sorption experiments consisted of reacting high-
surface area goethite with a pH 7.2, 0.08 M NaNO3 solution containing 1.3x10 5 M Np(V). Their
results indicate that Np(V) sorbs on goethite as a mononuclear, inner-sphere complex. The surface
complex may also contain hydroxide or carbonate ligands. There was no indication that the Np(V)
sorption on goethite occurred as an ordered neptunium oxide or hydroxide or as a coprecipitate with
iron oxide/hydroxide solid.
Nakayama and Sakamoto (1991) studied the sorption of neptunium on several minerals in 0.1 M
NaNO3 solution at pH values between 4 and 11. The minerals included samples of naturally
occurring hematite, magnetite, goethite, lepidocrocite, biotite, bohmite, and corundum. The
sorption of neptunium on goethite differed as a function of pH relative to that for hematite,
magnetite, and biotite. The experiments indicated strong sorption of neptunium on goethite at pH
values greater than 6, while sorption on hematite, magnetite, and biotite increased at pH values
above 9. The pH dependence of neptunium sorption on lepidocrocite and boehmite was similar to
that of goethite, and neptunium sorption on corundum was similar to that for hematite. Nakayama
and Sakamoto (1991) attributed the difference in neptunium sorption between the various minerals
to the pH * surface charge density, and the pH dependency of the different neptunium aqueous
species.
Although others have reported the enhanced adsorption of neptunium by Fe(II)-containing minerals
by surface-mediated heterogeneous reduction and sorption reactions, Nakayama and Sakamoto
(1991) observed no significant difference in the retardation by the Fe(II)-containing biotite and
magnetite versus that for Fe (III)-containing hematite.
Torstenfelt et al. (1988) used batch experiments to measure neptunium Kd values for several minerals
in groundwater as a function of pH. Natural groundwater containing a total concentrations of
10~7 M 237Np was equilibrated under aerobic conditions with crushed samples of granite from
three sites and six minerals which included hornblende, three feldspar minerals (albite, anorthite,
microcline), and two clays (illite and bentonite). The groundwater contained 143 mg/1 HCO3, and
its pH was adjusted to values between approximately 2 and 11 using NaOH or HC1 solutions.
Torstenfelt et al. (1988) observed greater sorption of Np(V) at higher pH than at lower pH values.
As the pH increases, mineral surfaces become increasingly more negatively charged. The pH where the surface is
uncharged (i.e., zero net charge) is referred to as the pH of zero-point-of-charge, pH,,pc.
G.5
-------
The results plotted1 by Torstenfelt et al. (1988) typically indicated that the Kd values for Np(V) were
near 0 ml/g measured at pH values near 2 for all solids, generally ranged from 50 to 100 ml/g for
the granite and clay samples at pH values near 7 to 8, and varied between 200 and 300 ml/g for
most of the mineral specimens at high pH values. At pH values from approximately 2 to 4, the Kd
values measured for the feldspar minerals were typically in the region of 10 to 30 ml/g.
For the experiments conducted with hornblende, the Kd values ranged from 200 to 250 ml/g at
pH 8.1 and 290 to 350 ml/g at pH 11.2. Because hornblende contains Fe(II), Torstenfelt et al.
(1988) speculated that these high Kd values might be affected by surface-mediated heterogeneous
reduction and sorption reactions on this Fe(II)-containing silicate mineral. For example,
autoradiography conducted by Beall et al. (1980) indicate that the sorption of Np-237(V) on a fresh
surface of granite in aerated groundwater shows a strong correlation to the Fe(II)-containing
minerals (pyrite and biotite) in the granite. Studies by Susak et al. (1983) and Meyer et al. (1984)
indicate that Fe(II)-containing geologic materials are able to reduce the concentrations of dissolved
neptunium and increase the concentrations of dissolved and sorbed Np(IV). Susak et al. (1983)
investigated the reduction of Np(VI) in the presence of olivine [(Fe(II),Mg)2SiO3]. Susak et al.
proposed the following sequence of reactions for the reduction and subsequent retardation of
neptunium:
NpOf + 4 FT - NpVI(adsorbed) + 2 H2O
NpVI (adsorbed) + [Fe(II)2SiO3] - Npv (adsorbed) + [Fe(II,III)2SiO3]
Npv (adsorbed) + 2 H2O - NpO+ + 4 FT.
Nakayama et al. (1988) used high performance liquid chromatography to study the migration of
Np(V) in columns packed with quartz. The distribution coefficient derived for Np(V) from these
column experiments decreased with increasing ionic strength and increased with increasing pH.
Keeney-Kennicutt et al. (1984) studied the sorption of Np(V) on several minerals in dilute aqueous
solutions and seawater. When normalized to the amount of adsorption per unit solid surface area,
their results indicated that Np(V) sorbs on the following minerals with decreasing affinity in the
sequence
aragonite > calcite > goethite » MnO2 ~ clays
Their measured dissolved/adsorbed ratios were constant over the range of 10~13 to 10~7 M Np(V).
Contrary to speculation at that time, their results indicated that Np(V) did not behave as a simple
monovalent ion with a low affinity for surfaces.
Relyea and Martin (1982) studied the adsorption of Np-237 on kaolinite over a range of
concentrations, pH, electrolyte strengths, and time. The adsorption of 237Np was determined to be
strongly affected by pH. Adsorption of Np-237 on kaolinite was weak at pH values less than pH 5,
but increased rapidly from pH 5 to 7.5. At pH values greater than 7.5, the adsorption of Np-237
was approximately 100 percent.
The coarse scale used for Kd axes by Torstenfelt et al. (1988) precluded interpolation of accurate values for the
experimentally-determined Kd values for each solid type.
G.6
-------
G.2 Kd Studies of Neptunium on Crushed Rock Materials
Numerous studies have been conducted for the sorption of neptunium on crushed rock as part of
national research programs to investigate the feasibility of geological disposal of high-level
radioactive waste (HLW). Some of these studies are listed in Table G.2.
Table G.3 Neptunium adsorption studies on
crushed rock and related materials.
Geologic Material
Basalt
Basalt secondary mineral
assemblage from vesicles,
vugs, and fractures
Diabase
Gneiss
Granite
Sandstone
Tuff
References
Meyer et al., (1984, 1983, 1985a, 1985b), Susak et al. (1983), Barney
etal. (1983), Salter et al. (1981a), Ames et al. (1981), Ames and
McGarrah (1980)
Salter etal. (1981b)
Andersson etal. (1982)
Andersson^/. (1982)
Torstenfelt etal. (1988), Andersson etal. (1982), Beall etal. (1980),
Allard (1979)
Barney (1982)
Baston etal. (1997), Rung and Triay (1994), Barney (1982)
G.7
-------
APPENDIX H
Radium Adsorption Studies
-------
Appendix H
Radium Adsorption Studies
The Kd values listed in the following pure mineral/oxide and crushed rock studies are not
necessarily relevant to the mobility and sorption of radium in soils. However, the studies and
associated references are listed below for completeness. The potential value of these references and
the sorption processes that were investigated is left to the reader. The references cited in this
appendix are listed in the main reference list in Chapter 6.
H.1 Adsorption Studies of Radium on Single Mineral Phases
Several studies have been conducted for the adsorption of radium on pure mineral and oxide phases.
These studies and the references they cite may be useful to the reader. Some of these studies are
listed in Table H.I.
The results of these studies demonstrate that the adsorption of radium on minerals and oxide phases
is dependent on pH, as expected for cations, and decreases with increasing concentrations of
competing ions. For example, Tachi eta!. (2001) used batch sorption experiments to study the
sorption of radium on bentonite and purified smectite as a function of pH, ionic strength, and
liquid-to-solid ratio. The measured Kd values ranged from 102 to more than 104 ml/g, and were
dependent on ionic strength and pH. Tachi eta!. (2001) found that most of the sorbed radium was
desorbed by a 1 mol/1 KC1 solution. Modeling results indicated that radium sorption on the
smectite was dominated by ion exchange at layer sites of smectite and may increase at the higher pH
values from surface complexation at edge sites. Radium sorption on bentonite was explained as
being due to ion exchange with smectite, and its pH dependency resulted from changes in calcium
concentrations caused from dissolution and precipitation of calcite.
Koulouris (1995) used batch and column techniques to investigate the sorption of Ra-226 on MnO2.
Radium sorption was studied as a function of initial Ra-226 concentrations, pH, chemical species,
and contact time. Koulouris (1995) observed no dependency in Ra-226 sorption on MnO2 as a
function of pH at values greater than 3.6. Radium sorption was essentially constant at
approximately 50 to 60 percent over the pH range 3.6 to 12.0, and decreased from 50 at pH 3.6
to 18 percent at pH 2.7. Koulouris (1995) attributed the decrease in Ra-226 sorption to a low pH2pc
for the MnO2 material used in the experiments. Essentially no dependency was observed in Ra-226
sorption on MnO2 as a function of initial Ra-226 concentrations and chemical species.
H.2
-------
Table H.1 Radium adsorption studies on pure mineral and oxide phases.
Mineral/Oxide
Amorphous ferric
oxyhydroxide
Anhydrite
Bentonite
Biotite
Calcite
Clinoptilolite
Dolomite
Feldspar
Ferric hydroxide
Glauconite
Illite
Kaolinite
MnO2
Montmorillonite
Muscovite
Nontronite
Opal
Phlogopite
Quartz
Silica
Smectite
References
Ames et al.
Maiti et al.
Tachi et al.
(1983c)
(1989)
(2001), Hanslik and Mansfeld (1990)
Hanslik and Mansfeld (1990), Ames et al. (1983a)
Maiti et al.
Ames et al.
Maiti et al.
(1989)
(1983b)
(1989)
Hanslik and Mansfeld (1990), Benes et al. (1986)
Benes et al
Ames et al.
Ames et al.
(1984),
(1983b)
(1983b)
Hanslik and Mansfeld (1990), Benes et al. (1985), Ames et al. (1983b),
Relyea and Martin (1982), Riese (1982)
Koulouris
Benes et al
Benes et al
Ames et al.
Ames et al.
(1995)
(1 985), Ames et al. (1983b)
(1 986), Ames et ai (198ty
(1983b)
(1983b)
Hanslik and Mansfeld (1990)
Benes et al
(1984), Riese (1982)
Hanslik and Mansfeld (1990)
Tachi et al.
(2001)
Benes and his co-investigators have studied the sorption of radium on sediments (Benes and Strejc,
1986), feme hydroxide (Benes et al, 1984), feldspar (Benes et al, 1986), kaolimte (Benes et al, 1985),
montmorillonite (Benes et al, 1985), muscovite (Benes et al, 1986), and quartz (Benes et al, 1984).
Benes and Strejc (1986) studied the adsorption and desorption of radium on stream sediments.
They determined that the radium adsorption on the sediment could not easily be predicted from the
composition and other properties of the sediments. No simple correlations could be made
H.3
-------
regarding the extent of radium sorption and the specific surface area, organic matter, mineral oxide
coatings, or other sediment components. As previously noted, Benes and Strejc (1986) observed
however that in the presence of dissolved sulfate, a strong dependency existed between the uptake
of radium by the sediment and the barite content of the sediments. Benes and his co-investigators
also indicated that the adsorption of radium was also strongly dependent on pH (Benes et al., 1984).
For example, the maximum adsorption of radium on ferric hydroxide occurred at pH values of 7
and greater, and radium adsorption decreased to almost negligible values at pH values less than 6.
The pH dependency of radium adsorption on kaolinite was less than that determined for ferric
hydroxide and quartz, and radium sorbed to a significant extent on kaolinite at pH values less than 6.
Riese (1982) studied the adsorption of radium on quartz and kaolinite. His measurements indicated
that radium was 50 percent adsorbed on quartz at pH 6 and 50 percent adsorbed on kaolinite at
pH 5. The adsorption of radium on quartz and kaolinite was also found to decrease with increasing
concentrations of dissolved calcium. Results from Riese (1982), as shown on p. 386 of Langmuir
(1997), for the adsorption of radium on kaolinite at 25°C in 0.01 M NaNO3 solution containing
10~4 M calcium indicate that significant adsorption of radium occurs on kaolinite over the pH
range 3 to 9. The adsorption of radium increases with increasing pH, especially at pH values greater
than 6.
Lowson and Evans (1984) measured the adsorption of radium in a 0.1 M NaNO3 solution on
alumina, kaolinite, montmorillonite, and illite. The adsorption of radium on alumina and kaolinite
was low at pH values less than 7, but increased rapidly at pH values greater than 8. Montmorillonite
and illite were found to be stronger adsorbents for radium than alumina and kaolinite.
Montmorillonite and illite exhibited significant adsorption (about 20 percent) for radium at pH 3,
and the adsorption of radium on these two minerals increased approximately at a linear rate with
increasing pH at higher pH values.
H.2 Kd Studies of Radium on Crushed Rock Materials
Numerous studies have been conducted for the sorption of radium on crushed rock as part of
national research programs to investigate the feasibility of geological disposal of high-level
radioactive waste (HLW). Some of these studies are listed in Table H.2. Because elevated
temperatures will exist in the near vicinity of the waste canisters and repository in general, many of
these studies were conducted over a range temperatures above ambient to determine the extent of
radionuclide adsorption as a function of temperature and with respect to changes in groundwater
chemistry (e.g., pH and Eh) and mineralogy that will occur with increasing temperature.
H.4
-------
Table H.2 Radium adsorption studies on
crushed rock and related materials.
Geologic Material
Basalt
Basalt secondary
mineral assemblage
from vesicles, vugs, and
fractures
Sandstone
Tuff
References
Puls etal. (1987), Salter et al. (1981a), Ames et al. (1981), Ames and
McGarrah (1980)
Salter etal (1981b)
Barney etal. (1982)
Barney^/. (1982)
H.5
-------
APPENDIX I
Technetium Adsorption Studies
-------
Appendix I
Technetium Adsorption Studies
The Kd values listed in the following pure mineral/oxide and crushed rock studies are not
necessarily relevant to the mobility and sorption of technetium in soils. However, the studies and
associated references are listed below for completeness. The potential value of these references and
the sorption processes that were investigated is left to the reader. The references cited in this
appendix are listed in the main reference list in Chapter 6.
1.1 Adsorption Studies of Technetium on Single Mineral Phases
Numerous technetium adsorption studies have been conducted of technetium on pure mineral,
oxide phases, and other geologic-related materials. Like those for iodine, many of these studies were
conducted because of extensive research interest in developing getters (adsorbents) that could be
added to waste streams and tailored barriers for removal and/or immobilization of dissolved iodine
[e.g., see Zhang et al. (2000), Balsley et al. (1997), Gu and Dowlen (1997), and others]. High sorption
of technetium is typically observed for minerals, such as sulfide minerals (e.g., chalcopyrite, pyrite),
that have the capacity to reduce Tc(VII) to Tc(IV). For example, Strickert et al. (1980) measured Kd
values in the range 100 to 2,000 ml/g for sorption of technetium on sulfide minerals such as
bournonite, chalcopyrite, pyrite, tennantite, and tetrahedrite. Technetium Kd values were however
less than 1 ml/g for non-sulfide materials, such as anhydrite, basalt, granite, and tuff. Sorption
studies for technetium on single mineral phases and the references they cite may be useful to the
reader. Some of these studies are listed in Tables I.la and Lib.
1.2 Kd Studies of Technetium on Crushed Rock Materials
Numerous studies have been conducted for the sorption of technetium on crushed rock as part of
national research programs to investigate the feasibility of geological disposal of high-level
radioactive waste (HLW). Some of these studies are listed in Table 1.2.
1.2
-------
Table 1.1 Technetium adsorption studies on
pure mineral and oxide phases.
Mineral/Oxide
Almandine
Anhydrite
Apatite
Augite
Bentonite
Biotite
Bornite
Bournonite
Chalcopyrite
Chysocolla
Dolomite
Enargite
Fe(II) Silicates
Ferrihydrite
Galena
Goethite
Hematite
Hornblende
Hypers thene
Ilmenite
Limonite
References
Bocketai (1989)
Stnckert et al. (1980)
Strickert^/. (1980)
Palmer and Meyer (1981)
Boston etai (1995)
Ticknor et al. (1996), Amaya et al. (1995), Meyer et al. (1983), Palmer
and Meyer (1981)
Stnckert et al. (1980)
Strickert et aL (1980)
Huie etal. (1988), Stnckert et al. (1980)
Stnckert et al. (1980)
Stnckert et aL (1980)
Stnckert et aL (1980)
Bocketai (1989)
Walton et al. (1986)
Bock etal. (1989), Huie etal. (1988), Meyer etal.
Meyer (1981), Strickert etal. (1980)
Ticknor et aL (1996), Walton et al. (1986), Meyer
Huie et al. (1988), Hames et al. (1987), Walton ei
al. (1984, 1985a). Palmer and Meyer (1981)
(1983), Palmer and
etal. (1984)
' aL (1986), Meyer et
Palmer and Meyer (1981)
Bocketai (1989)
Bock etai (1989), Meyer et aL (1984, 1985a), Palmer and Meyer (1981)
Meyer et aL (1984), Palmer and Meyer (1981)
1.3
-------
Continuation of Table 1.1.
Mineral/Oxide
Lolling! te (Loellingite)
Magnetite
Microcline
Molybdenite
Montmorillonite
Mordenite
Olivine
Plagioclase
Potassium Feldspar
Pyrite
Pyrrhotite
Quartz
Smectite
Stibnite
Tennantite
Tetrahedrite
Tiemannite
Witherite
References
Rocket al. (1989)
Byegard et al. (1992), Bock et al (1989), Lieser and Bauscher (1988),
Hames et al. (1987), Meyer et aL (1984), Palmer and Meyer (1981)
Meyer et al. (1984, 1985a)
Hme et al. (1988)
Ticknor et al. (1996)
Meyer et al. (1985a)
Bock etal (1989), Palmer and Meyer (1981)
Amzyz et al. (1995)
Amzyzetal. (1995)
Bock etal (1989), Lieser and Bauscher (1988), Hme etal. (1988),
Meyer et aL (1983), Palmer and Meyer (1981), Strickert et al (1980)
Bock etal. (1989), Lieser and
Ticknor et al. (1996), Amaya
Bauscher (1988)
etaL (1995), Wmkler et al. (1988)
Wmkler et al. (1988)
Zhuang etal. (1995), Bock et
al. (1989), Huie tf al. (1988)
Stnckert et al. (1980)
Stnckert et al. (1980)
Zhuang etal. (1995)
Stnckert et al. (1980)
1.4
-------
Table 1.2 Technetium adsorption studies on
crushed rock and related materials.
Geologic Material
Alluvium
Argillite
Basalt
Basalt secondary
mineral assemblage
from vesicles, vugs, and
fractures
Chalk
Gabbro
Granite
Granodiorite
Limestone
Sandstone
Shale
Siltstone
Tuff
References
EtddeiaL (1979)
Erdal^/. (1979)
Wood etaL (1987), Meyer et aL (1984, 1985), Barney et aL (1983),
Palmer and Meyer (1981), Salter etaL (1981a), Ames etaL (1981),
Strickert etaL (1980), Ames and McGarrah (1980)
Salter etaL (1981b)
Strickert ^«/. (1980)
Vandergraaf et al. (1984)
Ticknor etaL (1996), Amaya etaL (1995), Enksen and Cm (1991),
Vandergraaf et al. (1984), Strickert etaL (1980), Allard etaL (1979a,
1979b)
Baston^/. (1995)
Strickert ^«/. (1980)
Bradbury and Stephen (1985), Barney (1982)
Ho (1992), Strickert etaL (1980)
Strickert e^/. (1980)
Amaya et al. (1995), Baston etaL (1995), Barney (1982), Strickert et al.
(1980), Erdal etaL (1979)
1.5
------- |