United States Office of Air and Radiation EPA 402-R-99-004B
Environmental Protection August 1999
Agency
UNDERSTANDING VARIATION IN
PARTITION COEFFICIENT, Kd, VALUES
Volume II:
Review of Geochemistry and Available Kd Values
for Cadmium, Cesium, Chromium, Lead, Plutonium,
Radon, Strontium, Thorium, Tritium (3H), and Uranium
Case I: Kd = imi/g
Continuous Source of Contamination Steady State
- C/C, = 0.9 Flow
C/CJj =0.1
uasell: Kd =10 ml/g
Continuous Source of Contamination
Steady State
Flow
c/q, =0.1
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UNDERSTANDING VARIATION IN
PARTITION COEFFICIENT, Kd, VALUES
Volume II:
Review of Geochemistry and Available Kd Values
for Cadmium, Cesium, Chromium, Lead, Plutonium,
Radon, Strontium, Thorium, Tritium (3H), and Uranium
August 1999
A Cooperative Effort By:
Office of Radiation and Indoor Air
Office of Solid Waste and Emergency Response
U.S. Environmental Protection Agency
Washington, DC 20460
Office of Environmental Restoration
U.S. Department of Energy
Washington, DC 20585
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NOTICE
The following two-volume 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 more 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 two-volume report 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.
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 this two-volume document include: chromium, cadmium,
cesium, lead, plutonium, radon, strontium, thorium, tritium (3H), and uranium. This two-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 absolute
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,
Volume II of this document: (1) provides a "thumb-nail sketch" of the key geochemical
processes affecting the sorption of the selected contaminants; (2) provides references to related
key experimental and review articles for further reading; (3) identifies the important aqueous-
and solid-phase parameters controlling the sorption of these contaminants in the subsurface
environment under oxidizing conditions; and (4) identifies, 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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This publication is the result of a cooperative effort between the EPA Office of Radiation
and Indoor Air, Office of Solid Waste and Emergency Response, and the Department of Energy
Office of Environmental Restoration (EM-40). 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.
Stephen D. Page, Director
Office of Radiation and Indoor Air
iv
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ACKNOWLEDGMENTS
Ronald G. Wilhelm from ORIA's Center for Remediation Technology and Tools was the
project lead and EPA Project Officer for this two-volume report. Paul Beam, Environmental
Restoration Program (EM-40), was the project lead and sponsor for the Department of Energy
(DOE). Project support was provided by both DOE/EM-40 and EPA's Office of Remedial and
Emergency Response (OERR).
EPA/ORIA wishes to thank the following people for their assistance and technical review
comments on various drafts of this report:
Patrick V. Brady, U.S. DOE, Sandia National Laboratories
David S. Brown, U.S. EPA, National Exposure Research Laboratory
Joe Eidelberg, U.S. EPA, Region 9
Amy Gamerdinger, Washington State University
Richard Graham, U.S. EPA, Region 8
John Griggs, U.S. EPA, National Air and Radiation Environmental Laboratory
David M. Kargbo, U.S. EPA, Region 3
Ralph Ludwig, U.S. EPA, National Risk Management Research Laboratory
Irma McKnight, U.S. EPA, Office of Radiation and Indoor Air
William N. O'Steen, U.S. EPA, Region 4
David J. Reisman, U.S. EPA, National Risk Management Research Laboratory
Kyle Rogers, U.S. EPA, Region 5
Joe R. Williams, U.S. EPA, National Risk Management Research Laboratory
OSWER Regional Groundwater Forum Members
In addition, special acknowledgment goes to Carey A. Johnston from ORIA's Center for
Remediation Technology and Tools for his contributions in the development, production, and
review of this document.
Principal authorship in production of this guide was provided by the Department of Energy's
Pacific Northwest National Laboratory (PNNL) under the Interagency Agreement Number
DW8993 7220-01-03. Lynnette Downing served as the Department of Energy's Project Officer
for this Interagency Agreement. PNNL authors involved in this project include:
Kenneth M. Krupka
Daniel I. Kaplan
Gene Whelan
R. Jeffrey Serne
Shas V. Mattigod
v
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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 (Kd) Values
401 M Street, SW (6602J)
Washington, DC 20460
or
radiation.questions@epa.gov
VI
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ABSTRACT
This two-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. Volume I
discusses the technical issues associated with the measurement of Kd values and its use in
formulating the retardation factor, Rf. The Kd concept and methods for measurement of Kd
values are discussed in detail in Volume I. 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 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 presented. The capabilities of EPA's geochemical reaction model
MINTEQA2 and its different conceptual adsorption models are also reviewed. Volume II
provides a "thumb-nail sketch" of the key geochemical processes affecting the sorption of
selected inorganic contaminants, and a summary of Kd values given in the literature for these
contaminants under oxidizing conditions. The contaminants chosen for the first phase of this
project include chromium, cadmium, cesium, lead, plutonium, radon, strontium, thorium, tritium
(3H), and uranium. Important aqueous speciation, (co)precipitation/dissolution, and adsorption
reactions are discussed for each contaminant. References to related key experimental and review
articles for further reading are also listed.
Vll
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CONTENTS
Page
NOTICE ii
FOREWORD iii
ACKNOWLEDGMENTS v
FUTURE UPDATES vi
ABSTRACT vii
LIST 01 FIGURES xiii
LIST 01 TABLES xv
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 3.1
3.3 Other Methods 3.2
3.4 Issues 3.2
4.0 Application of Chemical Reaction Models 4.1
5.0 Contaminant Geochemistry and Kd Values 5.1
5.1 General 5.1
5.2 Cadmium Geochemistry and Kd Values 5.5
5.2.1 Overview: Important Aqueous- and Solid-Phase Parameters
Controlling Retardation 5.5
5.2.2 General Geochemistry 5.5
5.2.3 Aqueous Speciation 5.6
5.2.4 Dissolution/Precipitation/Coprecipitation 5.8
5.2.5 Sorption/Desorption 5.9
5.2.6 Partition Coefficient, Kd, Values 5.10
5.2.6.1 General Availability of Kd Values 5.10
5.2.6.2 Look-Up Tables 5.11
5.2.6.2.1 Limits of Kd Values with Aluminum/Iron-Oxide Concentrations 5.11
5.2.6.2.2 Limits of Kd Values with Respect to CEC 5.12
5.2.6.2.3 Limits of Kd Values with Respect to Clay Concentrations 5.12
5.2.6.2.4 Limits of Kd Values with Respect to Concentration of
Organic Matter 5.12
5.2.6.2.5 Limits of Kd Values with Respect to Dissolved Calcium,
viii
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Magnesium, and Sulfide Concentrations, and Redox Conditions
5.12
5.3 Cesium Geochemistry and Kd Values 5.13
5.3.1 Overview: Important Aqueous- and Solid-Phase Parameters
Controlling Retardation 5.13
5.3.2 General Geochemistry 5.13
5.3.3 Aqueous Speciation 5.13
5.3.4 Dissolution/Precipitation/Coprecipitation 5.14
5.3.5 Sorption/Desorption 5.14
5.3.6 Partition Coefficient, Kd, Values 5.15
5.3.6.1 General Availability of Kd Data 5.15
5.3.6.2 Look-Up Tables 5.16
5.3.6.2.1 Limits of Kd with Respect to pH 5.18
5.3.6.2.2 Limits of Kd with Respect to Potassium, Ammonium,
and Aluminum/Iron-Oxide Concentrations 5.18
5.4 Chromium Geochemistry and Kd Values 5.18
5.4.1 Overview: Important Aqueous- and Solid-Phase Parameters
Controlling Retardation 5.18
5.4.2 General Geochemistry 5.18
5.4.3 Aqueous Speciation 5.19
5.4.4 Dissolution/Precipitation/Coprecipitation 5.19
5.4.5 Sorption/Desorption 5.20
5.4.6 Partition Coefficient, Kd, Values 5.21
5.4.6.1 General Availability of Kd Data 5.21
5.4.6.2 Look-Up Tables 5.22
5.4.6.2.1 Limits of Kd with Respect to pH 5.23
5.4.6.2.2 Limits of Kd with Respect to Extractable Iron Content 5.23
5.4.6.2.3 Limits of Kd with Respect to Competing Anion Concentrations .. 5.23
5.5 Lead Geochemistry and Kd Values 5.25
5.5.1 Overview: Important Aqueous- and Solid-Phase Parameters
Controlling Retardation 5.25
5.5.2 General Geochemistry 5.25
5.5.3 Aqueous Speciation 5.26
5.5.4 Dissolution/Precipitation/Coprecipitation 5.27
5.5.5 Sorption/Desorption 5.30
5.5.6 Partition Coefficient, Kd, Values 5.31
5.5.6.1 General Availability of Kd Data 5.31
5.5.6.2 Kd Look-Up Tables 5.33
5.5.6.2.1 Limits of Kd with Respect to pH 5.33
5.5.6.2.2 Limits of Kd with Respect to Equilibrium Lead
Concentrations Extractable Iron Content 5.34
ix
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5.6 Plutonium Geochemistry and Kd Values 5.34
5.6.1 Overview: Important Aqueous- and Solid-Phase Parameters
Controlling Retardation 5.34
5.6.2 General Geochemistry 5.34
5.6.3 Aqueous Speciation 5.35
5.6.4 Dissolution/Precipitation/Coprecipitation 5.37
5.6.5 Sorption/Desorption 5.40
5.6.6 Partition Coefficient, Kd, Values 5.41
5.6.6.1 General Availability of Kd Data 5.41
5.6.6.2 Kd Look-Up Tables 5.43
5.6.6.2.1 Limits of Kd with Respect to Clay Content 5.43
5.6.6.2.2 Limits of Kd with Respect to Dissolved Carbonate
Concentrations 5.44
5.7 Radon Geochemistry and Kd Values 5.44
5.7.1 Overview: Important Aqueous- and Solid-Phase Parameters
Controlling Retardation 5.44
5.7.2 General Geochemistry 5.45
5.7.3 Aqueous Speciation 5.45
5.7.4 Dissolution/Precipitation/Coprecipitation 5.46
5.7.5 Sorption/Desorption 5.46
5.7.6 Partition Coefficient, Kd, Values 5.46
5.8 Strontium Geochemistry and Kd Values 5.46
5.8.1 Overview: Important Aqueous- and Solid-Phase Parameters
Controlling Retardation 5.46
5.8.2 General Geochemistry 5.47
5.8.3 Aqueous Speciation 5.47
5.8.4 Dissolution/Precipitation/Coprecipitation 5.48
5.8.5 Sorption/Desorption 5.49
5.8.6 Partition Coefficient, Kd, Values 5.51
5.8.6.1 General Availability of Kd Data 5.51
5.8.6.2 Look-Up Tables 5.51
5.8.6.2.1 Limits of Kd with Respect to pH, CEC, and
Clay Concentrations Values 5.52
5.8.6.2.2 Limits of Kd with Respect to Dissolved Calcium
Concentrations 5.52
5.8.6.2.3 Limits of Kd with Respect to Dissolved Stable
Strontium and Carbonate Concentrations 5.53
5.9 Thorium Geochemistry and Kd Values 5.53
5.9.1 Overview: Important Aqueous- and Solid-Phase Parameters
Controlling Retardation 5.53
5.9.2 General Geochemistry 5.54
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5.9.3 Aqueous Speciation 5.55
5.9.4 Dissolution/Precipitation/Coprecipitation 5.58
5.9.5 Sorption/Desorption 5.60
5.9.6 Partition Coefficient, Kd, Values 5.61
5.9.6.1 General Availability of Kd Data 5.61
5.9.6.2 Look-Up Tables 5.62
5.9.6.2.1 Limits of Kd with Respect to Organic Matter and
Aluminum/Iron-Oxide Concentrations 5.63
5.9.6.2.2 Limits of Kd with Respect to Dissolved Carbonate
Concentrations 5.63
5.10 Tritium Geochemistry and Kd Values 5.64
5.10.1 Overview: Important Aqueous- and Solid-Phase Parameters
Controlling Retardation 5.64
5.10.2 General Geochemistry 5.64
5.10.3 Aqueous Speciation 5.65
5.10.4 Dissolution/Precipitation/Coprecipitation 5.65
5.10.5 Sorption/Desorption 5.65
5.10.6 Partition Coefficient, Kd, Values 5.65
5.11 Uranium Geochemistry and Kd Values 5.65
5.11.1 Overview: Important Aqueous- and Solid-Phase Parameters
Controlling Retardation 5.65
5.11.2 General Geochemistry 5.66
5.11.3 Aqueous Speciation 5.67
5.11.4 Dissolution/Precipitation/Coprecipitation 5.69
5.11.5 Sorption/Desorption 5.72
5.11.6 Partition Coefficient, Kd, Values 5.74
5.11.6.1 General Availability of Kd Data 5.74
5.11.6.2 Look-Up Table 5.74
5.11.6.2.1 Limits Kd Values with Respect to Dissolved
Carbonate Concentrations 5.75
5.11.6.2.2 Limits of Kd Values with Respect to Clay Content and CEC ... 5.76
5.11.6.2.3 Use of Surface Complexation Models to Predict
Uranium Kd Values 5.76
5.12 Conclusions 5.77
6.0 References 6.1
Appendix A - Acronyms and Abbreviations A. 1
Appendix B - Definitions B.l
XI
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Appendix C - Partition Coefficients for Cadmium C.l
Appendix D - Partition Coefficients for Cesium D.l
Appendix E - Partition Coefficients for Chromium E. 1
Appendix F - Partition Coefficients for Lead F.l
Appendix G - Partition Coefficients for Plutonium G. 1
Appendix H - Partition Coefficients for Strontium H. 1
Appendix I - Partition Coefficients for Thorium 1.1
Appendix J - Partition Coefficients for Uranium J. 1
xii
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LIST OF FIGURES
Page
Figure 5.1. Calculated distribution of cadmium aqueous species as a function of pH
for the water composition in Table 5.1 5.7
Figure 5.2. Calculated distribution of lead aqueous species as a function of
pH for the water composition listed in Table 5.1 5.29
Figure 5.3. Calculated distribution of plutonium aqueous species as a function of
pH for the water composition in Table 5.1 5.39
Figure 5.4. Calculated distribution of thorium hydrolytic species as a function of pH. . . 5.57
Figure 5.5. Calculated distribution of thorium aqueous species as a function of
pH for the water composition in Table 5.1 5.59
Figure 5.6a. Calculated distribution of U(VI) hydrolytic species as a function of
pH at 0.1 |_ig/l total dissolved U(VI) 5.70
Figure 5.6b. Calculated distribution of U(VI) hydrolytic species as a function of pH
at 1,000 |_ig/l total dissolved U(VI) 5.71
Figure 5.7. Calculated distribution of U(VI) aqueous species as a function of pH
for the water composition in Table 5.1 5.72
Figure C. 1. Relation between cadmium Kd values and pH in soils C.5
Figure D. 1. Relation between cesium Kd values and CEC D.7
Figure D.2. Relation between CEC and clay content D.8
Figure D.3. Kd values calculated from an overall literature Fruendlich equation for
cesium (Equation D.2) D.12
Figure D.4. Generalized cesium Freundlich equation (Equation D.3) derived
from the literature D.16
Figure D.5. Cesium Kd values calculated from generalized Fruendlich equation
(Equations D.3 and D.4) derived from the literature D. 16
Xlll
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Figure E. 1. Variation of Kd for Cr(VI) as a function of pH and DCB extractable
Iron content without the presence of competing anions E.10
Figure F. 1. Correlative relationship between Kd and pH F.6
Figure F.2. Variation of Kd as a function of pH and the equilibrium
lead concentrations F.7
Figure G. 1. Scatter plot matrix of soil properties and the partition
coefficient (Kd) of plutonium G.12
Figure G.2. Variation of Kd for plutonium as a function of clay content and
dissolved carbonate concentrations G.14
Figure H. 1. Relation between strontium Kd values and CEC in soils H.5
Figure H.2. Relation between strontium Kd values for soils with CEC
values less than 15 meq/100 g H.7
Figure H.3. Relation between strontium Kd values and soil clay content H.7
Figure H.4. Relation between strontium Kd values and soil pH H.9
Figure 1.1. Linear regression between thorium Kd values and pH for the pH
range from 4 to 8 1.5
Figure 1.2. Linear regression between thorium Kd values and pH for the pH
range from 4 to 8 1.8
Figure J. 1. Field-derived Kd values for 238U and 235U from Serkiz and Johnson (1994)
plotted as a function of porewater pH for contaminated
soil/porewater samples J.8
Figure J.2. Field-derived Kd values for 238U and 235U from Serkiz and Johnson (1994)
plotted as a function of the weight percent of clay-size particles in the
contaminated soil/porewater samples J.9
Figure J.3. Field-derived Kd values for 238U and 235U plotted from Serkiz and Johnson (1994)
as a function of CEC (meq/kg) of the contaminated
soil/porewater samples J. 10
Figure J.4. Uranium Kd values used for development of Kd look-up table J. 19
xiv
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LIST OF TABLES
Page
Table 5.1. Estimated mean composition of river water of the world from Hem (1985) 5.3
Table 5.2. Concentrations of contaminants used in the aqueous species
distribution calculations 5.4
Table 5.3. Cadmium aqueous species included in the speciation calculations 5.6
Table 5.4. Estimated range of Kd values for cadmium as a function of pH 5.11
Table 5.5. Estimated range of Kd values (ml/g) for cesium based on CEC
or clay content for systems containing <5 percent mica-like minerals
in clay-size fraction and <10"9 M aqueous cesium 5.17
Table 5.6. Estimated range of Kd values (ml/g) for cesium based on CEC
or clay content for systems containing >5 percent mica-like minerals
in clay-size fraction and <10"9 M aqueous cesium 5.17
Table 5.7. Estimated range of Kd values for chromium (VI) as a function of soil pH,
extractable iron content, and soluble sulfate 5.24
Table 5.8. Lead aqueous species included in the speciation calculations 5.28
Table 5.9. Estimated range of Kd values for lead as a function of soil pH, and
equilibrium lead concentrations 5.33
Table 5.10. Plutonium aqueous species included in the speciation calculations 5.38
Table 5.11. Estimated range of Kd values for plutonium as a function of the soluble
carbonate and soil clay content values 5.43
Table 5.12. Strontium aqueous species included in the speciation calculations 5.48
Table 5.13. Look-up table for estimated range of Kd values for strontium based on
CEC (meq/100 g), clay content (wt.%), and pH 5.53
Table 5.14. Thorium aqueous species included in the speciation calculations 5.56
Table 5.15. Look-up table for thorium Kd values (ml/g) based on pH and
dissolved thorium concentrations 5.63
xv
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Table 5.16. Uranium(VI) aqueous species included in the speciation calculations 5.69
Table 5.17. Look-up table for estimated range of Kd values for uranium based on pH 5.75
Table 5.18. Selected chemical and transport properties of the contaminants 5.78
Table 5.19. Distribution of dominant contaminant species at 3 pH
values for an oxidizing water described in Tables 5.1 and 5.2 5.79
Table 5.20. Some of the more important aqueous- and solid-phase parameters
affecting contaminant sorption 5.81
Table C.l. Descriptive statistics of the cadmium Kd data set for soils C.3
Table C.2. Correlation coefficients (r) of the cadmium Kd data set for soils C.4
Table C.3. Look-up table for estimated range of Kd values for cadmium based on pH C.5
Table C.4. Cadmium Kd data set for soils C.6
Table D.l. Descriptive statistics of cesium Kd data set including
soil and pure mineral phases D.3
Table D.2. Descriptive statistics of data set including soils only D.4
Table D.3. Correlation coefficients (r) of the cesium Kd value data set that
included soils and pure mineral phases D.6
Table D.4. Correlation coefficients (r) of the soil-only data set D.6
Table D.5. Effect of mineralogy on cesium exchange D.9
Table D.6 Cesium Kd values measured on mica (Fithian illite) via adsorption
and desorption experiments D.10
Table D.7. Approximate upper limits of linear range of adsorption isotherms on
various solid phases D.l 1
Table D.8. Fruendlich equations identified in literature for cesium D. 13
Table D.9. Descriptive statistics of the cesium Freundlich equations (Table D.8)
reported in the literature D. 15
Table D.10. Estimated range of Kd values (ml/g) for cesium based on CEC
xvi
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or clay content for systems containing <5% mica-like
minerals in clay-size fraction and <10"9 M aqueous cesium D.18
Table D. 11. Estimated range of Kd values (ml/g) for cesium based on CEC
or clay content for systems containing >5% mica-like
minerals in clay-size fraction and <10"9 M aqueous cesium D.18
Table D. 12. Calculations for values used in look-up table D.19
Table D.13. Cesium Kd data base for soils and pure mineral phases D.20
Table D. 14. Cesium Kd data set for soils only D.27
Table E. 1. Summary of Kd values for Cr(VI) adsorption on soils E.5
Table E.2. Data from Rai et al. (1988) for the adsorption of Cr(VI) as a function of pH .... E.8
Table E.3. Estimated range of Kd values for Cr(VI) as a function of soil pH,
extractable iron content, and soluble sulfate E.9
Table E.4. Data from Rai et al. (1988) on effects of competing anions on Cr(VI)
adsorption on Cecil/Pacolet soil E.l 1
Table E.5. Data from Rai et al. (1988) on effects of competing anions on Cr(VI)
adsorption on Kenoma soil E.12
TableF.l. Summary of Kd values for lead adsorption on soils F.5
Table F.2. Estimated range of Kd values for lead as a function of soil pH, and
equilibrium lead concentrations F.8
Table G.l. Plutonium adsorption data for soil samples G.10
Table G.2. Regression models for plutonium adsorption G.13
Table G.3. Estimated range of Kd values for plutonium as a function of the soluble
carbonate and soil clay content values G.13
Table H.l. Descriptive statistics of strontium Kd data set for soils H.3
Table H.2. Correlation coefficient (r) of the strontium Kd data set for soils H.4
Table H.3. Simple and multiple regression analysis results involving
strontium Kd values, CEC (meq/100 g), pH, and clay content (percent) H.8
xvii
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Table H.4. Look-up table for estimated range of Kd values for strontium based
on CEC and pH H.10
Table H.5. Look-up table for estimated range of Kd values for strontium based on
clay content and pH H.10
Table H.6. Calculations of clay content using regression equations containing
CEC as a independent variable H. 11
Table H.7. Strontium Kd data set for soils H. 12
Table H.8. Strontium Kd data set for pure mineral phases and soils H. 16
Table 1.1. Descriptive statistics of thorium Kd value data set presented in Section 1.3 1.3
Table 1.2. Correlation coefficients (r) of the thorium Kd value data set presented
in Section 1.3 1.4
Table 1.3. Calculated aqueous speciation of thorium as a function of pH 1.5
Table 1.4. Regression coefficient and their statistics relating thorium Kd values and pH 1.6
Table 1.5. Look-up table for thorium Kd values (ml/g) based on pH and
dissolved thorium concentrations 1.7
Table 1.6. Data set containing thorium Kd values 1.9
Table J.l. Uranium Kd values (ml/g) listed by Warnecke et al. (1994, Table 1) J. 12
Table J.2. Uranium Kd values 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 J. 17
xviii
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Table J.3. Geometric mean uranium Kd values derived by Thibault et al.
(1990) for sand, loam, clay, and organic soil types J. 18
Table J.4. Look-up table for estimated range of Kd values for uranium based on pH J.22
Table J.5. Uranium Kd values selected from literature for development
of look-up table J.29
xix
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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,L2 values and
their use in formulating the retardation factor, Rp The contaminant retardation factor (R,) 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 two-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 Kd 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 of this volume noted, "Kd's are the coin of the
realm in this business." For better or worse, the Kd model is 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 Volume n.
Particular attention is directed at providing an understanding of: (1) the use of Kd values in formulating
Rfi (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.
1 Throughout this report, the term "partition coefficient" will be used to refer to the Kd "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.
2 A list of acronyms, abbreviations, symbols, and notation is given in Appendix A. A list of
definitions is given in Appendix B
3 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.
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The Kd parameter is very important in estimating the potential for the adsorption of dissolved
contaminants in contact with soil. As typically used in fate and contaminant transport calculations, the
Kd 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 errors 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 conservatism into their calculations by selecting limiting or bounding conservative Kd values.
For example, the most conservative (i.e.., maximum) estimate from the perspective of off-site risks due
to contaminant migration through the 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). Consequently, 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. Due to
groundwater flow paths, site characteristics, or environmental uncertainties, 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. For those
situations that are sensitive to the selected Kd value, site-specific Kd values are essential.
The Kd is usually a measured parameter that is obtained from laboratory experiments. The 5 general
methods used to measure Kd values are reviewed. These methods include the batch laboratory
method, the column laboratory method, field-batch method, field modeling method, and Koc method.
The summary identifies what the 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 and briefly
described in Chapter 4 of Volume n. 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 is 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
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reviewed.
The main focus of Volume II 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 include cadmium, cesium, chromium, lead, plutonium, radon, strontium, thorium,
tritium (3H), and uranium. The selection of these contaminants by EPA and PNNL project staff was
based on 2 criteria. First, the contaminant had to be of high priority to the 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 - cadmium, cesium, plutonium, strontium, thorium, and uranium(VI).
Anions - chromium(VI) (as chromate) and uranium(VI).
Radionuclides - cesium, plutonium, radon, strontium, thorium, tritium (3H), and uranium.
Conservatively transported contaminants - tritium (3H) and radon.
Nonconservatively transported contaminants - other than tritium (3H) and radon.
Redox sensitive elements - chromium, plutonium, 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 3 solutes show
no or very little adsorption.
The concentration of contaminants in groundwater is controlled primarily by the amount of contaminant
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 soils. Dissolution/precipitation is more likely to be the
key process where chemical nonequilibium 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
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a concentration gradient {i.e., Fick'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 K, Model
The simplest and most common method of estimating contaminant retardation is based on the partition
(or distribution) coefficient, Kd. 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. Particular
attention is directed at issues relevant to the selection of Kd 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 + Ci = A (2.1)
the mass action expression for Kd is
Kj = Mass of Adsorbate Sorbed = _A; (2.1)
Mass of Adsorbate in Solution C;
where A = free or unoccupied surface adsorption sites
C, = total dissolved adsorbate remaining in solution at equilibrium
A; = 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 C; and that the activity of A; is equal to 1.
Chemical retardation, Rg is defined as,
Rf=vp/vc (2.2)
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
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is actually less than 1, and such circumstances are thought to be caused by anion exclusion (See
Volume I, Section 2.8). 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 moisture conditions,
the Rf is defined as
Rf=l+(pb/ne)Kd (2.3)
where pb = porous media bulk density (mass/length3)
r^ = 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 use in 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 1 sorbate (contaminant) and 1 sorbent (soil), and (6) all adsorption sites are
accessible and have equal strength. The last point is especially limiting for groundwater contaminant
models because it requires that Kd values should be used only to predict transport in systems chemically
identical to those used in the laboratory measurement of 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-Kd 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) can
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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 R, , when called 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 Kd 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
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 is
presented in Volume I primarily to provide a paradigm for sorption processes.
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3.0 Methods, Issues, and Criteria for Measuring Kd Values
There are 5 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. These
methods and the associated technical issues are described in detail in Chapter 3 of Volume I. 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. The 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. The primary advantage of the method is that such experi-
ments 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 experi-
ments 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).
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 experiment 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
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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, than 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 (nonequilibium) 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 1 mechanism of adsorption or desorption over another cannot be predicted a priori
and therefore generalizing the results from 1 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.
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 organic compounds. 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 and source term data to calculate a
Kd value. One key drawback to this technique is that it is very model dependent. Because the
calculated Kd value are model dependent and highly site specific, the Kd values must 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 include: using simple versus complex systems to measure Kd values, field
variability, the "gravel issue," and the "colloid issue." Soils are a complex mixture containing solid,
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gaseous, and liquid phases. Each phase contains several different constituents. The use of simplified
systems containing single mineral phases and aqueous phases with 1 or 2 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. 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 <1 percent of the
soil volume but is the most chemically reactive fraction, would 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 has a Kd = 0 ml/g. Although this assumption is incorrect (ie., 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 in the smaller fraction are similar to those that occur
in the larger fraction.
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 heterogenous 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
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 appreciable different Kd values. Examples are presented
in this volume.
Contaminant transport models generally treat the subsurface environment as a 2-phase system in which
contaminants are distributed between a mobile aqueous phase and an immobile solid phase (e.g., soil).
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An increasing body of evidence indicates that under some subsurface conditions, components of the
solid phase may exist as 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 2 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. Secondly, it is difficult to unambiguously
delineate between the contaminants in the mobile-aqueous and mobile-solid phases.
Often Kd values used in transport models are selected to provide a conservative estimate of
contaminant migration or health effects. However, the same Kd value would not provide a conservative
estimate for clean-up calculations. 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.
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 6 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.
1 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.
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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.
The MINTEQA2 computer code is an equilibrium chemical reaction model. It was developed with
EPA funding by originally combining the mathematical structure of the MDS1EQL 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.
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 [chromium, cadmium, cesium,
tritium (3H), lead, plutonium, radon, strontium, thorium, and uranium], the MINTEQA2 thermodynamic
database contains speciation and solubility reactions for chromium, including the valence states Cr(II),
Cr(m), and Cr(VI); cadmium; lead; strontium; and uranium, including the valence states U(m), U(IV),
U(V), and U(VI). Some of the thermodynamic data in the EPA version have been superseded in other
users' databases by more recently published data.
The MINTEQA2 code includes 7 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 7 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. The user
however must make assumptions concerning the composition and mass of the dominant sorptive
substrate, and supply the adsorption parameters for surface-complexation constants for the
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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 several waste disposal issues. The EPA in
its "Soil Screening Guidance" determined MINTEQA2-estimated Kd values for barium, beryllium,
cadmium, Cr(m), 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.
Chemical reaction models can be used, however, to support evaluations of Kd values and related
contaminant migration and risk assessment modeling predictions. Chemical reaction codes can be used
to calculate aqueous complexation to determine the ionic state and composition of the dominant 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.
Chemical reaction models can be used to predict bounding, 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. Chemical
reaction models can also be used to analyze initial and final geochemical conditions associated with
laboratory Kd measurements to determine if the measurement had been affected by processes such as
mineral precipitation which might have compromised the derived Kd values. Although chemical reaction
models cannot be used to predict Kd values, they 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.
4.2
-------
5.0 Contaminant Geochemistry and Kd Values
The important geochemical factors affecting the sorption1 of cadmium (Cd), cesium (Cs), chromium
(Cr), lead (Pb), plutonium (Pu), radon (Rn), strontium (Sr), thorium (Th), tritium (3H), and uranium (U)
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) identify, when
possible, minimum and maximum conservative Kd values for each contaminant as a function key
geochemical processes affecting their sorption.
5.1 General
Important chemical speciation, (co)precipitation/dissolution, and adsorption/desorption processes of
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 2 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 al., 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 convenience as a proxy for the composition of a shallow groundwater. Obviously, there are
significant differences between surface waters and groundwaters, and considerable variability in the
concentrations of various constituents in surface and groundwaters. For example, the concentrations of
1 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
3-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, Kd.
5.1
-------
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 groundwaters,
especially deep groundwaters, will likely be more reducing that 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, 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, 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 materials. The species distributions of
thorium and uranium were also modeled using pure water, free of any ligands other than hydroxyl ions,
to show the effects of hydrolysis in the absence of other complexation reactions. The concentrations
used for the dissolved contaminants in the species distribution calculations are presented in Table 5.2
and are further discussed in the following sections. The species distributions of cesium, radon, and
tritium were not determined because only 1 aqueous species is likely to exist under the environmental
conditions under consideration; namely, cesium would exist as Cs+, radon as Rn°(gas), and tritium as
tritiated water, HTO (T = tritium, 3H).
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, the
variables had include parameters that were readily available to modelers. For instance, particle size and
pH are often available to modelers whereas such parameters as iron oxide or surface area are not as
frequently available.
1 "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).
2 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).1
Dissolved Constituent
Total Concentration
mg/1
mol/l
Silica, as HSiC^
20.8
2.16 x 10"4
Ca
15
3.7 x 10"4
Mg
4.1
1.7 x 10"4
Na
6.3
2.7 x 10"4
K
2.3
5.9 x 10"5
Inorganic Carbon, as C03
57
9.5 x 10"4
S04
11
1.1 x 10"4
CI
7.8
2.2 x 10"4
F
1
5 x 10"5
no3
1
2x 10"5
P04
0.0767
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 micro g/1 for average
concentration of total dissolved phosphorous in river water estimated by
Meybeck (1982). This concentration of total phosphorus was converted
to total phosphate (P04) listed above.
5.3
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Table 5.2. Concentrations of contaminants used in the aqueous
species distribution calculations.
Element
Total Cone,
(micro g/1)
Reference for Concentration of Contaminant
Used in Aqueous Speciation Calculations
Cd
1.0
Hem (1985, p. 142) lists this value as a median concentration of dissolved cadmium
based on the reconnaissance study of Duram etal. (1971) of metal concentrations in
surface waters in the United States.
Cs
--
Distribution of aqueous species was not modeled, because mobility of dissolved
cesium is not significantly affected by complexation (see Section 5.3).
Cr
1.4
Hem (1985, p. 138) lists this value as an average concentration estimated by Kharkar
et al. (1968) for chromium in river waters.
Pb
1.0
Hem (1985, p. 144) lists this value as an average concentration estimated by Duram
etal. (1971) for lead in surface-water samples from north- and southeastern sections
of the United States.
Pu
3.2 xlO"7
This concentration is based on the maximum activity of 239'240pu measured by
Simpson etal. (1984) in 33 water samples taken from the highly alkaline Mono Lake
in California.
Rn
--
Aqueous speciation was not calculated, because radon migrates as a dissolved gas
and is not affected by complexation (see Section 5.7).
Sr
110
Hem (1985, p. 135) lists this value as the median concentration of strontium for larger
United States public water supplies based on analyses reported by Skougstad and
Horr (1963).
Th
1.0
Hem (1985, p. 150) gives 0.01 to 1 micro g/1 as the range expected for thorium
concentrations in fresh waters.
3H
--
Aqueous speciation was not calculated, because tritium (3H) migrates as tritiated
water.
U
0.1 and
1,000
Because dissolved hexavalent uranium can exist as polynuclear hydroxyl complexes,
the hydrolysis of uranium under oxic conditions is therefore dependent on the
concentration of total dissolved uranium. To demonstrate this aspect of uranium
chemistry, 2 concentrations (0.1 and 1,000 micro g/1) of total dissolved uranium were
used to model the species distributions. Hem (1985, p. 148) gives 0.1 to 10 microg/1
as the range for dissolved uranium in most natural waters. For waters associated
with uranium ore deposits, Hem states that the uranium concentrations may be
greater than 1,000 microg/1.
5.4
-------
5.2 Cadmium Geochemistry and Kd Values
5.2.1 Overview: Important Aqueous-and Solid-Phase Parameters
Controlling Retardation
The dominant cadmium aqueous species in groundwater at pH values less than 8.2 and containing
moderate to low concentrations of sulfate (<10"2 5 M SO4") is the uncomplexed Cd2+ species. The
dominant cadmium solution species in groundwater at pH values greater than 8.2 are CdC03 (aq) and
to a smaller extent CdCl . Both precipitation/coprecipitation/dissolution and adsorption/desorption
reactions control cadmium concentrations. Several researchers report that otavite (CdC03) limits
cadmium solution concentrations in alkaline soils. The solid Cd3(P04)2 has also been reported to be a
solubility-controlling solid for dissolved cadmium. Under low redox conditions, sulfide concentrations
and the formation of CDs precipitates may play an important role in controlling the concentrations of
dissolved cadmium. At high concentrations of dissolved cadmium (>10"7 M Cd), either cation
exchange or (co)precipitation are likely to control dissolved cadmium concentrations. Precipitation with
carbonate is increasingly important in systems with a pH greater than 8, and cation exchange is more
important in lower pH systems. At lower environmental concentrations of dissolved cadmium, surface
complexation with calcite and aluminum- and iron-oxide minerals may be the primary process
influencing retardation. Transition metals (e.g., copper, lead, zinc) and alkaline earth (e.g., calcium,
magnesium) cations reduce cadmium adsorption by competition for available specific adsorption and
cation exchange sites. In conclusion, the key aqueous- and solid-phase parameters influencing
cadmium adsorption include pH, cadmium concentration, competing cation concentrations, redox,
cation exchange capacity (CEC), and mineral oxide concentrations.
5.2.2 General Geochemistry
Cadmium (Cd) exists in the +2 oxidation state in nature. It forms a number of aqueous complexes,
especially with dissolved carbonate. Its concentration may be controlled by either adsorption or
precipitation/coprecipitation processes. The extent to which cadmium is associated with or bound to
soils varies greatly with type of mineral, oxidation state of the system, and presence of competing
cations in solution.
Cadmium concentrations in uncontaminated soils is typically less than 1 mg/kg. However,
concentrations may be significantly elevated by some human activities or by the weathering of parent
materials with high cadmium concentrations, e.g., black shales (Jackson and Alloway, 1992).
Approximately 90 percent of all the cadmium consumed goes into 4 use categories: plating (35
percent), pigments (25 percent), plastic stabilizers (15 percent), and batteries (15 percent) (Nriagu,
1980b). Cadmium may also be introduced into the environment by land applications of sewage sludge.
Cadmium concentrations in sewage sludge are commonly the limiting factor controlling land disposal
(Juste and Mench, 1992). Nriagu (1980a) has edited an excellent review on the geochemistry and
toxicity of cadmium.
5.5
-------
5.6
-------
5.2.3 Aqueous Speciation
Cadmium forms soluble complexes with inorganic and organic ligands resulting in an increase of
cadmium mobility in soils (McLean and Bledsoe, 1992). The distribution of cadmium aqueous species
was calculated using the water composition described in Table 5.1 and a concentration of 1 micro g/1
total dissolved cadmium (Table 5.2). Hem (1985, p. 142) lists this value as a median concentration of
dissolved cadmium based on the reconnaissance study of Duram el al. (1971) of metal concentrations
in surface waters in the United States. These MINTEQA2 calculations indicate that cadmium
speciation is relatively simple. In groundwaters of pH values less than 6, essentially all of the dissolved
cadmium is expected to exist as the uncomplexed Cd2+ ion (Figure 5.1). The aqueous species included
in the MINTEQA2 calculations are listed in Table 5.3. As the pH increases between 6 and 8.2,
cadmium carbonate species [CdHC03 and CdCO^ (aq)] become increasingly important. At pH
values between 8.2 and 10, essentially all of the cadmium in solution is expected to exist as the neutral
complex CdCC>3 (aq). The species CdSOij (aq), CdHC03, CdCf, and CdOH+ are also present, but
at much lower concentrations. The species distribution illustrated in Figure 5.1 does not change if the
concentration of total dissolved cadmium is increased from 1 to 1,000 micro g/1.
Table 5.3. Cadmium aqueous species included
in the speciation calculations.
Aqueous Species
Cd2+
CdOH+, Cd(OH)2 (aq), Cd(OH)i, Cd(OH)^", Cd2OH3+
CdHCOs, CdCO; (aq), Cd(C03)^
CdSC>4 (aq), Cd(S04)i"
CdNOs
CdCl+, CdClj (aq), CdCli, CdOHCl" (aq)
CdF+, CdF; (aq)
5.7
-------
5.8
-------
5
¦g
*c
5/5
5
fi
(J
~-
CL
pH
Figure 5.1. Calculated distribution of cadmium aqueous species as a function of pH for the water
composition in Table 5.1. [The species distribution is based on a concentration of
1 micro g/1 total dissolved cadmium and thermodynamic data supplied with the
MINTEQA2 geochemical code.]
Information available in the literature regarding interactions between dissolved cadmium and naturally
occurring organic ligands (humic and fulvic acids) is ambiguous. Weber and Posselt (1974) reported
that cadmium can form stable complexes with naturally occurring organics, whereas Hem (1972) stated
that the amount of cadmium occurring in organic complexes is generally small and that these complexes
are relatively weak. Pittwell (1974) reported that cadmium is complexed by organic carbon under all
pH conditions encountered in normal natural waters. Levi-Minzi el al. (1976) found cadmium
adsorption in soils to be correlated with soil organic matter content. In a critical review of the literature,
Giesy (1980) concluded that the complexation constants of cadmium to naturally occurring organic
matter are weak because of competition for binding sites by calcium, which is generally present in much
higher concentrations.
5.9
-------
5.2.4 Dissolution/Precipitation/Coprecipitation
Lindsay (1979) calculated the relative stability of cadmium compounds. His calculations show that at
pH values less than 7.5, most cadmium minerals are more soluble than cadmium concentrations found in
oxic soils (10"7 M), indicating that cadmium at these concentrations is not likely to precipitate. At pH
levels greater than 7.5, the solubilities of Cd3(P04)2 or CdC03 may control the concentrations of
cadmium in soils. Cavallaro and McBride (1978) and McBride (1980) demonstrated that otavite,
CdC03, precipitates in calcareous soils (pH > 7.8), whereas in neutral or acidic soils, adsorption is the
predominate process for removal of cadmium from solution. Jenne el al. (1980), working with the
waters associated with abandoned lead and zinc mines and tailings piles, also indicate that the upper
limits on dissolved levels of cadmium in most waters were controlled by CdC03. Santillan-Medrano
and Jurinak (1975) observed that the activity of dissolved cadmium in cadmium-amended soils was
lowest in calcareous soils. Baes and Mesmer (1976) suggested that cadmium may coprecipitate with
calcium to form carbonate solid solutions, (Ca,Cd)C03. This may be an important mechanism in
controlling cadmium concentrations in calcareous soils.
Although cadmium itself is not sensitive to oxidation/reduction conditions, its concentration in the
dissolved phase is generally very sensitive to redox state. There are numerous studies (reviewed by
Khalid, 1980) showing that the concentrations of dissolved cadmium greatly increase when reduced
systems are oxidized, such as when dredged river sediments are land filled or rice paddies are drained.
The following 2 mechanisms appear to be responsible for this increase in dissolved cadmium
concentrations: (1) very insoluble CDs (greenockite) dissolves as sulfide [S(II)] that is oxidized to
sulfate [S(VI)], and (2) organic materials binding cadmium are decomposed through oxidization,
releasing cadmium into the environment (Gambrell et al., 1977; Giesy, 1980). This latter mechanism
appears to be important only in environments in which moderate to high organic matter concentrations
are present (Gambrell el al., 1977). Serne (1977) studied the effect of oxidized and reduced sediment
conditions on the release of cadmium from dredged sediments collected from the San Francisco Bay.
Greater than 90 percent of the cadmium in the reduced sediment [sediment incubated in the presence of
low 02 levels (Eh<100 mV)] was complexed with insoluble organic matter or precipitated as sulfides.
The remainder of the cadmium was associated with the oxide minerals, clay lattices, or exchangeable
sites. Dissolved cadmium concentrations greatly increased when the sediments were incubated under
oxidizing conditions (Eh>350 mV). Cadmium concentrations released in the elutriate increased with
agitation time. These data suggested that this kinetic effect was due to slow oxidation of sulfide or
cadmium bound to organic matter bound in the reduced sediment prior to steady state equilibrium
conditions being reached. In a similar type of experiment in which Mississippi sediments were slowly
oxidized, Gambrell el al. (1977) reported that the insoluble organic- and sulfide-bound cadmium
fractions in sediment decreased dramatically (decreased >90 percent) while the exchangeable and
water-soluble cadmium fractions increased. Apparently, once the cadmium was released from the
sulfide and organic matter fractions, the cadmium entered the aqueous phase and then re-adsorbed
onto other sediment phases.
5.10
-------
A third mechanism involves pyrite that may be present in soils or sediments and gets oxidized when
exposed to air.1 The pyrite oxidizes to form FeS04, which generates high amounts of acidity when
reacted with water. The decrease in the pH results in the dissolution of cadmium minerals and increase
in the dissolved concentration of cadmium. This process is consistent with the study by Kargbo (1993)
of acid sulfate clays used as waste covers.
5.2.5 Sorption/Desorption
At high solution concentrations of cadmium (>10 mg/1), the adsorption of cadmium often correlates with
the CEC of the soil (John, 1971; Levi-Minzi el al., 1976; McBride et al., 1981; Navrot et al., 1978;
Petruzelli el al., 1978). During cation exchange, cadmium generally exchanges with adsorbed calcium
and magnesium (McBride el al., 1982). The ionic radius of Cd2+ is comparable to that of Ca2+ and, to
a lesser extent, Mg2+. At low solution concentrations of cadmium, surface complexation to calcite
(McBride, 1980) and hydrous oxides of aluminum and iron (Benjamin and Leckie, 1981) may be the
most important adsorption mechanism. Both Cd2+ and possibly CdOH+ may adsorb to aluminum- and
iron-oxide minerals (Balistrieri and Murray, 1981; Davis and Leckie, 1978).
As with other cationic metals, cadmium adsorption exhibits pH dependency. The effect of pH on
cadmium adsorption by soils (Huang el al., 1977), sediment (Reid and McDuffie, 1981), and iron
oxides (Balistrieri and Murray, 1982; Levy and Francis, 1976) is influenced by the solution
concentration of cadmium and the presence of competing cations or complexing ligands. At low
cadmium solution concentrations, sharp adsorption edges (the range of pH where solute adsorption
goes from ~0 to -100 percent) suggests that specific adsorption {i.e., surface complexation via a strong
bond to the mineral surface) occurs. Under comparable experimental conditions, the adsorption edge
falls at pH values higher than those for lead, chromium, and zinc. Thus, in lower pH environments,
these metals, based on their propensity to adsorb, would rank as follows: Pb > Cr > Zn > Cd. This
order is inversely related to the pH at which hydrolysis of these metals occurs (Benjamin and Leckie,
1981).
Competition between cations for adsorption sites strongly influences the adsorption behavior of
cadmium. The presence of calcium, magnesium, and trace metal cations reduce cadmium adsorption
by soils (Cavallaro and McBride, 1978; Singh, 1979), iron oxides (Balistrieri and Murray, 1982),
manganese oxides (Gadde and Laitinen, 1974), and aluminum oxides (Benjamin and Leckie, 1980).
The extent of competition between cadmium and other ions depends on the relative energies of
interaction between the ions and the adsorbing surface, the concentrations of the competing ions, and
solution pH (Benjamin and Leckie, 1981; Sposito, 1984). The addition of copper or lead, which are
more strongly adsorbed, slightly reduces cadmium adsorption by iron and aluminum oxides, suggesting
that copper and lead are preferentially adsorbed by different surface sites (Benjamin and Leckie,
1 D. M. Kargbo (1998, personal communication).
5.11
-------
1980). In contrast, zinc almost completely displaces cadmium, indicating that cadmium and zinc
compete for the same group of binding sites (Benjamin and Leckie, 1981).
Although organic matter may influence adsorption of cadmium by soils (John, 1971; Levi-Minzi el al.,
1976), this effect is probably due to the CEC of the organic material rather than to complexation by
organic ligands (Singh and Sekhon, 1977). In fact, removal of organic material from soils does not
markedly reduce cadmium adsorption and may enhance adsorption (Petruzelli el al., 1978). Clay
minerals with adsorbed humic acids (organo-clay complexes) do not adsorb cadmium in excess of that
expected for clay minerals alone (Levy and Francis, 1976).
5.2.6 Partition Coefficient, Kd, Values
5.2.6.1 General Availability of Kd Data
A total of 174 cadmium Kd values were found in the literature and included in the data base used to
create the look-up tables.1 The cadmium Kd values as well as the ancillary experimental data are
presented in Appendix C. Data included in this table were from studies that reported Kd values (not
percent adsorption or Langmuir constants) and were conducted in systems consisting of natural soils (as
opposed to pure mineral phases), low ionic strength (< 0.1 M), pH values between 4 and 10, low
humic material concentrations (<5 mg/1), and no organic chelates (e.g., EDTA). At the start of the
literature search, attempts were made to identify cadmium Kd studies that reported ancillary data on
aluminum/iron-oxide concentrations, calcium and magnesium solution concentrations, CEC, clay
content,2 pH, redox status, organic matter concentrations and sulfide concentrations. Upon reviewing
the data and determining the availability of cadmium Kd studies reporting ancillary data, we selected
data on clay content, pH, CEC, and total organic carbon. The selection of these parameters was
based on availability of data and the possibility that the parameter may impact cadmium Kd values. Of
the 174 cadmium Kd values included in the compiled data, only 62 values had associated clay content
data, 174 values had associated pH data, 22 values had associated CEC data, 63 values had total
organic carbon data, and 16 had associated aluminum/iron-oxide data. Descriptive statistics and a
correlation coefficient matrix are presented in Appendix C.
1 Since the completion of our review and analysis of Kd data for the selected contaminants and
radionuclides, the study by Wang el al. (1998) was identified and may be of interest to the reader.
2 Unless specified otherwise, "clay content" refers to the particle size fraction of soil that is less
than 2 micro m.
5.12
-------
5.2.6.2 Look-Up Tables
One cadmium Kd look-up table was created. The table requires knowledge of the pH of the system
(Table 5.4). The pH was selected as the key independent variable because it had a highly significant (P
< 0.001) correlation with cadmium Kd, a correlation coefficient value of 0.75. A detailed explanation
of the approach used in selecting the Kd values used in the table is presented in Appendix C. Briefly, it
involved conducting a regression analysis between pH and Kd values). The subsequent regression
equation was used to provide central estimates. Minimum and maximum values were estimated by
plotting the data and estimating where the limits of the data existed.
There is an unusually wide range of possible cadmium Kd values for each of the 3 pH categories. The
cause for this is likely that there are several other soil parameters influencing the Kd in addition to pH.
Unfortunately, the correlations between the cadmium Kd values and the other soil parameters in this
data set were not significant (Appendix C).
5.2.6.2.1 Limits of Kd Values With Respect to Aluminum/Iron-Oxide Concentrations
The effect of iron-oxide concentrations on cadmium Kd values was evaluated using the data presented
in Appendix C. Of the 174 cadmium Kd values in the data set presented in Appendix C, only 16
values had associated iron oxide concentration data. In each case iron, and not aluminum, oxide
concentration data were measured. The correlation coefficient describing the linear relationship
between cadmium Kd values and iron oxide concentration was 0.18, which is nonsignificant at the 5
percent level of probability. It was anticipated that there would be a positive correlation between iron
or aluminum oxide concentrations and cadmium Kd values because oxide minerals provide adsorption
(surface complexation) sites.
Table 5.4. Estimated range of Kd values for cadmium as a function of pH.
[Tabulated values pertain to systems consisting of natural soils (as
opposed to pure mineral phases), low ionic strength (< 0.1 M),
low humic material concentrations (<5 mg/1), no organic chelates
(e.g., EDTA), and oxidizing conditions.]
pH
Kd (ml/g)
3-5
5-8
8-10
Mnimum
1
8
50
Maximum
130
4,000
12,600
5.13
-------
5.2.6.2.2 Limits of Kd Values with Respect to CEC
The effect of CEC on cadmium Kd values was evaluated using the data presented in Appendix C.
Of the 174 cadmium Kd values in the data set presented in Appendix C, only 22 values had associated
CEC data. The correlation coefficient describing the linear relationship between cadmium Kd values and
CEC was 0.40, which is nonsignificant at the 5 percent level of probability. It was anticipated that
there would be a positive correlation between CEC and cadmium Kd values because cadmium can
adsorb to minerals via cation exchange.
5.2.6.2.3 Limits of Kd Values with Respect to Clay Content
The effect of clay content on cadmium Kd values was evaluated using the data presented in Appendix
C. Of the 174 cadmium Kd values in the data set presented in Appendix C, 64 values had associated
clay content data. The correlation coefficient describing the linear relationship between cadmium Kd
values and clay content was -0.04, which is nonsignificant at the 5 percent level of probability. It was
anticipated that there would be a positive correlation between clay content and cadmium Kd values,
because clay content is often highly correlated to CEC, which in turn may be correlated to the number
of sites available for cadmium adsorption.
5.2.6.2.4 Limits of Kd Values with Respect to Concentration of Organic Matter
The effect of organic matter concentration, as approximated by total organic carbon, on cadmium Kd
values was evaluated using the data presented in Appendix C. Of the 174 cadmium Kd values in the
data set presented in Appendix C, 63 values had associated total organic carbon concentration data.
The correlation coefficient describing the linear relationship between cadmium Kd values and total
organic carbon concentration was 0.20, which is nonsignificant at the 5 percent level of probability. It
was anticipated that there would be a positive correlation between total organic carbon concentration
and cadmium Kd values because soil organic carbon can have extremely high CEC values, providing
additional sorption sites for dissolved cadmium.
5.2.6.2.5 Limits of Kd Values with Respect to Dissolved Calcium, Magnesium, and Sulfide
Concentrations, and Redox Conditions
Calcium, magnesium, and sulfide solution concentrations were rarely, if at all, reported in the
experiments used to comprise the cadmium data set. It was anticipated that dissolved calcium and
magnesium would compete with cadmium for adsorption sites, thereby decreasing Kd values. It was
anticipated that sulfides would induce cadmium precipitation, thereby increasing cadmium Kd values.
Similarly, low redox status was expected to provide an indirect measure of sulfide concentrations,
which would in turn induce cadmium precipitation. Sulfides only exist in low redox environments; in
high redox environments, the sulfides oxidize to sulfates that are less prone to form cadmium
precipitates.
5.14
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5.15
-------
5.3 Cesium Geochemistry and Kd Values
5.3.1 Overview: Important Aqueous-and Solid-Phase Parameters
Controlling Retardation
The aqueous speciation of cesium in groundwater is among the simplest of the contaminants being
considered in this study. Cesium forms few stable complexes and is likely to exist in groundwater as
the uncomplexed Cs+ ion, which adsorbs rather strongly to most minerals, especially mica-like clay
minerals. The extent to which adsorption will occur will depend on (1) the concentration of mica-like
clays in the soil, and (2) the concentration of major cations, such as K+ which has a small ionic radius as
Cs+, that can effectively compete with Cs+ for adsorption sites.
5.3.2 General Geochemistry
Cesium (Cs) exists in the environment in the +1 oxidation state. Stable cesium is ubiquitous in the
environment with concentrations in soils ranging between 0.3 and 25 mg/kg (Lindsay, 1979). The only
stable isotope of cesium is 133Cs. Fission products include 4 main cesium isotopes. Of these, only
134Cs [half life (/.J = 2.05 y], 135Cs (tVl = 3 x 106 y), and 137Cs (tVl = 30.23 y) are at significant
concentrations 10 y after separation from nuclear fuels (Schneider and Piatt, 1974).
Contamination includes cesium-containing soils and cesium dissolved in surface- and groundwaters. Of
the contaminated sites considered in EPA/DOE/NRC (1993), radioactive contamination of soil, surface
water, and/or groundwater by 134Cs, 135Cs and/or 137 Cs has been identified at 9 of the 45 Superfund
National Priorities List (NPL) sites.
5.3.3 Aqueous Speciation
There is little, if any, tendency for cesium to form aqueous complexes in soil/water environments. Thus,
the formation of inorganic complexes is not a major influence on cesium speciation and the dominant
aqueous species in most groundwater is the uncomplexed Cs+ ion. Baes and Mesmer (1976) report
that cesium may be associated with OH" ions in solution, but that the extent of this association cannot be
estimated accurately. The uncomplexed Cs+ ion forms extremely weak aqueous complexes with
sulfate, chloride, and nitrate. Cesium also can form weak complexes with humic materials, as shown by
the following ranking of cations by their propensity to form complexes with humic materials (Bovard el
al., 1970):
Ce > Fe > Mn > Co > Ru > Sr > Cs
5.16
-------
Further, complexation of cesium by common industrial chelates (e.g., EDTA) is believed to be poor
due to their low stabilities and the presence of competing cations (e.g., Ca2+) at appreciably higher
concentrations than that of cesium. Therefore, aqueous complexation is not thought to greatly influence
cesium behavior in most groundwater systems.
5.3.4 Dissolution/Precipitation/Coprecipitation
Neither precipitation nor coprecipitation are expected to affect the geochemistry of cesium in
groundwater. The solubility of most cesium compounds in water is very high.
5.3.5 Sorption/Desorption
In general, most soils sorb cesium rather strongly (Ames and Rai, 1978). Some mica-like minerals,
such as illite {(K,H3O)(Al,Mg,Fe)2(Si,Al)4O10[(OH)2 H20]} and vermiculite
[(Mg,Fe,Al)3(Si,Al)4O10(OH)2-4H2O], tend to intercalate (fix) cesium between their structural layers
(Bruggenwert and Kamphorst, 1979; Douglas, 1989; Smith and Comans, 1996). These silicate
minerals can be thought of as having a crystal lattice composed of continuous sheet structures. The
distance between the silicate layers is controlled by the type of cation associated with the adsorption
sites on the layers. Large hydrated cations, such as Na+, Li+, Ca2+, and Mg2+, tend to pry the layers
further apart, whereas small hydrated cations, such as K+, have the opposite effect. The interlayer
distance between the sheets of mica-like minerals excludes the absorption of the majority of cations by
size, while permitting the Cs+ ion to fit perfectly between the layers. Consequently, these mica-like
minerals commonly exhibit a very high selectivity for Cs+ over other cations, including cations existing at
much higher concentrations. Even a small amount (e.g., 1-2 weight percent) of these mica-like minerals
in a soil may strongly absorb a large amount of dissolved cesium (Coleman el al., 1963; Douglas,
1989). Some researchers have considered the exchange of trace cesium on these mica-like minerals to
be nearly irreversible (Douglas, 1989; Routson, 1973), meaning that cesium absorbs at a much faster
rate than it desorbs.
The effect of cesium concentration and pH on cesium adsorption by a calcareous soil containing mica-
like minerals has been studied by McHenry (1954). The data indicate that trace cesium concentrations
are essentially completely adsorbed above pH 4.0. When placed in a high-salt solution, 4 M NaCl,
only up to 75 percent of the trace cesium was adsorbed, and the adsorption was essentially
independent of pH over a wide range. At cesium loadings on the soil of less than 1 percent of the soil
CEC, the effect of competing cations on cesium adsorption was slight. Low concentrations of
dissolved cesium are typical of cesium-contaminated areas. Thus competition may not play an
important role in controlling cesium adsorption in most natural groundwater environments. The results
of McHenry (1954) also indicate that trace concentrations of cesium were adsorbed to a greater
degree and were more difficult to displace from the soil by competing cations than when the cesium was
adsorbed at higher loadings.
5.17
-------
Cesium may also adsorb to iron oxides (Schwertmann and Taylor, 1989). Iron oxides, unlike mica-like
minerals, do not "fix" cesium. Instead they complex cesium to sites whose abundance is pH dependent;
i.e., iron oxides have variable charge surfaces. Iron oxides dominate the adsorption capacity of many
soils in semi-tropical regions, such as the southeastern United States. In these soils, many mica-like
minerals have been weathered away, leaving minerals with more pH-dependent charge. As the pH
decreases, the number of negatively charged complexation sites also decreases. For example, Prout
(1958) reported that cesium adsorption to iron-oxide dominated soils from South Carolina decreased
dramatically when the suspension pH was less than 6.
Cesium adsorption to humic materials is generally quite weak (Bovard el al., 1970). This is consistent
with cation ranking listed above showing that cesium forms relatively weak complexes with organic
matter.
5.3.6 Partition Coefficient, Kd, Values
5.3.6.1 General Availability of Kd Data
Three generalized, simplifying assumptions were established for the selection of cesium Kd values for
the look-up table. These assumptions were based on the findings of the literature review we conducted
on the geochemical processes affecting cesium sorption.1 The assumptions are as follows:
• Cesium adsorption occurs entirely by cation exchange, with the exception when mica-like
minerals are present. Cation exchange capacity (CEC), a parameter that is frequently not
measured, can be estimated by an empirical relationship with clay content and pH.
• Cesium adsorption into mica-like minerals occurs much more readily than desorption. Thus, Kd
values, which are essentially always derived from adsorption studies, will greatly overestimate
the degree to which cesium will desorb from these surfaces.
• Cesium concentrations in groundwater plumes are low enough, less than approximately 10"7
M, such that cesium adsorption follows a linear isotherm.
These assumptions appear to be reasonable for a wide range of environmental conditions. However,
these simplifying assumptions are clearly compromised in systems with cesium concentrations greater
than approximately 10"7 M, ionic strength levels greater than about 0.1 M, and pH levels greater than
about 10.5. These 3 assumptions will be discussed in more detail in the following sections.
1 Since the completion of our review and analysis of Kd data for the selected contaminants and
radionuclides, the studies by Cygan et al. (1998), Fisher el al. (1999), and Oscarson and Hume
(1998) were identified and may be of interest to the reader.
5.18
-------
Based on the assumptions and limitation described in above, cesium Kd values and some important
ancillary parameters that influence cation exchange were collected from the literature and tabulated.
Data included in this table were from studies that reported Kd values (not percent adsorbed or
Freundlich or Langmuir constants) and were conducted in systems consisting of: (1) low ionic strength
(< 0.1 M), (2) pH values between 4 and 10.5, (3) dissolved cesium concentrations less than 10"7 M,
(4) low humic material concentrations (<5 mg/1), and (5) no organic chelates (e.g., EDTA). Initially,
attempts were made to include in the Kd data set all the key aqueous and solid phase parameters
identified above. The key parameters included aluminum/iron-oxide mineral concentration, CEC, clay
content, potassium concentration, mica-like mineral content, ammonium concentration, and pH. The
ancillary parameters for which data could be found in the literature that were included in these tables
were clay content, mica content, pH, CEC, surface area, and solution cesium concentrations. This
cesium data set included 176 cesium Kd values. The descriptive statistics of the cesium Kd data set are
presented in Appendix D.
5.3.6.2 Look-Up Tables
Linear regression analyses were conducted with data collected from the literature. These analyses were
used as guidance for selecting appropriate Kd values for the look-up table. The Kd values used in the
look-up tables could not be based entirely on statistical consideration because the statistical analysis
results were occasionally nonsensible. For example, the data showed a negative correlation between
pH and CEC, and pH and cesium Kd values. These trends contradict well established principles of
surface chemistry. Instead, the statistical analysis was used to provide guidance as to the approximate
range of values to use and to identify meaningful trends between the cesium Kd values and the solid
phase parameters. Thus, the Kd values included in the look-up table were in part selected based on
professional judgment. Again, only low-ionic strength solutions, such as groundwaters, were
considered; thus no solution variables were included.
Two look-up tables containing cesium Kd values were created. The first table is for systems containing
low concentrations of mica-like minerals: less than about 5 percent of the clay-size fraction (Table 5.5).
The second table is for systems containing high concentrations of mica-like minerals (Table 5.6). For
both tables, the user will be able to reduce the range of possible cesium Kd values with knowledge of
either the CEC or the clay content. A detailed description of the assumptions and the procedures used
in coming up with these values is presented in Appendix D.
5.19
-------
Table 5.5. Estimated range of Kd values (ml/g) for cesium based on
CEC or clay content for systems containing <5 percent
mica-like minerals in clay-size fraction and <10"9 M
aqueous cesium. [Table pertains to systems consisting of
natural soils (as opposed to pure mineral phases), low
ionic strength (<0.1 M), low humic material
concentrations (<5 mg/1), no organic chelates (e.g.,
EDTA), and oxidizing conditions.]
Kd (ml/g)
CEC (meq/100 g) / Clay Content (wt.%)
<3/<4
3-10/4-20
10-50/20-60
Minimum
10
30
80
Maximum
3,500
9,000
26,700
Table 5.6. Estimated range of Kd values (ml/g) for cesium based on
CEC or clay content for systems containing >5 percent
mica-like minerals in clay-size fraction and <10"9 M
aqueous cesium. [Table pertains to systems consisting of
natural soils (as opposed to pure mineral phases), low
ionic strength (<0.1 M), low humic material concentrations
(<5 mg/1), no organic chelates (e.g., EDTA), and oxidizing
conditions.]
Kd (ml/g)
CEC (meq/100 g) / Clay Content (wt.%)
<3/<4
3-10/4-20
10-50/20-60
Minimum
30
70
210
Maximum
9,000
22,000
66,700
5.20
-------
5.3.6.2.1 Limits of Kd Values with Respect to pH
Of the 177 cesium Kd values obtained from the literature, 139 of them had associated pH values for the
system under consideration (Appendix D). The average pH of the systems described in the data set
was pH 7.4, ranging from pH 2.4 to 10.2. The correlation coefficient (r) between pH and cesium Kd
values was 0.05. This is clearly an insignificant correlation. This poor correlation may be attributed to
the fact that other soil properties having a greater impact on cesium Kd values were not held constant
throughout this data set.
5.3.6.2.2 Limits of Kd Values with Respect to Potassium, Ammonium, and Aluminum/Iron-Oxides
Concentrations
Potassium, ammonium, and aluminum/iron-oxide mineral concentrations were rarely, if at all, reported in
the experiments used to comprise the cesium Kd data set (Appendix D). It was anticipated that
dissolved potassium and ammonium would compete with cesium for adsorption sites, thereby
decreasing Kd values. The presence of aluminum and/or iron oxides in the solid phase was expected to
increase cesium Kd values.
5.4 Chromium Geochemistry and Kd Values
5.4.1 Overview: Important Aqueous- and Solid-Phase Parameters
Controlling Retardation
A plume containing high concentrations of chromium is more likely to be composed of Cr(VI) than
Cr(III) because the former is less likely to adsorb or precipitate to the solid phase. Chromium(VI) is
also appreciably more toxic than Cr(m). It exhibits significant subsurface mobility in neutral and basic
pH environments. In acid environments, Cr(VI) may be moderately adsorbed by pH-dependent
charge minerals, such as iron- and aluminum-oxide minerals. The reduction of Cr(VI) to Cr(m) by
ferrous iron, organic matter, and microbes is generally quite rapid whereas the oxidation of Cr(m) to
Cr(VI) by soil manganese oxides or dissolved oxygen is kinetically slower. The most important
aqueous- and solid-phase parameters controlling retardation of chromium include redox status, pH, and
the concentrations of aluminum- and iron-oxide minerals and organic matter.
5.4.2 General Geochemistry
Chromium is found in the environment primarily in the +3 and +6 oxidation states. The geochemical
behavior and biological toxicity of chromium in these 2 oxidation states are profoundly different.
Chromium(VI) tends to be soluble, forms anionic or neutral dissolved species, can be very mobile, and
is acutely toxic (Nriagu and Nieboer, 1988). In contrast, Cr(III) tends to precipitate, forms cationic
dissolved species, is immobile under moderately alkaline to slightly acidic conditions, and is relatively
nontoxic. The primary human activities leading to the introduction of chromium into the environment are
5.21
-------
ore processing, plating operations, and manufacturing (reviewed by Nriagu and Nieboer, 1988).
Discussions of the production, uses, and toxicology of chromium have been presented by Nriagu and
Nieboer (1988). Good review articles describing the geochemistry of chromium have been written by
Rai et al. (1988), Palmer and Wittbrodt (1991), Richard and Bourg (1991), and Palmer and Puis
(1994). A critical review of the thermodynamic properties for chromium metal and its aqueous ions,
hydrolysis species, oxides, and hydroxides was published by Ball and Nordstrom (1998).
5.4.3 Aqueous Speciation
Chromium exists in the +2, +3, and +6 oxidation states in water, of which only the +3 and +6 states are
found in the environment. Chromium(m) exists over a wide range of pH and Eh conditions, whereas
Cr(VI) exists only under strongly oxidizing conditions. According to Baes and Mesmer (1976), Cr(m)
exists predominantly as Ci3+ below pH 3.5 in a Cr(in)-H20 system. With increasing pH, hydrolysis of
Ci3+ yields CrOH2+, Cr(OH)2, Cr(OH);(aq), and Cr(OH);, Cr2(OH)t, and Cr3(OH)^+. At higher
chromium concentrations, polynuclear species, such as Cr2(OH)2+ and Cr3(OH)4+, can form slowly at
25 "C (Baes and Mesmer, 1976). Chromium(VI) hydrolyses extensively, forming primarily anionic
species. These species are HCrO; (bichromate), Cr04" (chromate), and Cr207" (dichromate) (Baes
and Mesmer, 1976; Palmer and Wittbrodt, 1991; Richard and Bourg, 1991). Palmer and Puis (1994)
presented some Cr(VI) speciation diagrams representative of groundwater conditions. They showed
that above pH 6.5, Cr04" generally dominates. Below pH 6.5, HCrO; dominates when the total
concentration of dissolved Cr(VI) is low (<30 mM). When Cr(VI) concentrations are greater than
30 mM, CroOy" is the dominant aqueous species relative to HCrO; at acidic conditions (Palmer and
Puis, 1994). These results are consistent with those of Baes and Mesmer (1976).
5.4.4 Dissolution/Precipitation/Coprecipitation
Several investigators have presented evidence suggesting the formation of solubility-controlling solids of
Cr(m) in soils. Rai and Zachara (1984) concluded that most Cr(m) solubility-controlling solids in
nature are either Cr(OH)3 or Cr(m) coprecipitated with iron oxides. Their conclusion was supported
by 3 observations: (1) the thermodynamic treatment of the data where the solubility of chromite
(FeCr204) is predicted to be the lowest among the chromium minerals for which data are available
(Hem, 1977), (2) the similarity of Cr(m) and Fe(m) ionic radii, and (3) the observations that aqueous
Cr(III) is removed by Fe(OH)3 precipitation and that chromium during weathering is found to associate
with ferric-rich materials (Nakayama et al., 1981). Hem (1977) reported that the total chromium
concentration in groundwater beneath Paradise Valley, Arizona was close to the solubility of Cr203.
Because Cr(III) minerals are sparingly soluble, the aqueous concentration of Cr(III) should be less than
EPA's maximum concentration level (MCL) for chromium (0.1 mg/1) between slightly acid to
moderately alkaline conditions (Palmer and Puis, 1994).
Several Cr(VI)-containing mineral phases may be present at chromium-contaminated sites. Palmer and
Wittbrodt (1990) identified PbCr04 (crocoite), PbCr04 H20 (iranite), and K2Cr04 (tarapacaite) in
5.22
-------
chromium sludge from a plating facility. They also reported that BaCr04 formed a complete solid
solution with BaS04. They concluded that these solid solutions can be a major impediment to the
remediation of chromium-contaminated sites by pump-and-treat technologies.
Chromium(VI) is a strong oxidant and is rapidly reduced in the presence of such common electron
donors as aqueous Fe(II), ferrous iron minerals, reduced sulfur, microbes, and organic matter (Bartlett
and Kimble, 1976; Nakayama et al., 1981). Studies indicate that Cr(VI) can be reduced to Cr(III) by
ferrous iron derived from magnetite (Fe304) and ilmenite (FeTi03) (White and Hochella, 1989),
hematite (Fe^) (Eary and Rai, 1989),1 and pyrite (FeS2) (Blowes and Ptacek, 1992).
The reduction of Cr(VI) by Fe(II) is very rapid. The reaction can go to completion in a matter of
minutes (Eary and Rai, 1989). The rate of reduction of Cr(VI) increases with decreasing pH and
increasing initial Cr(VI) and reductant concentrations (Palmer and Puis, 1994). Interestingly, this
reaction does not appear to be slowed by the presence of dissolved oxygen (Eary and Rai, 1989).
When the pH is greater than 4, Cr(m) can precipitate with Fe(m) to form a solid solution with the
general composition CrxFe1.x(OH)3 (Sass and Rai, 1987). The solubility of chromium in this solid
solution decreases as the mole fraction of Fe(III) increases. The oxidation reaction proceeds much
more slowly than the reduction reaction; the former reaction requires months for completion (Eary and
Rai, 1987; Palmer and Puis, 1994). Only 2 constituents in the environment are known to oxidize
Cr(m): dissolved oxygen and manganese-dioxide minerals [e.g., pyrolusite (P-Mn02)]. Eary and Rai
(1987) reported that the rate of Cr(m) oxidation was much greater in the presence of manganese-
dioxide minerals than dissolved oxygen.
5.4.5 Sorption/Desorption
The extent to which Cr(III) sorbs to soils is appreciably greater than that of Cr(VI) because the former
exists in groundwater as a cation, primarily as Ci3+ (and its complexed species), whereas the latter
exists as an anion, primarily as Cr04" or HCrO;. Most information on Cr(VI) adsorption comes from
studies with pure mineral phases (Davis and Leckie, 1980; Griffin el al., 1977; Leckie el al., 1980).
These studies suggest that Cr(VI) adsorbs strongly to gibbsite (a-Al203) and amorphous iron oxide
[Fe203H20(am)] at low to medium pH values (pH 2 to 7) and adsorbs weakly to silica (Si02) at all
but very low pH values (Davis and Leckie, 1980; Griffin et al., 1977; Leckie el al., 1980). These
results can be explained by considering the isoelectric points (IEP)2 of these minerals. When the pH of
the system is greater than the isoelectric point, the mineral has a net negative charge. When the pH is
1 Eary and Rai (1989) attributed the reduction of Cr(VI) to Cr(III) by hematite (Fe203) as containing
having trace quantities of Fe(II).
2 The isoelectric point (IEP) of a mineral is the pH at which it has a net surface charge of zero. More
precisely, it is the pH at which the particle is electrokinetically uncharged.
5.23
-------
below the isoelectric point, the mineral has a net positive charge. Hence, anion adsorption generally
increases as the pH becomes progressively lower than the isoelectric point. The isoelectric point of
gibbsite (a-Al203) is 9.1, amorphous iron oxide [Fe203 H20 (am)] is 8.1, and silica is 2.0 (Stumm and
Morgan, 1981).
The presence of competing and, less commonly, complexing ions may significantly alter chromate
adsorption. Although sulfate is adsorbed less strongly on Fe^ H20(am) than chromate, sulfate may
compete for adsorption sites when present in higher concentration (Leckie el al., 1980). Phosphate
exhibits a greater competitive effect on chromate adsorption (MacNaughton, 1977), reducing sorption
by around 50 percent when present at equal normality. Information on effects of complexing ions on
Cr(VI) sorption is almost nonexistent, though adsorption of ion pairs [e.g., CaCrO^aq) and
KHCrO^aq)] is suggested as 1 possible mechanism for removal of Cr(VI) by Fe^ H20 (am)
(Leckie et al., 1980).
Adsorption of Cr(m) to soils has received only a nominal amount of research attention. The reason for
this may be that sorption of Cr(m) by soil is commonly ascribed to solid phase formation.
Chromium(m) rapidly hydrolyzes, and precipitates as the hydroxide Cr(OH)3 and/or coprecipitates
with Fe(OH)3 (Artiola and Fuller, 1979; Hem, 1977,). Adsorption may be an especially important
mechanism of sorption at lower pH (pH <4.5) and total chromium concentrations (<10"6 M). Limited
studies infer that Cr(m), like other +3 cationic metals, is strongly and specifically absorbed by soil iron
and manganese oxides (Korte el al., 1976). However, when Cr(III) is present in solution at high
concentrations, it may undergo exchange reactions with aluminosilicates (Griffin el al., 1977).
Chromium(m) adsorption may also be influenced by the presence of manganese-oxide minerals.
Manganese oxides may catalyze oxidation to Cr(VI), thereby decreasing the tendency for chromium to
adsorb to the soils (Bartlett and James, 1979; Nakayama el al., 1981).
5.4.6 Partition Coefficient, Kd, Values
5.4.6.1 General Availability of Kd Data
The review of chromium Kd data obtained for a number of soils (Appendix E) indicated that a number
of factors influence the adsorption behavior of chromium. These factors and their effects on chromium
adsorption on soils were used as the basis for generating a look-up table. These factors are:
• Concentrations of Cr(m) in soil solutions are typically controlled by dissolution/precipitation
reactions.
• Increasing pH decreases adsorption (decrease in Kd) of Cr(VI) on minerals and soils. The
data are quantified for only a limited number of soils.
5.24
-------
• The redox state of the soil affects chromium adsorption. Ferrous iron associated with iron
oxide/hydroxide minerals in soils can reduce Cr(VI) which results in precipitation (higher Kd).
Soils containing Mn oxides oxidize Cr(m) into Cr(VI) form thus resulting in lower Kd values.
The relation between oxide/hydroxide contents of iron and manganese and their effects on Kd
have not been adequately quantified except for a few soils.
• The presence of competing anions reduce Cr(VI) adsorption. These effects have been
quantified as a function of pH for only 2 soils.
The factors which influence chromium adsorption were identified from studies by Leckie el al. (1980),
Davis and Leckie (1980), Griffin el al. (1977), and Rai el al. (1986), and studies discussed below. A
description and assessment of these data are provided in Appendix E.
Adsorption data also show that iron and manganese oxide contents of soils significantly affect the
adsorption of Cr(VI) on soils (Korte el al., 1976). However, these investigators did not publish either
Kd values or any correlative relationships between Kd and the oxide contents. Studies by Stollenwerk
and Grove (1985) and Sheppard et al. (1987) using soils showed that Kd decreases as a function of
increasing equilibrium concentration of Cr(VI). Another study conducted by Rai el al. (1988) on
4 different soils confirmed that Kd values decrease with increasing equilibrium Cr(VI) concentration.
The adsorption data obtained by Rai el al. (1988) also showed that quantities of sodium dithionite-
citrate-bicarbonate (DCB) extractable iron content of soils is a good indicator of a soil's ability to
reduce Cr(VI) to the Cr(III) oxidation state. The reduced Cr has been shown to coprecipitate with
ferric hydroxide. Therefore, observed removal of Cr(VI) from solution when contacted with
chromium-reductive soils may stem from both adsorption and precipitation reactions. Similarly, Rai el
al. (1988) also showed that certain soils containing manganese oxides may oxidize Cr(m) to Cr(VI).
Depending on solution concentrations, the oxidized form (+6) of chromium may also precipitate in the
form of Ba(S,Cr)04 Such complex geochemical behavior chromium in soils implies that depending on
the properties of a soil, the measured Kd values may reflect both adsorption and precipitation reactions.
Adsorption studies have shown that competing anions such as SO4", COfTHCOj, HPO4", H2PO; NOj
and CI", significantly reduce Cr(VI) adsorption on oxide minerals and soils (Leckie el al., 1980;
MacNaughton, 1977; Rai el al., 1986; Rai et al, 1988; Stollenwerk and Grove, 1985).
The data regarding the effects of soil organic matter on Cr(VI) adsorption are rather sparse. In 1 study
(Stollenwerk and Grove, 1985) which evaluated the effects of soil organic matter on adsorption of
Cr(VI), the results indicated that organic matter did not influence Cr(VI) adsorption properties (see
Appendix E).
5.4.6.2 Kd Look-Up Tables
5.25
-------
Among all available data for Cr(VI) adsorption on soils, the most extensive data set was developed by
Rai el al. (1988). These investigators studied the adsorption behavior of 4 different well-characterized
subsurface soil samples. They investigated the adsorption behavior of Cr(VI) on these 4 soil samples as
a function of pH. Additionally, they also investigated the effects of competing anions such as SO4", and
CO3VHCO3. The adsorption data developed by these investigators was used to calculate the Kd
values (Appendix E). These Kd values were used as the basis to develop the look-up Table 5.7.
5.4.6.2.1 Limits of Kd Values with Respect to pH
Natural soil pH typically ranges from about 4 to 11 (Richards, 1954). The 2 most common methods of
measuring soil pH are either using a soil paste or a saturation extract. The standard procedure for
obtaining saturation extracts from soils has been described by Rhoades (1996). The saturation extracts
are obtained by saturating and equilibrating the soil with distilled water followed by collection using
vacuum filtration. Saturation extracts are usually used to determine the pH, the electrical conductivity,
and dissolved salts in soils.
The narrow pH ranges in the look-up table (Table 5.7) were selected from the observed rate of change
of Kd with pH. The Kd values for all 4 soils were observed to decline with increasing pH and at pH
values beyond about 9, Kd values for Cr(VI) are < 1 ml/g (see Appendix E).
5.4.6.2.2 Limits of Kd Values with Respect to Extractable Iron Content
The soil characterization data provided by Rai el al. (1988) indicate the soils with DCB extractable
iron contents above -0.3 mmol/g can reduce Cr(VI) to Cr(III). Therefore the measured Kd values for
such soils reflect both redox-mediated precipitation and adsorption phenomena. The data also show
that soils with DCB extractable iron contents of about 0.25 mmol/g or less do not appear to reduce
Cr(VI). Therefore, 3 ranges of DCB extractable iron contents were selected which represent the
categories of soils that definitely reduce (>0.3 mmol/g), probably reduce (0.26 - 0.29 mmol/g), and do
not reduce (<0.25 mmol/g) Cr(VI) to Cr(m) form.
5.4.6.2.3 Limits of Kd Values with Respect to Competing Anion Concentrations
The adsorption data (Rai el al., 1988) show that when total sulfate concentration in solution is about 2
x 10"3 M (191.5 mg/1), the chromium Kd values are reduced by about an order of magnitude as
compared to a noncompetitive condition. Therefore, a sulfate concentration of about 2 x 10"3 M
(191.5 mg/1) has been used as a limit at which an order of magnitude reduction in Kd values are
expected. Four ranges of soluble sulfate concentrations (0 -1.9, 2 -18.9, 19 - 189, and >190 mg/1)
have been used to develop the look-up table. The soluble sulfate concentrations in soils can be
assessed from saturation extracts (Richards, 1954).
5.26
-------
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5.27
-------
5.5 Lead Geochemistry and Kd Values
5.5.1 Overview: Important Aqueous- and Solid-Phase Parameters Controlling Retardation
Lead has 3 known oxidation states, 0, +2, and +4, and the most common redox state encountered in
the environment is the divalent form. Total dissolved lead concentrations in natural waters are very low
(~10"8 M). Dissolved lead in natural systems may exist in free ionic form and also as hydrolytic and
complex species. Speciation calculations show that at pH values exceeding 7, aqueous lead exists
mainly as carbonate complexes [PbC03(aq), and Pb(C03)2"]. Important factors that control aqueous
speciation of lead include pH, the types and concentrations of complexing ligands and major cationic
constituents, and the magnitude of stability constants for lead-ligand aqueous complexes.
A number of studies and calculations show that under oxidizing conditions depending on pH and ligand
concentrations, pure-phase lead solids, such as PbC03, Pb3(0H)2(C03)2, PbS04, Pb5(P04)3(Cl),
and Pb4S04(C03)2(0H)2, may control aqueous lead concentrations. Under reducing conditions,
galena (PbS) may regulate the concentrations of dissolved lead. It is also possible that lead
concentrations in some natural systems are being controlled by solid solution phases such as barite
(Ba(1_x)PbxS04), apatite [Can.x)Pbx(P04)30H], calcite (Can_x)PbxC03), and iron sulfides (Fe(1_x)PbxS).
Lead is known to adsorb onto soil constituent surfaces such as clay, oxides, hydroxides,
oxyhydroxides, and organic matter. In the absence of a distinct lead solid phase, natural lead
concentrations would be controlled by adsorption/desorption reactions. Adsorption data show that
lead has very strong adsorption affinity for soils as compared to a number of first transition metals.
Lead adsorption studies on bulk soils indicate that the adsorption is strongly correlated with pH and the
CEC values of soils. Properties that affect CEC of soils, such as organic matter content, clay content,
and surface area, have greater affect on lead adsorption than soil pH.
5.5.2 General Geochemistry
Lead is an ubiquitous heavy metal and its concentration in uncontaminated soil ranges from 2 to
200 mg/kg and averages 16 mg/kg (Bowen, 1979). Annual anthropogenic lead input into soils has
been estimated to be from 0.04 to 4 micro g/kg (Ter Haar el al., 1967). In contaminated soils, lead
concentrations may be as high as 18 percent by weight (Mattigod and Page, 1983; Ruby el al., 1994).
Lead in nature occurs in 4 stable isotopic forms (204Pb, 206Pb, 207Pb, and 208Pb). The isotopes, 206Pb,
207Pb, and 208Pb are the stable end products of the 238U, 235U, and 232Th thorium decay series,
respectively (Robbins, 1980). Additionally, heavier isotopes of lead (2"'Pb, 211Pb, 212Pb, and 214Pb)
are known to occur in nature as intermediate products of uranium and thorium decay (Robbins, 1978).
The
5.28
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most common valence state of lead encountered in the environment is the divalent form (Baes and
Mesmer, 1976). Extensive studies of lead biogeochemistry have been conducted due to its known
adverse effects on organisms (Hammond, 1977). Comprehensive descriptions of environmental
chemistry of lead have been published by Boggess and Wixson (1977) and Nriagu (1978).
5.5.3 Aqueous Speciation
Lead exhibits typical amphoteric1 metal ion behavior by forming hydrolytic species (Baes and Mesmer,
1976). Formation of monomelic hydrolytic species, such as PbOH+, Pb(OH)2(aq) and Pb(OH)j, is
well established. Although several polymeric hydrolytic species such as Pb2OH3+, Pb3(OH)|+,
Pb4(OH)4+, and Pb6(OH)g+ are known to form at high lead concentrations, calculations show that these
types of species are unlikely to form at concentrations of dissolved lead (~10"9M) typically
encountered even in contaminated environments (Rickard and Nriagu, 1978). These investigators also
showed that computation models of speciation of dissolved lead in fresh- or seawater predicted that at
pH values exceeding about 6.5, the dominant species are lead-carbonate complexes. Lead is known
to form aqueous complexes with inorganic ligands such as carbonate, chloride, fluoride, nitrate, and
sulfate.
To examine the distribution of dissolved lead species in natural waters, MINTEQA2 model calculations
were completed using the water composition described in Table 5.1. The total lead concentration was
assumed to be 1 micro g/1 based on the data for natural waters tabulated by Duram el al. (1971) and
Hem (1985). A total of 21 aqueous species (uncomplexed Pb2+, and 20 complex species, listed in
Table 5.8) were used in the computation. Results of the computation are plotted as a species
distribution diagram (Figure 5.2). The data show that, under low pH (<6) conditions, free ionic Pb2+
appears to be the dominant species, and the neutral species, PbSO^aq), accounts for about 5 percent
of the total dissolved lead. Within the pH range of 6.5 to 7.5, the main species of lead appear to be
free ionic species, Pb2+, and the neutral complex species, PbCO^aq) with minor percentage of the
species consisting of PbHCC>3 (about 15 percent), PbSO^aq) (<5 percent), and PbOH+ (<5 percent).
Between the pH range 7 to 9, the neutral complex species PbCO^aq) dominates dissolved lead
speciation. At pH values exceeding 9, in addition to PbCO^aq), a significant fraction of soluble lead is
present as the anionic carbonate complex, Pb(C03)2". These calculations also confirm Rickard and
Nriagu's (1978) observation that polymeric species are not significant in the chemistry of lead in natural
waters. The species distribution illustrated in Figure 5.2 does not change if the concentration of total
dissolved lead is increased from 1 to 1,000 micro g/1.
This speciation calculation demonstrates that the important factors that control aqueous speciation of
lead include pH and the types of complexing ligands. Aqueous speciation of lead has a direct bearing
1 Amphoteric behavior is the ability of an aqueous complex or solid material to have a negative, neutral,
or positive charge.
5.29
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on dissolution/precipitation of lead-solid phases and the adsorption/desorption reactions. Complexation
enhances the solubility of lead-bearing solid phases. This enhancement in solubility is dependent on the
strength of complexation [indicated by the magnitude of stability constant] and the total concentrations
of complexing ligands. Also, as will be discussed shortly, adsorption of lead is affected by the type,
charge, and the concentration of lead complexes present in solution. Cationic lead species, especially
Pb2+ and its hydrolysis species, adsorb more commonly than anionic lead complexes.
5.5.4 Dissolution/Precipitation/Coprecipitation
Lead solids in the environment may occur in a number of mineral forms (Rickard and Nriagu 1978;
Mattigod el al., 1986; Zimdahl and Hassett, 1977). However, these authors have identified a limited
number of secondary lead minerals that may control the concentrations of dissolved lead in soil/water
environments. If the concentration of dissolved lead in a pore water or groundwater exceeds the
solubility of any of these phases, the lead-containing solid phase will precipitate and thus control the
maximum concentration of lead that could occur in the aqueous phase. According to Rickard and
Nriagu (1978), under oxidizing conditions, depending on pH and ligand concentrations, cerussite
(PbC03), hydrocerussite [Pb3(0H)2(C03)2], anglesite (PbS04), or chloropyromorphite [Pb5(P04)3Cl]
may control aqueous lead concentrations. A review paper by McLean and Bledsoe (1992) included
data which showed that lead concentrations in a calcareous soil was controlled by lead-phosphate
compounds at lower pH and by mixed mineral phases at pH values exceeding 7.5. A study conducted
by Mattigod et al. (1986) indicated that the mineral leadhillite [Pb4S04(C03)2(0H)2] may be the
solubility controlling solid for lead in a mine-waste contaminated soil.
5.30
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Table 5.8. Lead aqueous species included in the
speciation calculations.
Aqueous Species
Pb2+
PbOFT, Pb(OH)2(aq), Pb(OH)i, Pb(OH)i"
Pb2(OH)+3, Pb3(OH)r
PbCO^aq), Pb(C03)i", PbHCO,
PbSO^aq), Pb(S04f
PbNO^
PbCl+, PbCl^aq), PbCli, PbCg"
PbF+, PbF^aq), PbFi, PbF^"
5.31
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100
¦Q
o>
u
•-
o>
a.
80
60
40
20
PbS04°(aq)
10
PH
Figure 5.2. Calculated distribution of lead aqueous species as a function of pH for the water
composition in Table 5.1. [The species distribution is based on a concentration
of 1 micro g/1 total dissolved lead.]
Lead may also exist in soils as solid-solution phases. Solid solutions are defined as solid phases in
which a minor element will substitute for a major element in the mineral structure. Depending on the
degree of substitution and the overall solubility of the solid-solution phase, the equilibrium solubility of
the minor element in the solid solution phase will be less than the solubility of the solid phase containing
only the minor element (pure phase). For instance, lead may occur as a minor replacement in barite
[Ba(1.x)PbxS04], apatite [Ca(i_x)Pbx(P04)30H], calcite [Can.x)PbxC03], and iron sulfides, [Fe(1_x)PbxS]
(Driesens, 1986; Goldschmidt, 1954; Nriagu and Moore, 1984; Rickard andNriagu, 1978).
Consequently, the equilibrium solubility of lead controlled by these phases will be less than the
concentrations controlled by corresponding pure phases, namely PbS04, Pb5(P04)30H, PbC03, and
PbS, respectively.
5.32
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Under reducing conditions, galena (PbS) may control the lead concentrations in the environment.
Rickard and Nriagu (1978) calculated that, within the pH range of 6-9, the equilibrium solubility of
galena would control total lead concentrations at levels less than approximately 10"10 M (<21 ng/1).
Therefore, if galena is present in a soil under reducing conditions, the aqueous concentrations of lead
will be controlled at extremely low concentrations.
5.5.5 Sorption/Desorption
Lead is known to adsorb onto soil constituent surfaces such as clays, oxides, hydroxides,
oxyhydroxides, and organic matter. Ion exchange reactions of lead on a number of clay minerals such
as montmorillonite, kaolinite, illite, and vermiculite have been studied by a number of investigators.
These studies showed that lead was preferentially adsorbed by exchange on clays, readily replacing
calcium and potassium (Bittel and Miller, 1974; Overstreet and Krishnamurthy, 1950). Studies
conducted by Lagerwerff and Brower (1973) on montmorillonitic, illitic, and kaolinitic soils confirmed
that lead would preferentially exchange for calcium. Another clay mineral, vermiculite, is also known to
exhibit very high ion exchange selectivity for lead (Rickard and Nriagu, 1978). Based on a number of
studies Rickard and Nriagu (1978) also concluded that beyond neutral pH, precipitation reactions may
control lead concentrations in solution rather than ion exchange and adsorption reactions involving clay
mineral surfaces.
Experimental data show that only hydrogen ions and unhydrolyzed aluminum ions are capable of
displacing lead from exchange sites on clay minerals (Lagerwerff and Brower, 1974; Zimdahl and
Hassett, 1977). Clay minerals also differ in their exchange preference for lead. Bittel and Miller
(1974) showed that the exchange preference for lead varies in the sequence,
kaolinite > illite > montmorillonite.
These studies also showed that, in neutral to high pH conditions, lead can preferentially exchange for
calcium, potassium, and cadmium. Under low pH conditions, hydrogen ions and aluminum ions would
displace lead from mineral exchange sites.
Studies of lead adsorption on oxide, hydroxide, and oxyhydroxide minerals show that the substrate
properties, such as the specific surface and degree of crystallinity, control the degree of adsorption
(Rickard and Nriagu, 1978). Experimental data by Forbes el al. (1976) showed that goethite
(FeOOH) has higher adsorption affinity for lead than zinc, cobalt, and cadmium. Data show that
manganese-oxide minerals also adsorb lead ions (Rickard and Nriagu, 1978). These investigators
concluded that the high specificity of lead adsorption on oxide and hydroxide surfaces and the relative
lack of desorbability (<10 percent) of adsorbed lead indicated that lead upon adsorption forms solid
solutions with oxide or hydroxide surfaces. Therefore, this lack of reversibility indicated that the
reaction is not a true adsorption phenomenon.
5.33
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A number of studies have confirmed that many natural and synthetic organic materials adsorb lead.
Data showing significant correlations between concentrations of organic matter and lead in soils indicate
that soil organic matter has a higher affinity for lead adsorption as compared soil minerals.
A number of lead adsorption studies on bulk soils indicate that the adsorption is strongly correlated with
pH and the CEC values of soils (Zimdahl and Hassett, 1977). A multiple regression analysis by
Hassett (1974) of lead adsorption data indicated that properties that affect CEC of soils, such as
organic matter content, clay content, and surface area, have a greater effect on lead adsorption than soil
pH. The results of a number of studies of lead adsorption on a variety of soil and mineral surfaces were
summarized by McLean and Bledsoe (1992). These data show that lead has very strong adsorption
affinity as compared to a number of first row transition metals (cobalt, nickel, copper, and zinc).
According to a recent study (Peters and Shem, 1992), the presence of very strong chelating organic
ligands dissolved in solution will reduce adsorption of lead onto soils. These data show that the
adsorption of lead in the environment is influenced by a number of factors such as the type and
properties of adsorbing substrate, pH, the concentrations of lead, and the type and concentrations of
other competing cations and complex forming inorganic and organic ligands.
5.5.6 Partition Coefficient, Kd, Values
5.5.6.1 General Availability of Kd Data
The review of lead Kd data reported in the literature for a number of soils (Appendix F) led to the
following important conclusions regarding the factors which influence lead adsorption on minerals and
soils.1 These principles were used to evaluate available quantitative data and generate a look-up table.
These conclusions are:
• Lead may precipitate in soils if soluble concentrations exceed about 4 mg/1 at pH 4 and about
0.2 mg/1 at pH 8. In the presence of phosphate and chloride, these solubility limits may be as
low as 0.3 mg/1 at pH 4 and 0.001 mg/1 at pH 8. Therefore, in experiments in which
concentrations of lead exceed these values, the calculated Kd values may reflect precipitation
reactions rather than adsorption reactions.
• Anionic constituents such as phosphate, chloride, and carbonate are known to influence lead
reactions in soils either by precipitation of minerals of limited solubility or by reducing
adsorption through complex formation.
1 Since the completion of our review and analysis of Kd data for the selected contaminants and
radionuclides, the studies by Azizian and Nelson (1998) and Yong and MacDonald (1998) were
identified and may be of interest to the reader.
5.34
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A number of adsorption studies indicate that within the pH range of soils (4 to 11), lead
adsorption increases (as does precipitation) with increasing pH.
Adsorption of lead increases with increasing organic matter content of soils.
Increasing equilibrium solution concentrations correlates with decreasing lead adsorption
(decrease in Kd).
The factors which influence lead adsorption were identified from the following sources of data. A
description and assessment of these data are provided in Appendix F. Lead adsorption behavior on
soils and soil constituents (clays, oxides, hydroxides, oxyhydroxides, and organic matter) has been
studied extensively. However, calculations by Rickard and Nriagu (1978) show that the solution lead
concentrations used in a number of adsorption studies may be high enough to induce precipitation. For
instance, their calculations show that lead may precipitate in soils if soluble concentrations exceed about
4 mg/1 at pH 4 and about 0.2 mg/1 at pH 8. In the presence of phosphate and chloride, these solubility
limits may be as low as 0.3 mg/1 at pH 4 and 0.001 mg/1 at pH 8. Therefore, in experiments in which
concentrations of lead exceed these values, the calculated Kd values may reflect precipitation reactions
rather than adsorption reactions.
Lead adsorption studies on manganese and iron oxides and oxyhydroxides indicate irreversible
adsorption which was attributed to the formation of solid solution phases (i.e., coprecipitation) (Forbes
et al., 1976; Grasselly and Hetenyi, 1971; Rickard and Nriagu, 1978). No correlations however have
been established between the type and content of oxides in soil and the lead adsorption characteristics
of soil.
Anionic constituents such as phosphate, chloride, and carbonate are known to influence lead reactions
in soils either by precipitation of minerals of limited solubility or by reducing adsorption through complex
formation (Rickard and Nriagu, 1978). Presence of synthetic chelating ligands, such as EDTA, has
been shown to reduce lead adsorption on soils (Peters and Shem, 1992). These investigators showed
that the presence of strongly chelating EDTA in concentrations as low as 0.01 M reduced Kd for lead
by about 3 orders of magnitude. By comparison quantitative data is lacking on the effects of more
common inorganic ligands (phosphate, chloride, and carbonate) on lead adsorption on soils.
A number of adsorption studies indicate that within the pH range of soils (4 to 11), lead adsorption
increases with increasing pH (Braids el al., 1972; Bittel and Miller, 1974; Griffin and Shimp, 1976;
Haji-Djafari el al., 1981; Hildebrand and Blum, 1974; Overstreet and Krishamurthy, 1950; Scrudato
and Estes, 1975; Zimdahl and Hassett, 1977). Griffin and Shimp (1976) also noted that clay minerals
adsorbing increasing amounts of lead with increasing pH may also be attributed to the formation of lead
carbonate precipitates which was observed when the solution pH values exceeded 5 or 6.
5.35
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Solid organic matter such as humic material in soils is known to adsorb lead (Rickard and Nriagu,
1978; Zimdahl and Hassett, 1977). Additionally, soluble organic matter such as fulvates and amino
acids are known to chelate soluble lead and affect its adsorption on soils (Rickard and Nriagu, 1978).
Correlative relationships between the organic matter content of soils and its effect on lead adsorption
have been established by Genitse el al. (1982) and Soldatini el al. (1976).
Lead adsorption by a subsurface soil sample from Hanford, Washington was investigated by Rhoads el
al. (1992). Adsorption data from these experiments showed that Kd values increased with decreasing
lead concentrations in solution (from 0.2 mg/1 to 0.0062 mg/1).
5.5.6.2 Kd Look-Up Tables
Among all available data, Genitse et al (1982) obtained adsorption data at lead concentrations (0.0001
- 0.01 mg/1) which apparently precluded precipitation reactions. Also, these concentrations are within
the range of lead concentrations most frequently encountered in ground waters (Chow, 1978).
Additionally, data obtained by Rhoads el al. (1992) indicated that Kd values vary log-linearly as a
function of equilibrium lead concentrations within the range of 0.00001 to 0.2 mg/1. The data generated
by Gerritse el al. (1982) and Rhoads el al. (1992) were used to develop a look-up table (Table 5.9)
of Kd as a function of soil pH and equilibrium lead concentrations.
5.5.6.2.1 Limits of Kd Values with Respect to pH
The pH ranges in the look-up table (Table 5.9) were selected from the rate of change that we noted in
the Kd data as a function of pH. The Kd values within this pH range increase with increasing pH, and
are greatest at the maximum pH limit (pH 11) of soils.
Table 5.9. Estimated range of Kd values for lead as a function of soil pH, and
equilibrium lead concentrations.
Equilibrium Lead
Concentration (micro
g/1)
Krf (ml/g)
Soil pH
4.0 - 6.3
6.4 - 8.7
8.8 - 11.0
0.1 -0.9
Minimum
940
4,360
11,520
Maximum
8,650
23,270
44,580
1.0-9.9
Minimum
420
1,950
5,160
Maximum
4,000
10,760
20,620
10-99.9
Minimum
190
900
2,380
5.36
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Maximum
1,850
4,970
9,530
100 -200
Minimum
150
710
1,880
Maximum
860
2,300
4,410
5.5.6.2.2 Limits of Kd Values with Respect to Equilibrium Lead Concentrations
The limits of equilibrium lead concentrations (0.0001 mg/1 to about 0.2 mg/1) were selected based on
the experimental data generated by Genitse el al. (1982) and Rhoads el al. (1992). These
investigators showed that within the range of initial lead concentrations used in their experiments the
principal lead removal reaction from solution was adsorption and not precipitation. Four concentration
ranges were selected to develop the Kd values.
5.6 Plutonium Geochemistry and Kd Values
5.6.1 Overview: Important Aqueous- and Solid-Phase Parameters
Controlling Retardation
In the ranges of pH and conditions typically encountered in the environment, plutonium can exist in all
4 oxidation states, namely +3, 4, +5, and +6. Under oxidizing conditions, Pu(IV), Pu(V), and Pu(VI)
are common, whereas, under reducing conditions, Pu(m) and Pu(TV) would exist. Dissolved plutonium
forms very strong hydroxy-carbonate mixed ligand complexes, therefore, its adsorption and mobility is
strongly affected by these complex species. Under conditions of low pH and high concentrations of
dissolved organic carbon, it appears that plutonium-organic complexes may be control adsorption and
mobility of plutonium in the environment.
If plutonium is present as a distinct solid phase (amorphous or partly crystalline Pu02xH20) or as a
solid solution, the upper limits of aqueous plutonium concentrations would be in the 10"12 to 10"9 M
range. Dissolved plutonium in the environment is typically present at < 10"15 M levels indicating that
adsorption may be the principal phenomenon that regulates the mobility of this actinide.
Plutonium can adsorb on geologic material from low to extremely high affinities with Kd values ranging
from 11 to 300,000 ml/g. Plutonium in the higher oxidation state adsorbed on iron oxide surfaces may
be reduced to the tetravalent state by Fe(II) present in the iron oxides.
Two factors that influence the mobilization of adsorbed plutonium under environmental pH conditions
(>7) are the concentrations of dissolved carbonate and hydroxyl ions. Both these ligands form very
strong mixed ligand complexes with plutonium, resulting in desorption and increased mobility in the
environment.
5.6.2 General Geochemistry
5.37
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Plutonium is produced by fissioning uranium fuel and is used in the construction of nuclear weapons.
Plutonium has entered the environment either through accidental releases or through disposal of wastes
generated during fuel processing and the production and detonation of nuclear weapons. Plutonium has
15 isotopes, but only 4 of these isotopes namely, 238Pu [tVl (half life) = 86 y], 239Pu {tVl = 24,400 y),
240Pu (tVl = 6,580 y), 241Pu (tVl = 13.2 y), are of environmental concern due to their abundances and
long-half lives.
In the range of pH and redox conditions typically encountered in the environment, plutonium can exist in
4 oxidation states, namely +3, +4, +5, and +6 (Allard and Rydberg, 1983). Plutonium oxidation states
are influenced by factors such as pH, presence of complexants and reductants, radiolysis, and
temperature (Choppin, 1983). Observations indicate that under very low plutonium concentrations and
oxidizing environmental conditions, the disproportionation1 reactions of plutonium are not significant
(Cleveland, 1979). Under reducing conditions, Pu(m) species would be dominant up to pH values
approaching about 8.5, beyond which the Pu(TV) species are known to be the dominant species.
However, under oxidizing conditions and at pH values greater than 4.0, plutonium can exist in +4,+5,
and +6 oxidation states (Keeney-Kennicutt and Morse, 1985). A number of investigators believe that
under oxidizing conditions, the +5 state to be the dominant redox state (Aston, 1980; Bondietti and
Trabalka, 1980; Nelson and Orlandini, 1979; Rai etal., 1980b).
Of the contaminated sites considered in EPA/DOE/NRC (1993), radioactive contamination by 238Pu,
239Pu, and/or 240Pu has been identified at 9 of the 45 Superfund National Priorities List (NPL) sites.
The reported contamination includes airborne particulates, plutonium-containing soils, and plutonium
dissolved in surface- and groundwaters.
5.6.3 Aqueous Speciation
Dissolved plutonium forms complexes with various inorganic ligands such as hydroxyl, carbonate,
nitrate, sulfate, phosphate, chloride, bromide, and fluoride; with many naturally occurring organic
ligands such as acetate, citrate, formate, fulvate, humate, lactate, oxalate, and tartrate; and with
1 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+ + 2H20 = 2Pu3+ + PuOf +4H+
3Pu02 + 4H+ = Pu3+ + 2PuOi +2H20.
5.38
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synthetic organic ligands such as EDTA and 8-hydroxyquinoline derivatives (Cleveland, 1979).
Plutonium(IV) hydrolyzes more readily than all other redox species of plutonium (Baes and Mesmer,
1976). The order of hydrolysis of plutonium redox species follows the sequence
Pu(rv) > Pu(ni) > Pu(vi) > Pu(v)
(Choppin, 1983). Plutonium hydrolytic species may have up to 4 coordinated hydroxyls.
The tendency of plutonium in various oxidation states to form complexes depends on the ionic potential
defined as the ratio (z/r) of the formal charge (z) to the ionic radius (r) of an ion. Among plutonium
redox species, Pu(TV) exhibits the highest ionic potential and therefore forms the strongest complexes
with various ligands. Based on the equilibrium constants (K'.'oyx) for the plutonium complexation
reactions, ligands, such as chloride and nitrate, form weak complexes (log K'^ 298 of 1 to 2) with
plutonium, whereas fluoride, sulfate, phosphate, citrate, and oxalate form stronger complexes
(log K j 298 of 6 to 30). Among the strongest complexes of plutonium are the hydroxy-carbonate mixed
ligand complexes [e.g., Pu(0H)2(C03)2"] (Tait etal., 1995; Yamaguchi etal., 1994). Additionally,
dissolved organic matter (fulvic and humic material) may also form complexes with plutonium. Although
the nature of these complexes and their stability constants have not been fully characterized, it is
believed that humic complexes of plutonium may be the dominant soluble species in natural
environments at lower pH (below 5 to 6) values (Allard and Rydberg, 1983).
Because dissolved plutonium can exist in multiple redox states and form hydrolytic and complex species
in solution, it is useful to assess the probable dominant plutonium aqueous species that may exist in
typical ground water. Therefore, the aqueous speciation of dissolved plutonium was calculated as a
function of pH using the MINTEQA2 code and a concentration of 3.2x10"10 mg/1 (1.36x10"15 M) total
dissolved plutonium. This concentration is based on the maximum activity of 239:240 pu measured by
Simpson et al. (1984) in 33 water samples taken from the highly alkaline Mono Lake in California.
The species distribution was calculated assuming that multiple plutonium valence states might be present
based on thermodynamic equilibrium considerations. This calculation is dependent on redox conditions
as well as the pH and composition of the water. Therefore, a set of oxic conditions that might be
associated with surface or near-surface disposal facilities or contaminated sites were selected for these
illustrative calculations. These redox conditions are based on an experimentally determined pHZEh
relationship described in Lindsay (1979) for suspensions of sandy loam and distilled water. In a series
of acid and base titrations, the pH/Eh response of the soil/water suspension was determined to vary
according to the equation
pe + pH =15.23 (5.1)
where pe = negative log of the electron activity.1
1 The electron activity is defined as unity for the standard hydrogen electrode.
5.39
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The pe is related to Eh by the equation
pe = ( 2.303RT/ F) pe
(5.2)
where R= universal gas constant (1.9872 cal/mol-K)
T = temperature in degrees kelvin
F = Faraday constant (96,487 coulombs/equivalent).
At 25.0 C (298 K),
Eh(mV) = 5.92 pe
(5.3)
Using Equations 5.1 and 5.3, an Eh value was calculated for each pH value used as an input for the
MDS1TEQA2 calculations of plutonium aqueous speciation. The plutonium aqueous species that were
included in the computation scheme are tabulated in Table 5.10. Thermodynamic data for these
species were taken primarily from Lemire and Tremaine (1980) and other secondary sources and
database modifications described by Krupka and Serne (1996).
Results are plotted as a species distribution diagram (Figure 5.3). The data show that, under very low
pH (~3 - 3.5) conditions, PuFf and Pu02 are the dominant species of plutonium. The free ionic
species, Pu02 appears to be the dominant form within the pH range of 4 to 5. Within the pH range of
5.5 to 6.5, the main species of plutonium appear to be PuO 2, and Pu(0H)2(C03)2", with minor species
being the neutral hydrolytic species Pu(OH)4(aq) and the phosphate complex Pu(HP04)4". At pH
values exceeding 6.5, the bulk of the dissolved plutonium (-90 percent) would be comprised of the
Pu(0H)2(C03)2" species with a minor percentage of Pu(OH)4(aq). These illustrative computations
indicate that, under pH conditions that typically exist in surface and groundwaters (>6.5), the dominant
form of dissolved plutonium would be the tetravalent complex species, Pu(0H)2(C03)2".
Polymeric species of plutonium may not occur under environmental conditions because the total
plutonium concentrations in nature are at least 7 orders of magnitude less than the concentrations
required for the formation of such species (Choppin, 1983). It is important to note that the speciation
of plutonium would change significantly with changing redox conditions, pH, the types and total
concentrations of complexing ligands and major cationic constituents.
5.6.4 Dissolution/Precipitation/Coprecipitation
Allard and Rydberg (1983) calculated that the aqueous concentrations of plutonium in nature may be
controlled by the solubility of the solid phase Pu02-xH20. Many observations show that plutonium
associated with soils and particulate organic matter is present in tetravalent oxidation state (Nelson and
Lovett, 1980; Nelson et al., 1987; Silver, 1983). Calculations by Allard and Rydberg (1983) based
on available thermodynamic data show that, under reducing conditions, the solubility of dissolved
5.40
-------
plutonium would be limited by the solid phase Pu02 at pH values greater than 8, and by the solid phase
Pu2(C03)3 of trivalent plutonium at lower pH values.
5.41
-------
Table 5.10. Plutonium aqueous species included in the speciation calculations.
Redox
State
Aqueous Species
Pu(ni)
Pu3+, PuOH2+, Pu(OH)2, Pu(OH);(aq)
PuCOi Pu(C03)2,Pu(C03)r
PuSO+4, Pu(S04)2
PuH2POf, PuCl2+
Pu(IV)
Pu4+, PuOH3+, Pu(OH>f, Pu(OH)i, Pu(OH)4(aq)
Pu(0H)4(C03)t Pu(OH)2(COJ-
PuSOf, Pu(S04)2(aq), PuHPOf, Pu(HP04)2(aq),
Pu(HP04)|", Pu(HP04)4"
PuCP, PuF3+, PuF^+, PuF3, PuF4(aq)
Pu(V)
PuO+2, Pu020H (aq), (Pu02)20H+
Pu(VI)
PuOi+' Pu020H+, Pu02(0H)2(aq),
Pu02(0H)3, (Pu02)2(0H)r, (Pu02)3(0H)+
Pu02C03(aq), Pu02(C03)i", Pu02(C03)i"
Pu02Cl+, Pu02F+, Pu02F2(aq), Pu02Fi, PuO.Fj"
Pu02S04(aq), Pu02H2P04
5.42
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pH
Figure 5.3. Calculated distribution of plutonium aqueous species as a function of pH for the water
composition in Table 5.1. [The species distribution is based on a concentration of
3.2 x 10"10 mg/1 (1.36 x 10"15 M) total dissolved plutonium.]
Laboratory studies conducted by Rai et al. (1980a), Delegard (1987), and Yamaguchi el al. (1994)
indicated that a freshly precipitated amorphous Pu02xH20 phase controls the equilibrium solubility of
plutonium. Solubility on aged precipitates by Rai el al. (1980a) and Delegard (1987) also showed that
equilibrium plutonium concentrations would be controlled by a partially crystallized Pu02xH20 phase
at concentrations about 2 orders of magnitude less than that of amorphous Pu02xH20. Therefore,
under oxidizing conditions, amorphous Pu02xH20, if present in soils, may control soluble plutonium
concentrations near 10"8 M. Under alkaline conditions with high dissolved carbonate concentrations,
dissolved plutonium concentrations may increase to micromolar levels. When dissolved carbonate is
not present, Pu02xH20 may control plutonium concentrations at about 10"10 M (Rai el al., 1980a).
5.43
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5.6.5 Sorption/Desorption
Plutonium is known to adsorb onto soil components such as clays, oxides, hydroxides, oxyhydroxides,
aluminosilicates and organic matter. Depending on the properties of the substrate, pH, and the
composition of solution, plutonium would adsorb with affinities varying from low (Kd = 11 ml/g) to
extremely high (Kd = 300,000 ml/g) (Baes and Sharp, 1983; Coughtrey el al., 1985; Thibault el al.,
1990).
A number of studies indicate that iron hydroxides adsorb and reduce penta- and hexavalent plutonium
to its tetravalent state at the solid surface. Experimental data showed that tetra- and pentavalent
plutonium aqueous species oxidize to hexavalent form upon adsorption onto manganese dioxide
surfaces whereas, pentavalent plutonium adsorbed on goethite disproportionate into tetra and
hexavalent forms (Keeney-Kennicutt and Morse, 1985). Subsequently, the hexavalent form of
plutonium was observed to have been reduced to tetravalent state. Additionally, these reactions were
found to occur faster under light conditions than under dark conditions suggesting photochemical
catalysis of adsorbed plutonium redox change reactions.
Laboratory studies have indicated that increasing carbonate concentrations decreased adsorption of
tetra- and pentavalent plutonium on goethite surfaces (Sanchez el al., 1985). Phenomenon similar to
the reduction and suppression of plutonium adsorption in the presence of carbonate ions have also been
observed for other actinides which also form strong hydroxy-carbonate mixed ligand aqueous species.
These data suggest that plutonium would be most mobile in high pH carbonate-rich groundwaters.
Some studies indicate that the mass of plutonium retarded by soil may not be easily desorbed from soil
mineral components. For example, Bunzl el al. (1995) studied the association of 239+240Pu from global
fallout with various soil components. They determined the fractions of plutonium present as readily
exchangeable, bound to carbonates, bound to iron and manganese oxides, bound to organic matter,
and residual minerals. For soils at their study site in Germany, the results indicated that 30-40 y after
deposition of the plutonium, the readily exchangeable fraction of plutonium was less than 1 percent.
More than 57 percent of the plutonium was sorbed to organic matter and a considerable mass sorbed
to the oxide and mineral fractions.
5.44
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5.6.6 Partition Coefficient, Kd, Values
5.6.6.1 General Availability of Kd Data
A number of studies have focused on the adsorption behavior of plutonium on minerals, soils, and other
geological materials.1 A review of data from diverse sources of literature indicated that Kd values for
plutonium typically range over 4 orders of magnitude (Thibault el al., 1990). Also, based on a review
of these data, a number of factors which influence the adsorption behavior of plutonium have been
identified. These factors and their effects on plutonium adsorption on soils were used as the basis for
generating a look-up table. These factors are:
Typically, in many experiments, the oxidation state of plutonium in solution was not determined
or controlled. Therefore it would be inappropriate to compare the Kd data obtained from
different investigations.
In natural systems with organic carbon concentrations exceeding -10 mg/kg, plutonium exists
mainly in trivalent and tetravalent redox states. If initial plutonium concentrations exceed ~10"7
M, the measured Kd values would reflect mainly precipitation reactions and not adsorption
reactions.
Adsorption data show that the presence of ligands influence plutonium adsorption onto soils.
Increasing concentrations of ligands decrease plutonium adsorption.
If no complexing ligands are present plutonium adsorption increases with increasing pH
(between 5.5 and 9.0).
Plutonium is known to adsorb onto soil components such as aluminum and iron oxides,
hydroxides, oxyhydroxides, and clay minerals. However, the relationship between the amounts
of these components in soils and the measured adsorption of plutonium has not been quantified.
The factors which influence plutonium adsorption were identified from the following sources of data. A
description and assessment of these data are provided in Appendix G. Because plutonium in nature
can exist in multiple oxidation states (m, IV, V, and VI), soil redox potential would influence the Pu
redox state and its adsorption on soils. However, our literature review found no plutonium adsorption
studies which included soil redox potential as a variable. Studies conducted by Nelson el al. (1987)
and Choppin and Morse (1987) indicated that the oxidation state of dissolved plutonium under natural
1 Since the completion of our review and analysis of Kd data for the selected contaminants and
radionuclides, the studies by Duff el al. (1999) and Fisher el al. (1999) were identified and may be of
interest to the reader.
5.45
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conditions depended on the colloidal organic carbon content in the system. Additionally, Nelson et al
(1987) also showed that plutonium precipitation occurred if the solution concentration exceeded 10"7
M.
Plutonium complexation by ligands, such as acetate (Nishita, 1978; Rhodes, 1957), oxalate (Bensen,
1960), and fulvate (Bondietti el al., 1975), are known to reduce adsorption of plutonium. Studies of
suspended particles from natural water systems also showed that increasing concentrations of dissolved
organic carbon decreased plutonium adsorption (Nelson el al., 1987). Experiments using synthetic
ligands such as EDTA (1 mmol/1), DTPA (1 mmol/1), and HEDTA (100 mmol/1) have shown that
plutonium adsorption onto soils was reduced due to complexing effects of these ligands (Delegard el
al., 1984; Relyea and Brown, 1978). However, it is unlikely that such concentrations of these synthetic
ligands would exist in soils. The effects of carbonate ions on Pu(TV) adsorption on goethite have been
quantified by Sanchez el al. (1985). They found that carbonate concentrations exceeding 100 mmol/1
significantly reduced adsorption of Pu(TV) on goethite. In contrast, under soil saturation extract
conditions in which carbonate concentrations typically range from 0.1 to 6 mmol/1 HC03", Pu(IV)
adsorption appears to increase with increasing carbonate concentration (Glover el al., 1976).
Rhodes (1957) and Prout (1958) conducted studies of plutonium adsorption as a function of pH. Both
these studies indicated that Pu exhibited an adsorption maxima between pH values 6.5 to 8.5. These
data however are unreliable because initial plutonium concentrations of 6.8xl0"7 to lxlO"6 M used in
the experiments may have resulted in precipitation reactions thus confounding the observations.
Even though the adsorption behavior of plutonium on soil minerals such as glauconite (Evans, 1956),
montmorillonite (Billon, 1982; Bondietti el al., 1975), attapulgite (Billon, 1982), and oxides,
hydroxides, and oxyhydroxides (Evans, 1956; Charyulu el al., 1991; Sanchez el al., 1985; Tamura,
1972; Tick nor, 1993; Van Dalen el al., 1975) has been studied, correlative relationships between the
type and quantities of soil minerals in soils and the overall plutonium adsorption behavior of the soils
have not been established.
Plutonium adsorption data for 14 soils have been collected by Glover el al. (1976) along with a
number of soil properties that included soil organic matter content. A multiple regression analyses of
these data showed that compared to other soil parameters such as clay mineral content, dissolved
carbonate concentration, electrical conductivity and pH, soil organic matter was not a significant
variable.
These criteria were used to evaluate and select plutonium adsorption data in developing a look-up
table. Only 2 adsorption studies using soils in which the initial concentrations of Pu(TV) used were less
than the concentration that would trigger precipitation reactions. Barney (1984) conducted adsorption
experiments in which initial plutonium concentrations of 10"11 to 10"9 M were used to examine plutonium
adsorption on to basalt interbed sediments from Hanford, Washington. Glover el al. (1976) conducted
a set of experiments using 10"8 M initial concentration to study the adsorption behavior of Pu(TV) on
5.46
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14 different soil samples from 7 DOE sites. A number of soil properties were also measured thus
providing a basis to correlate the adsorption behavior with a number of soil parameters. This is the best
available data set for Pu(IV) adsorption on a number of well characterized soils therefore, it was used
to develop correlative relationships and a look-up table for Kd values.
5.6.6.2 Kd Look- Up Table
The look-up table for plutonium Kd values (Table 5.11) was generated using the a piece-wise
regression model with clay content and dissolved carbonate as the independent variables (See
Appendix G for details).
5.6.6.2.1 Limits of Kd Values with Respect to Clay Content
The clay contents of the soils used for developing the regression relationship ranged from 3 to 64
percent by weight. Therefore the range of clay contents for the look-up table was set between 0 and
70 percent. Extending the regression relationship for high clay soils (>70 percent) would result in a
higher degree of uncertainty for predicted Kd values. Clay contents of soils are typically measured as
part of textural analysis of soil. Clay content of a soil is defined as the mass of soil particles with
average particle size of < 2 micro m.
Table 5.11. Estimated range of Kd values for plutonium as a function of the soluble
carbonate and soil clay content values.
Kd(ml/g)
Clay Content (wt.%)
0-30
31-50
51-70
Soluble Carbonate
(meq/1)
Soluble Carbonate
(meq/1)
Soluble Carbonate
(meq/1)
0.1-2
3-4
5-6
0.1-2
3-4
5-6
0.1 -2
3-4
5-6
Minimum
5
80
130
380
1,440
2,010
620
1,860
2,440
Maximum
420
470
520
1,560
2,130
2,700
1,980
2,550
3,130
5.47
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5.6.6.2.2 Limits of Kd Values with Respect to Dissolved Carbonate Concentrations
The dissolved carbonate content of the soils used for the regression relationships ranged from about 0.1
to 6 meq/1 (0.1 to 6 mmol/1 of HC03"). The dissolved carbonate values were measured on saturation
extracts obtained from these soils. The standard procedure for obtaining saturation extracts from soils
has been described by Rhoades (1996). The saturation extracts are obtained by saturating and
equilibrating the soil with distilled water followed by vacuum filtration to collect the extract. Saturation
extracts are usually used to determine the pH, the electrical conductivity, and dissolved salts in soils.
For soils with pH values less than 8.5, the saturation extracts typically contain less than 8 mmol/1 of
dissolved carbonate (Richards, 1954).
The regression relationship indicates that within the range of 0.1 to 6 mmol/1 of dissolved carbonate, the
Kd values increase with increasing dissolved carbonate values. Adsorption experiments conducted by
Sanchez el al. (1985) showed however that very high concentrations (100 to 1,000 meq/1) of dissolved
carbonate in matrix solution decreases Pu adsorption on goethite. The dissolved carbonates in soil
saturation extracts are 3 to 4 orders of magnitude less than the concentrations used in experiments by
Sanchez el al. (1985). The data by Glover el al. (1976) show that within very low concentration
range of dissolved carbonate (0.1 to 6 mmol/1) found soil saturation extracts, Kd values for Pu increase
as a function of dissolved carbonate. This correlation may be strictly serendipitous and a more likely
variable that would lead to an increased Kd would be increasing pH.
5.7 Radon Geochemistry and Kd Values
5.7.1 Overview: Important Aqueous-and Solid-Phase Parameters
Controlling Retardation
The migration of radon, an inert gas, in soil/water systems is not affected itself by aqueous speciation,
precipitation/dissolution, or adsorption/desorption processes. Therefore, the mobility of radon is not
affected by issues associated with the selection of appropriate "adsorption" Kd values for modeling
contaminant transport and risks in soil /water systems. Radon is soluble in water, and the hydrostatic
pressure on ground water below the water table is sufficient to keep dissolved radon in solution.
The generation of radon is however affected by the concentrations of its parent elements which, along
with radon's decay products, are of regulatory concern. Because aqueous speciation,
precipitation/dissolution, or adsorption/desorption processes can affect the movement of radon's
parents and decay products in soils, these processes should be considered when modeling contaminant
transport in a total environmental system, including air transport pathways.
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5.7.2 General Geochemistry
Radon is a colorless, odorless, essentially inert gas. All radon isotopes are radioactive. The longest-
lived isotope of radon is 222Rn which has a half life (tVl) of 3.8 d. The main health risk is from inhalation
of radon gas and its daughter products which are usually adsorbed on dust in the air. Detailed
descriptions of the geologic controls, migration, and detection of radon have been included in published
proceedings such as Graves (1987), Gesell and Lowder (1980), and elsewhere. Of the 45 Superfimd
National Priorities List (NPL) sites considered in EPA/DOE/NRC (1993), radioactive contamination of
air, soil, surface water, and/or groundwater by 220Rn and/or 222Rn has been identified at 23 sites.
Twenty isotopes of radon are known (Weast and Astle, 1980). Environmental radon contamination
typically results from radioactive decay of isotopes in the uranium-thorium series. These include the
formation of:
• 222Rn by alpha decay from 226Ra in the 238U decay series
• 220Rn (tv= 54 sec) by alpha decay from 224Ra in the 232Th decay series
• 219Rn (I, =3.9 sec) by alpha decay from 223Ra in the 238U decay series.
The final, stable daughter products in these 3 decay series are 206Pb, 208Pb, and 207Pb, respectively.
Some noble gases (i.e., krypton, xenon, and radon) have very limited chemical reactivity with other
elements. The chemical reactivity of radon is difficult to assess because of its short half life.
Geologic and hydrogeologic processes that might influence radon mobility are discussed in detail by
Tanner (1980). As an inert gas, radon is not immobilized by precipitation processes along migration
pathways. According to data cited by Tanner (1980), the ratio (i.e., solubility distribution coefficient)
of 222Rn in a water phase to that in a gas phase ranges from 0.52 at 0"C to 0.16 at 40" C. This ratio
has been used, for example, for the solubility of radon in water in mathematical models designed to
calculate radon diffusion coefficients in soils (e.g., Nielson el al., 1984). The solubility of radon in
organic liquids is greater than that in water.
5.7.3 Aqueous Speciation
The existence of radon aqueous species was not identified in any of the references reviewed for this
study. Given the inertness of radon and the short half life (/, =3.8 d) for 222Rn, aqueous speciation and
complexation of dissolved radon would not be expected to be important.
However, as noted above, radon is soluble in water. The hydrostatic pressure on ground water below
the water table is sufficient to keep dissolved radon in solution. Above the water table, the radon
5.49
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present in vadose zone pore water will exsolve from solution, enter the vapor phase, and migrate as
part of the air through the open rock and soil pore spaces.
5.7.4 Dissolution/Precipitation/Coprecipitation
Because radon exists as a dissolved gas, dissolution/precipitation processes are not important relative
to the geochemical behavior of radon and its movement through aqueous environments. These
processes are, however, important relative to the geochemical behavior of radon's parent elements
(e.g., radium) and associated mechanisms by which the radon gas escapes from the solid phases into
ground- and soil waters.
Rama and Moore (1984) studied the mechanism for the release of 222Rn and 220Rn from solid aquifer
material. They determined that radon and other decay products from the U-Th series were released by
alpha recoil1 from the walls of nanometer-size pores in the aquifer solids. Radon diffused into the
intergranular water for release to the atmosphere or decay to more long-lived products. These decay
products may in turn diffuse from the intergranular water and become adsorbed onto the walls of the
nanometer-size pores.
5.7.5 Adsorption/Desorption
Adsorption processes are not expected to be important relative to the geochemical behavior of
gaseous radon and its movement through aqueous environments. The lack of importance of sorption
processes is also supported by studies conducted at cryogenic temperatures (Tanner, 1980).
However, as noted by Tanner (1980), "adsorption effects on the release of radon isotopes from
geologic materials have not been studied sufficiently to determine unambiguously whether they are an
important factor."
5.7.6 Partition Coefficient, Kd, Values
Because adsorption processes are not important relative to the movement of gaseous radon through
aqueous environments, a review of Kd values for radon was not conducted. Compilations, such as
Thibault el al. (1990), do not list any Kd values for radon. A Kd value of zero should be considered for
radon.
1 Alpha recoil refers to the 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).
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5.8 Strontium Geochemistry and Kd Values
5.8.1 Overview: Important Aqueous-and Solid-Phase Parameters
Controlling Retardation
Strontium in solution is expected to be predominantly present as the uncomplexed Sr2+ ion. Only in
highly alkaline soils could strontianite (SrC03) control strontium concentrations in solutions. The extent
to which strontium partitions from the aqueous phase to the solid phase is expected to be controlled
primarily by the CEC of the solid phase. In environments with a pH greater than 9 and dominated by
carbonates, coprecipitation with CaC03 and/or precipitation as SrC03 may become an increasingly
important mechanism controlling strontium removal from solution (Lefevre et al., 1993). A direct
correlation between solution pH and strontium Kd has been reported (Prout, 1958; Rhodes, 1957).
This trend is likely the result of hydrogen ions competing with Sr2+ for exchange sites and the result of
pH increasing the CEC. Strontium Kd values may decrease from 100 to 200 ml/g in low ionic strength
solutions to less than 5 ml/g in high ionic strength solutions (Routson el al., 1980). Calcium is an
important competing cation affecting 90Sr Kd values (Kokotov and Popova, 1962; Schulz, 1965). The
most important ancillary parameters affecting strontium Kd values are CEC, pH, and concentrations of
calcium and stable strontium.
5.8.2 General Geochemistry
Strontium exists in nature only in the +2 oxidation state. The ionic radius of Sr2 is 1.12 A, very
close to that of Ca2+ at 0.99 A (Faure and Powell, 1972). As such, strontium can behave chemically
as a calcium analog, substituting for calcium in the structure of a number of minerals. Strontium has
4 naturally occurring isotopes: 84Sr (0.55 percent), 86Sr (9.75 percent), 87Sr (6.96 percent), and 88Sr
(82.74 percent). The other radioisotopes of strontium are between 80Sr and 95Sr. Only 90Sr [half life
(/.J = 28.1 y], a fission product, is of concern in waste disposal operations and environmental
contamination. The radionuclide 89Sr also is obtained in high yield, but the half-life is too short (ty2
= 52 d) to create a persistent environmental or disposal problem. Because of atmospheric testing of
nuclear weapons, 90Sr is distributed widely in nature. The average 90Sr activity in soils in the United
States is approximately 100 mCi/mi2 As a calcium analog, 90Sr tends to accumulate in bone
(UNSCEAR, 1982).
Contamination includes airborne particulates, strontium-containing soils and strontium dissolved in
surface- and groundwaters. Of the contaminated sites considered in EPA/DOE/NRC (1993),
radioactive contamination by 90Sr has been identified at 11 of the 45 Superfund National Priorities List
(NPL).
5.8.3 Aqueous Speciation
5.51
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There is little tendency for strontium to form complexes with inorganic ligands (Faure and Powell,
1972). The solubility of the free Sr2+ ion is not greatly affected by the presence of most inorganic
anions. Dissolved strontium forms only weak aqueous complexes with carbonate, sulfate, chloride, and
nitrate. For example, Izrael and Rovinskii (1970) used electrodialysis to study the chemical state of
strontium leached by groundwater from rubble produced in a nuclear explosion. They found that 100
percent of the strontium existed as uncomplexed Sr2+, with no colloidal or anionic strontium present in
the leachate. Stevenson and Fitch (1986) concluded that strontium should not form strong complexes
with fulvic or humic acids based on the assumptions that strontium would exhibit similar stability with
organic ligands as calcium and that strontium could not effectively compete with calcium for exchange
sites because calcium would be present at much greater concentrations. Thus, organic and inorganic
complexation is not likely to greatly affect strontium speciation in natural groundwaters.
Species distribution of strontium was calculated using the water composition described in Table 5.1 and
a concentration of 0.11 mg/1 total dissolved strontium. Hem (1985, p. 135) lists this value as a median
concentration of dissolved strontium for larger United States public water supplies based on analyses
from Skougstad and Horr (1963). The strontium aqueous species included in the speciation
calculations are listed in Table 5.12. These MINTEQA2 calculations support the contention that
strontium will exist in groundwaters predominantly as the uncomplexed Sr2 ion. The Sr2 ion
dominates the strontium speciation throughout the pH range of 3 to 10. Between pH 3 and 8.5, the
Sr2+ species constitutes approximately 98 percent of the total dissolved strontium. The remaining 2
percent is composed of the neutral species SrSO^aq). Between pH 9 and 10, SrC03(aq) is
calculated to be between 2 and 12 percent of the total dissolved strontium. As the pH increases above
9, the SrCO'^aq) complex becomes increasingly important. The species distribution for strontium
does not change if the concentration of total dissolved cadmium is increased from 1 to 1,000 micro g/1.
5.8.4 Dissolution/Precipitation/Coprecipitation
Strontium is an alkaline-earth element, which also includes beryllium, magnesium, calcium, strontium,
barium and radium, and can form similar solid phases as calcium. For instance, the 2 most prevalent
strontium minerals, celestite (SrS04) and strontianite (SrC03), have calcium counterparts, anhydrite
(CaS04), and calcite (CaC03). In an acidic environment, most of the strontium solids will be highly
soluble, and, if the activity of Sr2 in solution exceeds approximately 10"4 mol/1, celestite may precipitate
to form a stable phase. However, in alkaline conditions, strontianite would be the stable solid phase
and could control strontium concentrations in soil solutions. However, the dissolved strontium
concentrations in most natural waters are generally well below the solubility limit of strontium-containing
minerals.
5.52
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Table 5.12
Strontium aqueous species included in the speciation calculations.
Aqueous Species
Sr*, SrOH+
SrCO^aq), SrSO^aq), SrNO+3
SrCr, SrF+
SrPOi, SrHPOl(aq), SrH2P04, SrP2Of
Because strontium generally exists in nature at much lower concentration than calcium, it commonly
does not form pure phases (Faure and Powell, 1972). Instead it forms coprecipitates (solid solutions)
with calcite and anhydrite. Calcite can allow the substitution of several hundred parts per million
strontium before there is any tendency for strontianite to form. Strontium can also coprecipitate with
barium to form (Ba(1.x),Srx)S04 in more-alkaline environments (Ainsworth and Rai, 1987; Felmy el al.,
1993).
5.8.5 Adsorption/Desorption
A great deal of research has been directed at understanding and measuring the extent to which
strontium adsorbs to soils [reviewed by Ames and Rai (1978) and Strenge and Peterson (1989)]. The
primary motivation for this research is the need to understand the environmental fate and mobility of
90Sr, particularly as it relates to site remediation and risk assessment. The mechanism by which
5.53
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strontium partitions from the dissolved phase to the solid phase at pH values less than 9 is commonly
believed to be cation exchange1 (Ames and Rai, 1978; Lefevre et al., 1993; McHenry, 1958).
Among the most important environmental parameters affecting the magnitude of a strontium Kd value is
the soil CEC (Ames and Rai, 1978; Lefevre et al., 1993; McHenry, 1958). This finding is consistent
with cation exchange proposed as the mechanism generally controlling strontium adsorption. The
results of Serne and LeGore (1996) also indicate that strontium adsorption is largely controlled by
cation exchange. They reported that 90Sr adsorption was reversible; that is, strontium could be easily
desorbed (exchanged) from the surfaces of soils. Natural soils that had been in contact with 90Sr for
approximate 27 y could be leached of adsorbed90 Sr as readily as similar soils containing recently
adsorbed strontium, indicating that90 Sr does not become more recalcitrant to leaching with time.
Furthermore, these studies suggested that cation exchange, and not (co)precipitation, was responsible
for90 Sr sorption because the latter would leach at a much slower rate.
Some studies indicate that a fraction of some 90Sr sorbed to soil components may not be readily
exchanged [see review in Brady et al. (1999)]. For example, Schulz and Riedel (1961) studied the
influence of aging on the sorption of carrier-free 90Sr into nonexchangeable forms by three soils. They
observed that less than 10% of the total applied carrier-free90 Sr was not easily exchanged which they
attributed to adsorption onto solid-phase carbonates or phosphates. A study by Wiklander (1964)
indicated that after 4 y, only 90 percent of the 90Sr added to the soil could be displaced by repeated
acidic ammonium acetate (pH 4.6) extractions. Wiklander proposed that the retention of 90Sr was due
to strontium substituting for calcium into or adsorbing onto calcium-bearing minerals. Studies by
Roberts and Menzel (1961) and Taylor (1968) showed that as much as 50% of the 90Sr in some acidic
soils was not readily exchangeable. In sediments sampled from the White Oak Creek watershed at
DOE's Oak Ridge Site, Cerling and Spalding (1982) determined that the majority of the 90Sr present in
the sediments was weakly adsorbed and exchangeable, but substantial mass was fixed in the sediments.
They found that approximately 80-90 percent of90 Sr present in these sediments was extracted by
warm 1 A' NaCl or NH4OAC solutions and quantitative extraction required hot 8 N nitric acid.
Some important ancillary soil properties include the natural strontium and calcium concentrations in the
aqueous and solid phases (Kokotov and Popova, 1962; Schulz, 1965), mineralogy (Ames and Rai,
1 Cation exchange is a 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) + 90Sr2+(aq) = 90SrX(s) + Ca2+ (aq)
where X designates an exchange surface site. Adsorption phenomena are discussed in more detail in
Volume I of this report.
5.54
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1978), pH (Juo and Barber, 1970; Prout, 1958; Rhodes, 1957), and solution ionic strength (Rhodes,
1957; Routson el al., 1980). Numerous studies have been conducted to elucidate the effects of
competing cations on strontium adsorption [reviewed by Ames and Rai (1978) and Strenge and
Peterson (1989)]. These experiments consistently show that, on an equivalence basis, strontium will
dominate most Group 1A and IB elements (alkaline and alkaline earth elements) in competition for
exchange sites.
A ranking of the most common groundwater cations by their ability to displace strontium from an
exchange site is:
Stable Sr > Ca > Mg> K > NH4 > Na (5.4)
(Kokotov and Popova, 1962). Calcium exists in groundwaters at concentrations typically 2 orders of
magnitude greater than stable strontium and typically more than 12 orders of magnitude greater than
90Sr (Table 5.1). Consequently, mass action would improve the likelihood of calcium out competing
90Sr for exchange sites.
Rhodes (1957) showed the effect of solution pH and ionic strength on the adsorption of strontium on
soils containing carbonate minerals and montmorillonite. The pH of the system was adjusted with
NaOH or HC1 and the ionic strength was adjusted by adding 4 M NaN03. For a dilute solution, the
strontium Kd increased from 5 ml/g at pH 6 to 10 ml/g at pH 8, and 120 ml/g at pH 10. Above pH 10,
strontium adsorption began to level off, and the sodium added in the NaOH used for pH adjustment
began to compete for exchange sites with the strontium. In 4 M NaN03 (an extremely high ionic
strength solution with respect to natural environments), strontium adsorption was much less affected by
pH. At pH 8, for example, the strontium Kd was about 5 ml/g and increased to about 10 ml/g at pH
10. Using kaolinitic soils from South Carolina, Prout (1958) reported very similar pH and ionic
strength effects as Rhodes (1957). A maximum strontium adsorption was reached at about pH 10,
although this maximum was much higher (Kd = 700 to 800 ml/g) than that reported by Rhodes (1957).
Prout (1958) also reported only a slight pH effect on strontium Kd values in high ionic strength
solutions. Rhodes (1957) and Prout (1958) reported that increases in ionic strength resulted in lower
strontium Kd values.
5.8.6 Partition Coefficient, Kd, Values
5.8.6.1 General Availability of Kd Data
Two simplifying assumptions underlying the selection of strontium Kd values included in the look-up
table were made. Strontium adsorption: (1) occurs by cation exchange, and (2) follows a linear
isotherm. These assumptions appear to be reasonable for a wide range of environmental conditions.
However, these simplifying assumptions are compromised in systems with strontium concentration
5.55
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greater than about 10"4 M, humic substance concentration greater than about 5 mg/1, ionic strength
levels greater than about 0.1 M, and pH levels greater than about 12.
Based on these assumptions and limitation, strontium Kd values and some important ancillary
parameters that influence cation exchange were collected from the literature and tabulated
(Appendix H).1 Data included in this table, were from studies that reported Kd values (not percent
adsorbed or Freundlich or Langmuir constants) and were conducted in systems consisting of (1) natural
soils (as opposed to pure mineral phases), (2) low ionic strength (<0.1 M), (3) pH values between 4
and 10, (4) strontium concentrations less than 10"4 M, (5) low humic material concentrations (<5 mg/1),
and (6) no organic chelates (e.g., as EDTA). Initially, attempts were made to include in the Kd data set
all the key aqueous- and solid-phase parameters identified above. These parameters included CEC,
pH, calcium concentration, stable strontium concentration, and carbonate concentration.
The ancillary parameters for which data could be found that was included in these tables were clay
content, pH, CEC, surface area, solution calcium concentrations, and solution strontium concentrations.
This table described 63 strontium Kd values. A second table containing strontium Kd values for soils as
well as pure mineral phases was prepared at the same time and this table contained 166 entries. These
data are included in Appendix H but were not used to provide guidance regarding the selection of Kd
values to be included in the look-up table.
5.8.6.2 Look-Up Table
The look-up table requires knowledge of the CEC (or clay content) and pH of the system in order to
select the appropriate strontium Kd value (Table 5.13). A detailed explanation of the approach used in
selecting these Kd values is presented in Appendix H. Briefly, it involves tabulating the Kd and ancillary
data found in the literature and then conducting regression analysis of the data with strontium Kd as the
dependent variable. Selection of independent variables used in the final look-up tables was based in
part on their correlation coefficients. Perhaps more importantly, the independent variables had to be a
parameter that is readily available to modelers. For instance, particle size and pH are often available to
modelers whereas such parameters as iron oxide or surface area are not as frequently available. The
estimated ranges for the minimum and maximum Kd values were based on regression estimates of the
95 percent error (P < 0.05). The central estimates were based primarily on values calculated using the
appropriate regression equations.
5.8.6.2.1 Limits of Kd Values with Respect to pH, CEC and Clay Content Values
1 Since the completion of our review and analysis of Kd data for the selected contaminants and
radionuclides, the studies by Chen et al. (1998), Fisher el al. (1999), Oscarson and Hume (1998), and
Wang el al. (1998) were identified and may be of interest to the reader.
5.56
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A full factorial table was created that included 3 pH categories and 3 CEC categories, resulting in
9 cells (Table 5.13). Each cell contains an estimated minimum and maximum Kd value. As the pH or
the CEC of a system increases, so does the strontium Kd values.
A second table was created based on Table 5.13, in which clay content replaced CEC as an
independent variable (subset of Table 5.13). This second table was created because it is likely that
clay content data will be more readily available for modelers than CEC data. To accomplish this, clay
contents associated with the CEC values used to delineate the different categories were calculated using
regression equations (see Appendix H). for additional details).
5.8.6.2.2 Limits of Kd Values with Respect to Dissolved Calcium Concentrations
Of the 63 experiments reporting strontium Kd values, 32 also reported dissolved calcium concentrations
(Appendix H). The mean calcium concentration in this data set was 56 mg/1, with a minimum of 0 mg/1
and a maximum of 400 mg/1. Calcium concentration had a correlation with strontium Kd values, r = -
0.17. Although this correlation is insignificant, it does show that the relationship between these
2 parameters is negative. This inverse relationship can be attributed to calcium competing with
strontium for adsorption sites on the solid phase.
5.57
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Table 5.13. Look-up table for estimated range of Kd values for strontium based on CEC
(meq/100 g), clay content (wt.%), and pH. [Tabulated values pertain to systems
consisting of natural soils (as opposed to pure mineral phases), low ionic strength
(< 0.1 M), low humic material concentrations (<5 mg/1), no organic chelates (e.g.,
EDTA), and oxidizing conditions.]
Kd (ml/g)
CEC (meq/100 g) / Clay Content (wt.%)
3 / < 4
3-10/4-20
10 - 50 / 20 - 60
pH
pH
pH
<5
5-8
8-10
<5
5-8
8-10
<5
5-8
8-10
Minimum
1
2
3
10
15
20
100
200
300
Maximum
40
60
120
150
200
300
1,500
1,600
1,700
5.8.6.2.3 Limits of Kd Values with Respect to Dissolved Stable Strontium and Carbonate
Concentrations
Of the 63 experiments reporting strontium Kd values, none reported stable strontium or carbonate
concentrations (Appendix H). It was anticipated that the presence of stable strontium would compete
with the 90Sr for exchange sites, thereby decreasing 90Sr Kd values. The presence of dissolved
carbonate would likely decrease90 Sr Kd values due to formation of the weaker strontium-carbonate
aqueous complex.
5.9 Thorium Geochemistry and Kd Values
5.9.1 Overview: Important Aqueous-and Solid-Phase Parameters
Controlling Retardation
Thorium occurs only in the +4 oxidation state in nature. In aqueous solutions, especially in natural
waters, the concentrations of dissolved thorium are very low. Dissolved thorium forms a variety of
hydroxyl species, and undergoes extensive chemical interaction with water and most anions. Thorium
can form various aqueous complexes with inorganic anions such as dissolved carbonate, fluoride,
phosphate, chloride, and nitrate. The formation of these complexes will increase the concentrations of
total dissolved thorium in soil- and groundwaters. Recent studies of carbonate complexation of
dissolved thorium indicate that the speciation of dissolved thorium may be dominated by mixed thorium
5.58
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carbonate and hydroxyl-carbonate complexes, such as Th(OH)3COi, at pH values greater than 7.5.
Species distributions calculated using the stability constants for thorium citrate, oxalate, and
ethylenediamine complexes indicate that thorium organic complexes likely predominate over inorganic
complexes in organic-rich waters and soils. This would have an important effect on the solubility and
adsorption of thorium in such waters.
Thorium-containing minerals, such as thorite, thorianite, monazite, and zircon, do not dissolve readily in
low-temperature surface- and groundwaters. Because these minerals form at temperature and pressure
conditions associated with igneous and metamorphic rocks, it is unlikely that the concentration of
thorium in soil/water environments is controlled by the solubility of any of these minerals. The rate at
which thorium is released to the environment may however be controlled by the rates of dissolution of
1 or more of these phases. The maximum possible concentration of thorium dissolved in low-
temperature aqueous systems can however be predicted with the solubility of hydrous thorium oxide,
because the solubility of this compound will result in higher concentrations of dissolved thorium than will
likely occur from the kinetically-hindered dissolution of resistant primary thorium minerals. Moreover,
hydrous thorium oxide solid is known to precipitate in laboratory experiments {i.e., short time periods)
conducted at low temperature, oversaturated conditions.
The concentrations of dissolved thorium in surface and groundwaters may also be controlled to low
values by adsorption processes. Humic substances are considered particularly important in the
adsorption of thorium. The available partition coefficient, Kd, data indicates significant retention of
thorium by most soil types.
5.9.2 General Geochemistry
Twelve isotopes of thorium are known. Their atomic masses range from 223 to 234, and all are
unstable (or radioactive) (Weast and Astle, 1980). Of these, 6 thorium isotopes exist in nature. These
include:
• 238U decay series: 234Th \tVz (half life) = 24.1 d) and 230Th (tVi = 8.0 x 104 y)
• 232Th decay series: 232Th (tVl = 1.41 x 1010 y) and 228Th (tVl = 1.913 y)
• 235U decay series: 231Th (tVi = 25.5 h) and 227Th (tVi = 18.5 d).
Natural thorium consists of essentially 1 isotope, 232Th, with trace quantities of the other isotopes.
Thorium is fertile nuclear material in that the principal isotope 232Th can be converted by capture of a
thermal neutron and 2 beta decays to fissionable 233U which does not exist in nature. The application of
thorium as a reactor fuel in the Th02 ceramic form is described in detail by Belle and Berman (1984).
Thorium occurs only in the +4 oxidation state in nature. The Th4+ ion is the largest tetravalent cation
known with a radius of approximately 1.0 A. Although the Th4+ ion is more resistant to hydrolysis than
other tetravalent ions, it forms a variety of hydroxyl species at pH values above 3 (Baes and Mesmer,
5.59
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1976; Cotton and Wilkinson, 1980). The thorium content in natural water is very low. The
concentration range in natural fresh water rarely exceeds 1 micro g/1 (0.1 pCi/1 232Th), although mg/1
concentrations of 232Th have been detected in high-acid groundwaters beneath uranium tailings sites
(Langmuir and Herman, 1980).
Although the normal ranges of thorium concentrations in igneous, metamorphic, and sedimentary rocks
are less than 50 ppm, thorium concentrations can be as high as 30 and 300 ppm, respectively, in
oceanic sand/clays and marine manganese nodules (Gascoyne, 1982). These anomalously high
concentrations of thorium have been explained by the tendency of thorium to strongly adsorb on clay
and oxyhydroxide phases (Langmuir and Herman, 1980).
The mineralogy of thorium-containing minerals is described by Frondel (1958). Most thorium-
containing minerals are considered fairly insoluble and resistant to erosion. There are few minerals in
which thorium is an essential structural constituent. Important thorium minerals include thorite
[(Th,U,Ce,Fe,efc.)Si04] and thorianite (crystalline Th02). Thorite is found in pegmatites, gneisses,
granites, and hydrothermal deposits. Thorianite is chiefly found in pegmatitic rocks, but is best known
as a detrital mineral.1 Thorium also occurs, however, as variable, trace concentrations in solid solution
in many rare-earth, zirconium, and uranium minerals. The 2 most important minerals of this type include
monazite [(Ce,La,Th)P04] and zircon (ZrSi04). Monazite and zircon are widely disseminated as
accessory minerals in igneous and metamorphic rocks. They also occur in commercial quantities in
detrital sands derived from regions of these rocks due to their resistance to erosion (Deer el al., 1967;
Frondel, 1958). Concentrations of thorium can be several weight percent in these deposits.
Because of their long half lives, 228Th (tVi = 1.913 y), 230Th (tVi = 8.0 x 104 y), and 232Th (tVi =
1.41 x 1010 y), which are all alpha-particle emitters, pose long-term health risks and are therefore
environmentally important. Contamination includes thorium-containing soils and thorium dissolved in
surface- and groundwaters. Of the contaminated sites considered in EPA/DOE/NRC (1993),
radioactive contamination of soil, surface water, and/or groundwater by 228Th, 230Th, and/or 232Th has
been identified at 21 of the 45 Superfund National Priorities List (NPL) sites and 23 of the 38 NRC
Site Decommissioning Management Plan (SDMP) sites. Some of the contamination resulted from the
separation and processing of uranium and from the use of monazite and zircon sands as source
materials for metallurgical processes.
5.9.3 Aqueous Speciation
Thorium occurs only in the +4 oxidation state in natural soil/water environments. Dissolved thorium
forms a variety of hydrolytic species, and, as a small, highly charged ion, undergoes extensive chemical
1 A detrital mineral is defined as "any mineral grain resulting from mechanical disintegration of parent
rock" (Bates and Jackson, 1980).
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interaction with water and most anions. The available thermodynamic data for thorium-containing
aqueous species and solids have been compiled and critically reviewed by Langmuir and Herman
(1980) for an analysis of the mobility of thorium in low-temperature, natural waters.
Thorium undergoes hydrolysis in aqueous solutions at pH values above 3. The distribution of thorium
hydrolytic species, shown in Figure 5.4, was calculated as a function of pH using the MINTEQA2
code and the thermodynamic data tabulated in Langmuir and Herman (1980). The aqueous species
included in the speciation calculations are listed in Table 5.14. The species distribution in Figure 5.4
was determined for a concentration of 1 micro g/1 total dissolved thorium for a water free of any
complexing ligands other than hydroxide ions. The chosen thorium concentration is based on Hem
(1985, p. 150) who gives 0.01 to 1 micro g/1 as the range expected for thorium concentrations in fresh
waters. The calculated species distribution shows that the uncomplexed ion Th4+ is the dominant ion at
pH values less than -3.5. At pH values greater than 3.5, the hydrolysis of thorium is dominated, in
order of increasing pH, by the aqueous species Th(OH)2+, Th(OH)3, and Th(OH)4(aq). The latter
2 hydrolytic complexes have the widest range of stability with pH.
The large effective charge of the Th4+ ion can induce hydrolysis to the point that polynuclear complexes
may form (Baes and Mesmer, 1976). Present knowledge of the formation of polynuclear hydrolyzed
species is poor because there is no unambiguous analytical technique to determine these species.
However, polynuclear species are believed to play a role in mobility of thorium in soil/water systems.
Langmuir and Herman (1980) list estimated thermodynamic values for the thorium polynuclear
hydrolyzed species Th2(OH)|+, Th4(OH)jj , and Th^OH)'^ based on the review of Bases and Mesmer
(1976).
5.61
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Table 5.14. Thorium aqueous species included in the
speciation calculations.
Aqueous Species
Th4+, ThOH3+, Th(OH>f, Th(OH)3, Th(OH);(aq),
Th2(OH)f, Th4(OH)r, Thg(OH)i5
Th(OH)3COi and Th(C03)f"
ThF3+, ThF2 , ThF3, ThF^aq)
ThCP, ThCl , ThCl3, ThClI(aq)
ThSOf, Th(S04)2(aq), Th(S04)?\ Th(S04#
ThH3POr, ThH2POl, Th(H2P04>i,
Th(HP04);(aq), Th(HP04)?"
In addition to hydrolytic complexes, thorium can also form various aqueous complexes with inorganic
anions such as dissolved fluoride, phosphate, chloride, and nitrate. Studies (e.g., LaFamme and
Murray, 1987) completed since the review by Langmuir and Herman (1980) indicate the presence of
dissolved thorium carbonate complexes and their importance to the solution chemistry of thorium. Due
to the lack of available data, no thorium carbonate species were listed by Langmuir and Herman
(1980). Osthols el al. (1994) have recently published thermodynamic constants for the thorium
carbonate complexes Th(OH)3COi and Th(C03)f" that are based on their solubility studies of
microcrystalline Th02 at different partial pressures of C02 in aqueous media.
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pH
Figure 5.4. Calculated distribution of thorium hydrolytic species as a function of pH. [The
species distribution is based on a concentration of 1 micro g/1 total dissolved
thorium in pure water (i.e., absence of complexing ligands other than OH") and
thermodynamic data from Langmuir and Herman (1980).]
The distribution of thorium aqueous species (Figure 5.5) was also calculated as a function of pH using
the MINTEQA2 for a concentration of 1 micro g/1 total dissolved thorium and the water composition in
Table 5.1. The thermodynamic data were principally from Langmuir and Herman (1980). The
thermodynamic constants for the aqueous species Tt^OH^COj and Th(C03)f" from Osthols el al.
(1994) were also included in these speciation calculations. Below pH 5, dissolved thorium is
dominated by thorium fluoride complexes. Between pH 5 and 7, dissolved thorium is predicted to be
dominated by thorium phosphate complexes. Although phosphate complexation is expected to have a
role in the mobility of thorium in this range of pH values, the adequacy of the thermodynamic constants
tabulated for thorium phosphate complexes in Langmuir and Herman (1980) are suspect, and may over
predict the stability of these complexes. At pH values greater than 7.5, more than 95 percent of the
dissolved thorium is predicted to be present as Th(OH)3COi. The species distribution illustrated in
5.63
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Figure 5.5 changes slightly in the pH range from 5 to 7 if the concentration of total dissolved thorium is
increased from 1 to 1,000 micro g/1. At the higher concentration of dissolved thorium, the stability of
Th(OH)3COi extends to a pH of approximately 5, the hydrolytic species Th(OH)3 becomes an
important species (about 30 percent of the dissolved thorium), and the thorium phosphate species are
no longer dominant.
Thorium organic complexes likely have an important effect on the mobility of thorium in soil/water
systems. Langmuir and Herman (1980) used citrate (QHjOy), oxalate (C2O4), and ethylenediamine
tetra-acetic acid (EDTA) (C10H12O8N2") to show the possible role of organic complexes in the mobility
of thorium in natural waters. Based on the stability constants available for thorium citrate, oxalate, and
ethylenediamine complexes, calculations by Langmuir and Herman (1980) indicate that thorium organic
complexes likely predominate over inorganic complexes in organic-rich waters and soils. For the
concentrations considered by Langmuir and Herman (1980), the ThEDTA (aq) complex dominates all
other thorium aqueous species over the pH range from 2 to 8. This would in turn have an important
effect on the solubility and adsorption of thorium in such waters.
5.9.4 Dissolution/Precipitation/Coprecipitation
The main thorium-containing minerals, thorite [(Th,U,Ce,Fe,efc.)Si04], thorianite (crystalline Th02),
monazite [(Ce,La,Th)P04) and zircon (ZrSi04), are resistant to chemical weathering and do not
dissolve readily at low-temperature in surface and groundwaters. Because these minerals form at
temperature and pressure conditions associated with igneous and metamorphic rocks, it is unlikely that
the thermodynamic equilibrium solubilities (where the rate of precipitation equals the rate of
dissolution) of these minerals will control the concentration of dissolved thorium in low-temperature
soil/water environments. The rate at which thorium is released to the environment, as might be needed
in a source-term component of a performance assessment model, may however be controlled by the
kinetic rates of aqueous dissolution {i.e., non-equilibrium conditions) of 1 or more of these phases.
5.64
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3456789 10
pH
Figure 5.5. Calculated distribution of thorium aqueous species as a function of pH for the
water composition in Table 5.1. [The species distribution is based on a
concentration of 1 micro g/1 total dissolved thorium and thermodynamic data
from Langmuir and Herman (1980) and Osthols et al. (1994, for Th(0H)3C03
and Th(C03)f"). The thermodynamic database used for these speciation
calculations did not include the constants for thorium humic acid complexes.]
The maximum concentration of dissolved thorium that may occur in a low-temperature aqueous system
can be predicted with the solubility of hydrous thorium oxide. This solid is known to precipitate in
laboratory experiments conducted at low temperature, oversaturated conditions over several weeks. If
this solid precipitates in a natural environment, it will likely alter with time to a more crystalline solid that
has a lower solubility. The solubility of hydrous thorium oxide has been studied experimentally by Rai
and coworkers (Felmy et al., 1991; Rai et al., 1995; Ryan and Rai, 1987). In 0.1 M NaC104
solutions, the measured solubility of hydrous thorium oxide ranges from about 10"8 5 mol/1 (0.0007 mg/1)
5.65
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to less than 10"9 mol/1 (0.0002 mg/1) in the pH range from 5 to 10 (Ryan and Rai, 1987). The
concentration of dissolved thorium increases to approximately 10"2 6 mol/1 (600 mg/1) as pH decreases
from 5.0 to 3.2.
Felmy el al. (1991) determined that the solubility of hydrous thorium oxide increases with increasing
ionic strength. At pH values above 7 in 3.0 M NaCl solutions, the solubility of hydrous thorium oxide
increased by approximately 2 to 3 orders of magnitude compared to that determined in 0.1 M NaC104
solutions. Moreover, the pH at which hydrous thorium oxide exhibits rapid increases in solubility with
decreasing pH changes from pH 5 in 0.1 MNaC104 to approximately pH 7 in 3.0 MNaCl. In studies
conducted at high hydroxide and carbonate concentrations, Rai el al. (1995) determined that the
solubility of hydrous thorium oxide increases dramatically in high carbonate solutions and decreases
with increases in hydroxide concentration at fixed carbonate concentrations. This supports the
assertion that soluble thorium-carbonate complexes likely dominate the aqueous speciation of thorium
dissolved in natural waters having basic pH values.
5.9.5 Adsorption/Desorption
Thorium concentrations in surface- and groundwaters may also be controlled to very low levels
(< few micro g/1) by adsorption processes. Humic substances are considered particularly important in
the adsorption of thorium (Gascoyne, 1982). Thibault el al. (1990) conducted a critical compilation
and review of published Kd data by soil type needed to model radionuclide migration from a nuclear
waste geological disposal vault to the biosphere. Thibault el al. list Kd values for thorium that range
from 207 to 13,000,000 ml/g. The range of thorium Kd values listed for organic soil was 1,579 to
1.3 x 107 ml/g. Based on our experience, the very high Kd values reported for thorium should be
viewed with caution. The studies resulting in these values should be examined to determine if the initial
concentrations of thorium used for these Kd measurements were too great and precipitation of a
thorium solid (e.g., hydrous thorium oxide) occurred during the equilibration of the thorium-spiked
soil/water mixtures. As noted in the letter report for Subtask IB, precipitation of solids containing the
contaminant of interest results in Kd values that are erroneously too high.
The adsorption of thorium on pure metal-oxide phases has also been studied experimentally in
conjunction with surface complexation models.1 Osthols (1995) studied the adsorption of thorium on
amorphous colloidal particles of silica (Si02). Their results indicate that the adsorption of thorium on
silica will only be important in the pH range from 3 to 6. In neutral and alkaline pH values, silica surface
sites are not expected to be efficient adsorbents for thorium.
Iron and manganese oxides are expected to be more important adsorbents of thorium than silica.
Hunter el al. (1988) studied the adsorption of thorium on goethite (a-FeOOH) and nsutite (y-Mn02)
1 Surface complexation models are discussed in Volume I of this report.
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in marine electrolyte solutions. Their experiments indicate that adsorption of thorium increases from
approximately 0 percent at pH 2.5-3.5 to 90-100 percent at pH 5-6.5. The adsorption of thorium
decreased with the addition of sulfate as a result of the formation of competitive aqueous complexes
with dissolved thorium. The addition of organic ligands EDTA and trans-1,2-diaminocyclohexane
tetra-acetic acid (CDTA) shifted the adsorption edges for y-Mn02 to higher pH values by more than
5-6 pH units, such that 100 percent adsorption of thorium was not observed until pH 12. LaFlamme
and Murray (1987) experimentally studied the effects of pH, ionic strength and carbonate alkalinity on
the adsorption of thorium by goethite. The adsorption edge {i.e., range in pH where metal adsorption
goes from 0 percent to approximately 90-100 percent) was measured to be in the pH range from 2
to 5. For conditions considered in their study, ionic strength was found to have no effect on the
adsorption of thorium on goethite. LaFlamme and Murray did however observe a strong influence of
carbonate alkalinity on thorium adsorption. In their experiments at pH 9.0±0.6, they observed a
decrease of thorium adsorption with the addition of 100 meq/1 carbonate alkalinity, and no measurable
adsorption of thorium at carbonate alkalinity greater than 300 meq/1. At the low particle concentrations
used in their experiments, LaFlamme and Murray attributed this reduction to the competition for surface
sites by CO3" and HCOj and the formation of soluble thorium-carbonate complexes with a net negative
charge.
5.9.6 Partition Coefficient, Kd, Values
5.9.6.1 General Availability of Kd Data
Two generalized, simplifying assumptions were established for the selection of thorium Kd values for the
look-up table. These assumptions were based on the findings of the literature review conducted on the
geochemical processes affecting thorium sorption. The assumptions are as follows:
• Thorium precipitates at concentrations greater than 10"9 M. This concentration is based on the
solubility of Th(OH)4 at pH 5.5. Although (co)precipitation is usually quantified with the
solubility construct, a very large Kd value will be used in the look-up table to approximate
thorium behavior in systems with high thorium concentrations.
• Thorium adsorption occurs at concentrations less than 10"9 M. The extent of thorium
adsorption can be estimated by soil pH.
These assumptions appear to be reasonable for a wide range of environmental conditions. However,
these simplifying assumptions are clearly compromised in systems containing high alkaline (LaFlamme
and Murray, 1987), carbonate (LaFlamme and Murray, 1987), or sulfate (Hunter el al., 1988)
concentrations, and high or low pH values (pH: 3 < x > 8: Hunter el al., 1988; LaFlamme and Murray
1987; Landa el al., 1995). These assumptions will be discussed in more detail in the following
sections.
5.67
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Based on the assumptions and limitations described above, thorium Kd values and some important
ancillary parameters that influence sorption were collected from the literature and tabulated
(Appendix I). Data included in this table, were from studies that reported Kd values (not percent
adsorbed or Freundlich or Langmuir constants) and were conducted in systems consisting of:
• Low ionic strength (< 0.1 M)
• pH values between 4 and 10.5
• Dissolved thorium concentrations less than 10"9 M
• Low humic material concentrations (<5 mg/1)
• No organic chelates (e.g., EDTA)
These aqueous chemistry constraints were selected to limit the thorium Kd values evaluated to those
that would be expected to exist in a far-field. The ancillary parameters included in these tables were
clay content, calcite concentration, pH, and CEC. Attempts were also made to include in the data set
the concentration of organic carbon and aluminum/iron oxides in the solid phase. However, these latter
ancillary parameters, which were identified above, were rarely included in the reports evaluated to
compile the data set. The data set included 17 thorium Kd values for soils and pure phase minerals.
5.9.6.2 Look-Up Tables
Linear regression analyses were conducted with data collected from the literature (described in
Appendix I). These analyses were used as guidance for selecting appropriate Kd values for the look-
up table. The Kd values used in the look-up tables could not be based entirely on statistical
consideration because the statistical analysis results were occasionally nonsensible. For example, the
data showed a negative correlation between clay content and thorium Kd values. This trend contradicts
well established principles of surface chemistry. Instead, the statistical analysis was used to provide
guidance as to the approximate range of values to use and to identify meaningful trends between the
thorium Kd values and the solid phase parameters. Thus, the Kd values included in the look-up table
were in part selected based on professional judgment. Again, only low-ionic strength solutions, similar
to that expected in far-field groundwaters, were considered in these analyses.
The look-up table for thorium Kd values was based on plume thorium concentrations and pH. These
2 parameters have an interrelated effect on thorium Kd values. The maximum concentration of
dissolved thorium may be controlled by the solubility of hydrous thorium oxides (Felmy el al., 1991;
Rai el al., 1995; Ryan and Rai, 1987). The dissolution of hydrous thorium oxides may in turn vary with
pH. Ryan and Rai (1987) reported that the solubility of hydrous thorium oxide is ~10"8 5 to ~10"9 in the
pH range of 5 to 10. The concentration of dissolved thorium increases to ~10"2 6 M (600 mg/L) as pH
decreases from 5.0 to 3.2. Thus, 2 categories based on thorium solubility were included in the look-up
table, pH 3 to 5, and pH 5 to 10. Although precipitation is typically quantified by the solubility
construct, a very large Kd value was used in the look-up table to describe high thorium concentrations
5.68
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(Table 5.15). See Appendix I for a detailed account of the process used to select the Kd values in
Table 5.15.
5.9.6.2.1 Limits of Kd Values with Respect to Organic Matter and Aluminum/Iron-Oxide
Concentrations
Of the 17 entries in the thorium Kd data set (Appendix I), none of them had accompanying organic
matter or aluminum- and iron-oxide mineral concentration data. It was anticipated that the presence of
organic matter would decrease thorium Kd values by forming thorium-organic matter complexes. These
complexes would be less prone to adsorb to surface than the uncomplexed thorium species.
Conversely, it was anticipated that the presence of aluminum- and/or iron-oxides would increase
thorium Kd values by increasing the number of adsorption (surface complexation) sites.
5.9.6.2.2 Limits of Kd Values with Respect to Dissolved Carbonate Concentrations
Of the 17 entries in the thorium Kd data set (Appendix I), none of them had accompanying carbonate
concentration data. However, 5 entries had calcite (CaC03) mineral concentrations. It was
anticipated that calcite concentrations could be used as an indirect measure, albeit poor measure, of the
amount of dissolved carbonate in the aqueous phase. Calcite concentrations had a correlation
coefficient (r) with thorium Kd value of 0.76 (Appendix I). Although this is a relatively high correlation
value, it is not significant at the 5 percent level of probability due to the small number of observations
(5 observations). Furthermore, it was anticipated that the presence of dissolved carbonate would
decrease thorium Kd values due to formation of the weaker forming carbonate-thorium complexes.
Table 5.15. Look-up table for thorium Kd values (ml/g) based on pH and dissolved thorium
concentrations. [Tabulated values pertain to systems consisting of low ionic strength (< 0.1
M), lowhumic material concentrations (<5 mg/1), no organic chelates (e.g., EDTA), and
oxidizing conditions.]
Kd (ml/g)
pH
3-5
5-8
8-10
Dissolved Th, M
Dissolved Th, M
Dissolved Th, M
<1026
>1026
<109
>109
<109
>109
Mnimum
62
300,000
1,700
300,000
20
300,000
Maximum
6,200
300,000
170,000
300,000
2,000
300,000
5.69
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5.10 Tritium Geochemistry And Kd Values
5.10.1 Overview: Important Aqueous-and Solid-Phase Parameters
Controlling Retardation
Tritium, a radioactive isotope of hydrogen with a half life (/.J of 12.3 y, readily combines with oxygen
to form water. Its behavior in aqueous systems is controlled by hydrologic processes and it migrates at
essentially the same velocity as surface- and groundwaters. Aqueous speciation, precipitation, and
sorption processes are not expected to affect the mobility of tritium in soil/water systems.
5.10.2 General Geochemistry
Tritium (3H) is a radioactive isotope of hydrogen. Three isotopes of hydrogen are known. These
include the 2 stable isotopes 'H (protium or H) and 2H (deuterium or D), and the radioactive isotope
3H (tritium or T). Tritium has a half life (/, ) of 12.3 y, and disintegrates into helium-3 (3He) by emission
of a weak beta (P") particle (Rhodehamel el al., 1971). Tritium is formed by natural and man-made
processes (Cotton and Wilkinson, 1980). Tritium is formed in the upper atmosphere mainly by the
nuclear interaction of nitrogen with fast neutrons induced by cosmic ray reactions. The relative
abundances of 'H, 2H, and 3H in natural water are 99.984, 0.016, and 0-10"15 percent, respectively
(Freeze and Cherry, 1979). Tritium can also be created in nuclear reactors as a result of processes
such as thermal neutron reactions with 6Li.
As an isotope of hydrogen, tritium in soil systems behaves like hydrogen and will exist in ionic, gaseous,
and liquid forms (e.g., tritiated water, HTO). Ames and Rai (1978) discuss the geochemical behavior
of tritium, and summarize field and laboratory studies of the mobility of tritium in soil systems. Because
tritium readily combines with oxygen to form water, its behavior in aqueous systems is controlled by
hydrologic processes. Because of these properties and its moderately long half life, tritium has been
used as an environmental isotopic indicator to study hydrologic flow conditions. Rhodehamel el al.
(1971) present an extensive bibliography (more than 1,200 references) and summarize the use of tritium
in hydrologic studies through 1966. Tritium has been used to study recharge and pollution of
groundwater reservoirs; permeability of aquifers; velocity, flow patterns, and stratification of surface-
and groundwater bodies; dispersion and mixing processes in surface- and groundwaters; movement of
soil moisture; chemisorption of soils and water-containing materials; biological uptake and release of
water; and secondary recovery techniques for petroleum resources. IAEA (1979) published the
proceedings from a 1978 conference dealing with the behavior of tritium in the environment. The
conference was designed to provide information on the residence time and distribution of tritium in
environmental systems and the incorporation of tritium into biological materials and its transfer along the
food chain.
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Tritium-contamination may include surface- and groundwater, soil, sediment, and air components at a
site. Of the contaminated sites considered in EPA/DOE/NRC (1993), tritium contamination has been
identified at 12 of the 45 Superfund National Priorities List (NPL) sites and 1 of the 38 NRC Site
Decommissioning Site Plan (SDMP) sites.
5.10.3 Aqueous Speciation
Because tritium oxidizes rapidly to form isotopic water, aqueous speciation reactions do not affect the
mobility of tritium in soil/water systems.
5.10.4 Dissolution/Precipitation/Coprecipitation
Neither precipitation or coprecipitation processes affect the mobility of tritium in soil/water systems.
5.10.5 Adsorption/Desorption
Because tritium readily combines with oxygen to form water, its behavior in aqueous systems is
controlled by hydrologic processes and it migrates at essentially the same velocity as surface and
groundwaters. Sorption processes are therefore not expected to be important relative to the movement
of tritium through aqueous environments. Typically, a partition coefficient, Kd, of 0 ml/g is used to
model the migration of tritium in soil and groundwater environments. As an exception, Thibault el al.
(1990), based on a review of published studies, list 0.04 to 0.1 ml/g as the range for Kd values for
tritium in sandy soils. Although tritium may substitute for hydrogen in water on clays and other hydrated
soil constituents, Ames and Rai (1978) indicate that this reaction is not important relative to the mobility
of tritium based on their review of published laboratory and field studies. Some laboratory studies
considered in their review describe fixation of isotopic water on clays and other hydrated minerals,
while others indicate minimal fixation. All field studies reviewed by Ames and Rai indicate that tritium
migrates at the same velocity as surface- and groundwaters.
5.10.6 Partition Coefficient, Kd, Values
A review of the literature pertaining to Kd values for tritium was not conducted given the limited
availability of Kd values for tritium (see section above) and limited importance of sorption processes
relative to the mobility of tritium in aqueous environments.
5.11 Uranium Geochemistry and Kd Values
5.11.1 Overview: Important Aqueous- and Solid-Phase Parameters
Controlling Retardation
In essentially all geologic environments, +4 and +6 are the most important oxidation states of uranium.
5.65
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Uranium(VI) species dominate in oxidizing environments. Uranium(VI) retention by soils and rocks in
alkaline conditions is poor because of the predominance of neutral or negatively charged species. An
increase in C02 pressure in soil solutions reduces U(VI) adsorption by promoting the formation of
poorly sorbing carbonate complexes. Uranium(TV) species dominate in reducing environments.
Uranium(IV) tends to hydrolyze and form strong hydrolytic complexes. Uranium(TV) also tends to
form sparingly soluble precipitates that commonly control U(TV) concentrations in groundwaters.
Uranium(TV) forms strong complexes with naturally occurring organic materials. Thus, in areas where
there are high concentrations of dissolved organic materials, U(IV)-organic complexes may increase
U(IV) solubility. There are several ancillary environmental parameters affecting uranium migration. The
most important of these parameters include redox status, pH, ligand (carbonate, fluoride, sulfate,
phosphate, and dissolved carbon) concentrations, aluminum- and iron-oxide mineral concentrations,
and uranium concentrations.
5.11.2 General Geochemistry
Uranium (U) has 14 isotopes; the atomic masses of these isotopes range from 227 to 240. All uranium
isotopes are radioactive. Naturally-occurring uranium typically contains 99.283 percent 238U, 0.711
percent 235U, and 0.0054 percent 234U by weight. The half-lives of these isotopes are 4.51 x 109 y, 7.1
x 108 y, and 2.47 x 105 y, respectively. Uranium can exist in the +3, +4, +5, and +6 oxidation states,
of which the +4 and +6 states are the most common states found in the environment.
The mineralogy of uranium-containing minerals is described by Frondel (1958). Uranium in the +4 and
+6 oxidation states exists in a variety of primary and secondary minerals. Important U(TV) minerals
include uraninite (U02 through UO2.25) and coffinite [USi04] (Frondel, 1958; Langmuir, 1978).
Aqueous U(TV) is inclined to form sparingly soluble precipitates, adsorb strongly to mineral surfaces,
and partition into organic matter, thereby reducing its mobility in groundwater. Important U(VI)
minerals include carnotite [(^(UC^MVC^], schoepite (U03 2H20), rutherfordine (U02C03),
tyuyamunite [Ca(U02)2(V04)2], autunite [Ca(U02)2(P04)2], potassium autunite [K2(U02)2(P04)2],
and uranophane [Ca(U02)2(Si030H)2] (Frondel, 1958; Langmuir, 1978). Some of these are
secondary phases which may form when sufficient uranium is leached from contaminated wastes or a
disposal system and migrates downstream. Uranium is also found in phosphate rock and lignite1 at
concentrations that can be commercially recovered. In the presence of lignite and other sedimentary
carbonaceous substances, uranium enrichment is believed to be the result of uranium reduction to form
insoluble precipitates, such as uraninite.
Contamination includes airborne particulates, uranium-containing soils, and uranium dissolved in
surface- and groundwaters. Of the contaminated sites considered in EPA/DOE/NRC (1993),
radioactive contamination by 234U, 235U, and/or 238U has been identified at 35 of the 45 Superfund
1 Lignite is a coal that is intermediate in coalification between peat and subbituminous coal.
5.66
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National Priorities List (NPL) sites and 26 of the 38 NRC Site Decommissioning Site Plan (SDMP)
sites.
5.11.3 Aqueous Speciation
Because of its importance in nuclear chemistry and technology, a great deal is known about the
aqueous chemistry of uranium [reviewed by Baes and Mesmer (1976), Langmuir (1978), and Wanner
and Forest (1992)]. Uranium can exist in the +3, +4, +5, and +6, oxidation states in aqueous
environments. Dissolved U(m) easily oxidizes to U(TV) under most reducing conditions found in
nature. The U(V) aqueous species (UOj) readily disproportionates to U(IV) and UCVI).1
Consequently, U(TV) and U(VI) are the most common oxidation states of uranium in nature. Uranium
will exist in the +6 and +4 oxidation states, respectively, in oxidizing and more reducing environments.
Both uranium species, UOf and U4+, hydrolyze readily. The U4+ ion is more readily hydrolyzed than
UC>2+, as would be expected from its higher ionic charge. Langmuir (1978) calculated U(TV)
speciation in a system containing typical natural water concentrations of chloride (10 mg/1), fluoride
(0.2 mg/1), phosphate (0.1 mg/1), and sulfate (100 mg/1). Below pH 3, UF22+ was the dominant uranium
species. The speciation of dissolved U(TV) at pH values greater than 3 is dominated by hydrolytic
species such as U(OH)3 and U(OH)5|(aq). Complexes with chloride, fluoride, phosphate, and sulfate
were not important above pH 3. The total U(TV) concentration in solution is generally quite low,
between 3 and 30 |xg/l, because of the low solubility of U(TV) solid phases (Bruno el al., 1988; Bruno
el al., 1991). Precipitation is discussed further in the next section.
Dissolved U(VI) hydrolyses to form a number of aqueous complexes. The distribution of U(VI)
species is presented in Figures 5.6a-b and 5.7. The distribution of uranyl hydrolytic species
(Figures 5.6a-b) was calculated as a function of pH using the MINTEQA2 code. The U(VI) aqueous
species included in the speciation calculations are listed in Table 5.16. The thermodynamic data for
these aqueous species were taken primarily from Wanner and Forest (1992). Because dissolved
uranyl ions can be present as polynuclear2 hydroxyl complexes, the hydrolysis of uranyl ions under oxic
conditions is therefore dependent on the concentration of total dissolved uranium. To demonstrate this
aspect of uranium chemistry, 2 concentrations of total dissolved uranium, 0.1 and 1,000 |ig/l, were used
in these calculations. Hem (1985, p. 148) gives 0.1 to 10 |ig/l as the range for dissolved uranium in
1 Disproportionation is defined in the glossary at the end of this letter report. This particular
disproportionation reaction can be described as:
2UO2 + 4H30+ = \JOf + U4+.
2 A polynuclear species contains more than 1 central cation moiety, e.g., (U02)2C03(0H)3 and
Pb4(OH)4+.
5.67
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most natural waters. For waters associated with uranium ore deposits, Hem states that the uranium
concentrations may be greater than 1,000 jig/1.
In a U(VI)-water system, the dominant species were U02 at pH values less than 5, U02(0H)2 (aq) at
pH values between 5 and 9, and U02(0H)3 at pH values between 9 and 10. This was true for both
uranium concentrations, 0.1 |ig/l (Figure 5.6a) and 1,000 |ig/l dissolved U(VI) (Figure 5.6b). At
1,000 |ig/l dissolved uranium, some polynuclear species, (U02)3(0H)5 and (U02)2(0H)2+, were
calculated to exist between pH 5 and 6. Morris el al. (1994) using spectroscopic techniques provided
additional proof that an increasing number of polynuclear species were formed in systems containing
higher concentrations of dissolved uranium.
A large number of additional uranyl species (Figure 5.7) are likely to exist in the chemically more
complicated system such as the water composition in Table 5.1 and 1,000 |ig/l dissolved U(VI). At
pH values less than 5, the U02F+ species dominates the system, whereas at pH values greater than 5,
carbonate complexes [U02C03(aq), U02(C03)2", U02(C03)3"] dominate the system. These
calculations clearly show the importance of carbonate chemistry on U(VI) speciation. For this water
composition, complexes with chloride, sulfate, and phosphate were relatively less important. Consistent
with the results in Figure 5.7, Langmuir (1978) concluded that the uranyl complexes with chloride,
phosphate, and sulfate were not important in a typical groundwater. The species distribution illustrated
in Figure 5.7 changes slightly at pH values greater than 6 if the concentration of total dissolved uranium
is decreased from 1,000 to 1 |ig/l. At the lower concentration of dissolved uranium, the species
(U02)2C03(0H)i is no longer present as a dominant aqueous species.
Sandino and Bruno (1992) showed that U02+-phosphate complexes [UOoHPO^aq) and U02PO;]
could be important in aqueous systems with a pH between 6 and 9 when the total concentration ratio
P04(total)/C03(total) is greater than 0.1. Complexes with sulfate, fluoride, and possibly chloride are
potentially important uranyl species where concentrations of these anions are high. However, their
stability is considerably less than the carbonate and phosphate complexes (Wanner and Forest, 1992).
Organic complexes may also be important to uranium aqueous chemistry. The uncomplexed uranyl ion
has a greater tendency to form complexes with fulvic and humic acids than many other metals with a +2
valence (Kim, 1986). This has been attributed to the greater "effective charge" of the uranyl ion
compared to other divalent metals. The effective charge has been estimated to be about +3.3 for
U(VI) in U02+. Kim (1986) concluded that, in general, +6 actinides, including U(VI), would have
approximately the same tendency to form humic- or fulvic-acid complexes as to hydrolyze or form
carbonate complexes. This suggests that the dominant reaction with the uranyl ion that will take place in
a groundwater will depend largely on the relative concentrations of hydroxide, carbonate, and organic
material concentrations. He also concluded, based on comparison of stability constants, that the
tendency for U4+ to form humic- or fulvic-acid complexes is less than its tendency to hydrolyze or form
carbonate complexes. Importantly, U(IV) and U(VI) can form stable organic complexes, thereby
increasing their solubility and mobility.
5.68
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Table 5.16. Uranium(VI) aqueous species included in the
speciation calculations.
Aqueous Species
UOf, U020H+, U02(0H)2(aq), U02(0H)i,, U02(0Htf,
(U02)20H3+, (U02)2(0H)r, (U02)3(0H)f, (U02)3(0H)5,
(U02)3(0H)7, (U02)4(0H)y, U6(OH)j5
U02C(%q), U02(C03)i", U02(C03)t U02(C0J\
(U02)3(C03)r, (U02)n(C03)6(0H)i2, (U02)2C03(0H)3
U02P0;, U02HP05|(aq), U02H2P04, U02H3P0f,
U02(H2P04)2(aq), U02(H2P04)(H3P04)+,
U02S04(aq), U02(S04)2"
U02N03
U02C1+, U02Cl2(aq), U02F+, U02F^(aq), U02Fi, UOoFj"
U02Si0(0H)3
5.11.4 Dissolution/Precipitation/Coprecipitation
Dissolution, precipitation, and coprecipitation have a much greater effect on the concentrations of
U(IV) than on the concentration of U(VI) in groundwaters. In most cases, these processes will likely
not control the concentration of U(VI) in oxygenated groundwaters far from a uranium source. Near a
uranium source, or in reduced environments, these processes tend to become increasingly important
and several (co)precipitates may form depending on the environmental conditions (Falck, 1991;
Frondel, 1958). Reducing conditions may exist in deep aquifers, marsh areas, or engineered barriers
that may cause U(TV) to precipitate. Important U(TV) minerals include uraninite (compositions ranging
from U02 to U02 25), coffinite (USi04), and ningyoite [CaU(P04)2-2H20] (Frondel, 1958; Langmuir,
1978). Important U(VI) minerals include carnotite [(K2(U02)2(V04)2], schoepite (U03-2H20),
rutherfordine (U02C03), tyuyamunite [Ca(U02)2(V04)2], autunite [Ca(U02)2(P04)2], potassium
autunite [K2(U02)2(P04)2], and uranophane [Ca(U02)2(Si030H)2] (Frondel, 1958; Langmuir, 1978).
Carnotite, a U(VI) mineral, is found in the oxidized zones of uranium ore deposits and uraninite, a
5.69
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U(TV) mineral, is a primary mineral in reducing ore zones (Frondel, 1958). The best way to model the
concentration of precipitated uranium is not with the Kd construct, but through the use of solubility
pH
constants.
Figure 5.6a. Calculated distribution of U(VI) hydrolytic species as a function of pH
at 0.1 |ig/l total dissolved U(VI). [The species distribution is based on U(VI)
dissolved in pure water (i.e., absence of complexing ligands other than OH")
and thermodynamic data from Wanner and Forest (1992).]
5.70
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pH
Figure 5.6b. Calculated distribution of U(VI) hydrolytic species as a function of pH at
1,000 |ig/l total dissolved U(VI). [The species distribution is based on U(VI)
dissolved in pure water and thermodynamic data from Wanner and Forest
(1992).]
5.71
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100
uo,.cco,v
=
s
xi
=
a>
u
—
a>
-
80
60
40
20
uo2r
UO2CO3 (aq) U02(C03)2*
10
pH
Figure 5.7. Calculated distribution of U(VI) aqueous species as a function of pH for the
water composition in Table 5.1. [The species distribution is based on a
concentration of 1,000 |ig/l total dissolved U(VI) and thermodynamic data from
Wanner and Forest (1992).]
5.11.5 Sorption/Desorption
In low ionic strength solutions with low U(VI) concentrations, dissolved uranyl concentrations will likely
be controlled by cation exchange and adsorption processes. The uranyl ion and its complexes adsorb
onto clays (Ames el al., 1982; Chisholm-Brause el al., 1994), organics (Borovec el al., 1979; Read
et al., 1993; Shanbhag and Choppin, 1981), and oxides (Hsi and Langmuir, 1985; Waite el al.,
1994). As the ionic strength of an oxidized solution increases, other ions, notably Ca2+, Mg2+, and K+,
will displace the uranyl ion from soil exchange sites, forcing it into solution. For this reason, the uranyl
5.72
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ion is particularly mobile in high ionic-strength solutions. Not only will other cations dominate over the
uranyl ion in competition for exchange sites, but carbonate ions will form strong soluble complexes with
the uranyl ion, further lowering the activity of this ion while increasing the total amount of uranium in
solution (Yeh and Tripathi, 1991).
Some of the sorption processes to which uranyl ion is subjected are not completely reversible.
Sorption onto iron and manganese oxides can be a major process for extraction of uranium from
solution (Hsi and Langmuir, 1985; Waite el al., 1994). These oxide phases act as a somewhat
irreversible sink for uranium in soils. Uranium bound in these phases is not generally in isotopic
equilibrium with dissolved uranium in the same system, suggesting that the reaction rate mediating the
transfer of the metal between the 2 phases is slow.
Naturally occurring organic matter is another possible sink for U(VI) in soils and sediments. The
mechanisms by which uranium is sequestered by organic matter have not been worked out in detail.
One possible process involves adsorption of uranium to humic substances through rapid ion-exchange
and complexation processes with carboxylic and other acidic functional groups (Boggs el al., 1985;
Borovec et al., 1979; Idiz el al., 1986; Shanbhag and Choppin, 1981; Szalay, 1964). These groups
can coordinate with the uranyl ion, displacing waters of hydration, to form stable complexes. A
process such as this probably accounts for a significant fraction of the organically bound uranium in
surface and subsurface soils. Alternatively, sedimentary organics may act to reduce dissolved U(VI)
species to U(TV) (Nash el al., 1981).
Uranium sorption to iron oxide minerals and smectite clay has been shown to be extensive in the
absence of dissolved carbonate (Ames et al., 1982; Hsi and Langmuir, 1985; Kent el al., 1988).
However, in the presence of carbonate and organic complexants, sorption has been shown to be
substantially reduced or severely inhibited (Hsi and Langmuir, 1985; Kent el al., 1988).
Aqueous pH is likely to have a profound effect on U(VI) sorption to solids. There are 2 processes by
which it influences sorption. First, it has a great impact on uranium speciation (Figures 5.6a-b and 5.7)
such that poorer-adsorbing uranium species will likely exist at pH values between about 6.5 and 10.
Secondly, decreases in pH reduce the number of exchange sites on variable charged surfaces, such as
iron-, aluminum-oxides, and natural organic matter.
5.73
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5.11.6 Partition Coefficient, Kd, Values
5.11.6.1 General Availability of Kd Values
More than 20 references (Appendix J) that reported Kd values for the sorption of uranium onto soils,
crushed rock material, and single mineral phases were identified during this review.1 These studies
were typically conducted to support uranium migration investigations and safety assessments associated
with the genesis of uranium ore deposits, remediation of uranium mill tailings, agriculture practices, and
the near-surface and deep geologic disposal of low-level and high-level radioactive wastes (including
spent nuclear fuel). These studies indicated that pH and dissolved carbonate concentrations are the
2 most important factors influencing the adsorption behavior of U(VI).
The uranium Kd values listed in Appendix J exhibit large scatter. This scatter increases from
approximately 3 orders of magnitude at pH values below pH 5, to approximately 3 to 4 orders of
magnitude from pH 5 to 7, and approximately 4 to 5 orders of magnitude at pH values from pH 7 to 9.
At the lowest and highest pH regions, it should be noted that 1 to 2 orders of the observed variability
actually represent uranium Kd values that are less than 10 ml/g. At pH values less than 3.5 and greater
than 8, this variability includes Kd values of less than 1 ml/g.
Uranium Kd values show a trend as a function of pH. In general, the adsorption of uranium by soils and
single-mineral phases in carbonate-containing aqueous solutions is low at pH values less than 3,
increases rapidly with increasing pH from pH 3 to 5, reaches a maximum in adsorption in the pH range
from pH 5 to 8, and then decreases with increasing pH at pH values greater than 8. This trend is
similar to the in situ Kd values reported by Serkiz and Johnson (1994), and percent adsorption values
measured for uranium on single mineral phases such as those reported for iron oxides (Hsi and
Langmuir, 1985; Tripathi, 1984; Waite el al., 1992, 1994), clays (McKinley el al., 1995; Turner el
al., 1996; Waite el al., 1992), and quartz (Waite el al., 1992). This pH-dependent behavior is related
to the pH-dependent surface charge properties of the soil minerals and complex aqueous speciation of
dissolved U(VI), especially near and above neutral pH conditions where dissolved U(VI) forms strong
anionic uranyl-carbonato complexes with dissolved carbonate.
5.11.6.2 Look-Up Table
Solution pH was used as the basis for generating a look-up table for the range of estimated minimum
and maximum Kd values for uranium. Given the orders of magnitude variability observed for reported
1 Since the completion of our review and analysis of Kd data for the selected contaminants and
radionuclides, the studies by Pabalan et al. (1998), Payne el al. (1998), Redden el al. (1998),
Rosentreter el al. (1998), and Thompson el al. (1998) were identified and may be of interest to the
reader.
5.74
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uranium Kd values, a subjective approach was used to estimate the minimum and maximum Kd values
for uranium as a function of pH. These values are listed in Table 5.17. For Kd values at non-integer
pH values, especially given the rapid changes in uranium adsorption observed at pH values less than 5
and greater than 8, the reader should assume a linear relationship between each adjacent pair of pH-Kd
values listed in Table 5.17.
Table 5.17. Look-up table for estimated range of Kd values for uranium based on pH.
K<1
(ml/g)
pH
3
4
5
6
7
8
9
10
Minimum
<1
0.4
25
100
63
0.4
<1
<1
Maximum
32
5,000
160,000
1,000,000
630,000
250,000
7,900
5
The boundary representing the minimum limit for uranium Kd values is based on values calculated for
quartz from data given in Waite el al. (1992) and the Kd values reported by Kaplan el al. (1996,
1998), Lindenmeirer el al. (1995), and Serne el al. (1993). It is unlikely that actual Kd values for
U(VI) can be much lower than those represented by this lower boundary. At the pH extremes along
this curve, the uranium Kd values are very small. Moreover, if one considers potential sources of error
resulting from experimental methods, it is difficult to rationalize uranium Kd values much lower than this
lower boundary.
The curve representing the maximum limit for uranium Kd values is based on Kd values calculated for
fenihydrite and kaolinite from data given in Waite el al. (1992). It is estimated that this maximum limit
is biased high, possibly by an order of magnitude or more especially at pH values greater than 5. This
estimate is partially based on the distribution of measured Kd values listed in Appendix J, and the
assumption that some of the very large Kd measurements may have included precipitation of uranium-
containing solids due to starting uranium solutions being oversaturated. Moreover, measurements of
uranium adsorption onto crushed rock materials may include U(VI)/U(IV) redox/precipitation reactions
resulting from contact of dissolved U(VI) with Fe(II) exposed on the fresh mineral surfaces.
5.11.6.2.1 Limits of Kd Values with Respect to Dissolved Carbonate Concentrations
As noted in several studies summarized in Appendix J and in surface complexation studies of uranium
adsorption by Tripathi (1984), Hsi and Langmuir (1985), Waite el al. (1992, 1994), McKinley el al.
(1995), Duff and Amrheim (1996), Turner el al. (1996), and others, dissolved carbonate has a
significant effect on the aqueous chemistry and solubility of dissolved U(VI) through the formation of
5.75
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strong anionic carbonato complexes. In turn, this complexation affects the adsorption behavior of
U(VI) at alkaline pH conditions.
No attempt was made to statistically fit the Kd values summarized in Appendix J as a function of
dissolved carbonate concentrations. Typically carbonate concentrations were not reported and/or
discussed, and one would have to make assumptions about possible equilibrium between the solutions
and atmospheric or soil-related partial pressures of C02 or carbonate phases present in the soil
samples. Given the complexity of these reaction processes, it is recommended that the reader consider
the application of geochemical reaction codes, and surface complexation models in particular, as the
best approach to predicting the role of dissolved carbonate in the adsorption behavior of uranium and
derivation of U(VI) Kd values when site-specific Kd values are not available.
5.11.6.2.2 Limits of Kd Values with Respect to Clay Content and CEC
No attempt was made to statistically fit the Kd values summarized in Appendix J as a function of clay
content or CEC. The extent of clay content and CEC data, as noted from information compiled during
this review, is limited to a few studies that cover somewhat limited geochemical conditions. Moreover,
Serkiz and Johnson (1994) found no correlation between their uranium in situ Kd values and the clay
content or CEC of their soils. Their systems covered the pH conditions from 3 to 7.
However, clays have an important role in the adsorption of uranium in soils. Attempts have been made
(e.g., Borovec, 1981) to represent this functionality with a mathematical expression, but such studies
are typically for limited geochemical conditions. Based on studies by Chisholm-Brause (1994), Morris
et al. (1994), McKinley et al. (1995), Turner et al. (1996), and others, uranium adsorption onto clay
minerals is complicated and involves multiple binding sites, including exchange and edge-coordination
sites. The reader is referred to these references for a detailed treatment of the uranium adsorption on
smectite clays and application of surface complexation modeling techniques for such minerals.
5.11.6.2.3 Use of Surface Complexation Models to Predict Uranium Kd Values
As discussed in Chapter 4 and in greater detail in Volume I of this report, electrostatic surface
complexation models (SCMs) incorporated into chemical reaction codes, such as EPA's MINTEQA2,
may be used to predict the adsorption behavior of some radionuclides and other metals and to derive
Kd values as a function of key geochemical parameters, such as pH and carbonate concentrations.
Typically, the application of surface complexation models is limited by the availability of surface
complexation constants for the constituents of interest and competing ions that influence their adsorption
behavior.
The current state of knowledge regarding surface complexation constants for uranium adsorption onto
important soil minerals, such as iron oxides, and development of a mechanistic understanding of these
reactions is probably as advanced as those for any other trace metal. In the absence of site-specific Kd
5.76
-------
values for the geochemical conditions of interest, the reader is encouraged to apply this technology to
predict bounding uranium Kd values and their functionality with respect to important geochemical
parameters.
5.12 Conclusions
One objective of this report is to provide a "thumb-nail sketch" of the geochemistry of cadmium,
cesium, chromium, lead, plutonium, radon, strontium, thorium, tritium, and uranium. These
contaminants represent 6 nonexclusive contaminant categories: cations, anions, radionuclides,
non-attenuated contaminants, attenuated contaminants, and redox-sensitive contaminants (Table 5.18).
By categorizing the contaminants in this manner, general geochemical behaviors of 1 contaminant may
be extrapolated by analogy to other contaminants in the same category. For example, anions, such as
N03" and CI", commonly adsorb to geological materials to a limited extent. This is also the case
observed for the sorption behavior of anionic Cr(VI).
Important solution speciation, (co)precipitation/dissolution, and adsorption reactions were discussed for
each contaminant. The species distributions for each contaminant were calculated using the chemical
equilibria code MINTEQA2 (Version 3.11, Allison el al., 1991) for the water composition described
in Tables 5.1 and 5.2. The purpose of these calculations was to illustrate the types of aqueous species
that might exist in a groundwater. A summary of the results of these calculations are presented in Table
5.19. The speciation of cesium, radon, strontium, and tritium does not change between the pH range of
3 and 10; they exist as Cs+, Rn°, Sr2+, and HTO, respectively (Ames and Rai, 1978; Rai and Zachara,
1984). Chromium (as chromate, CrOj), cadmium, and thorium have 2 or 3 different species across
this pH range. Lead, plutonium, and uranium have several species. Calculations show that lead forms a
large number of stable complexes. The aqueous speciation of plutonium is especially complicated
because it may exist in groundwaters in multiple oxidation states [Pu(in), Pu(TV), Pu(V), and Pu(VI)]
and it forms stable complexes with a large number of ligands. Because of redox sensitivity, the
speciation of uranium exhibits a large number of stable complexes. Uranium(VI) also forms polynuclear
complex species [complexes containing more than 1 mole of uranyl [e.g., (U02)2C030H"].
One general conclusion that can be made from the results in Table 5.19 is that, as the pH increases, the
aqueous complexes tend to become increasingly more negatively charged. For example, lead,
plutonium, thorium, and uranium 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. 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.
5.77
-------
Table 5.18. Selected chemical and transport properties of the contaminants.
Elemen
t
Radio-
nuclide 1
Primary Species at pH 7
and Oxidizing Conditions
Redox
Sensitive
2
Transport Through
Soils at pH 7
Cationic
Anionic
Neutral
Not
Retarded5
Retarded
3
Cd
X
X
X
Cs
X
X
X
Cr
X
X
X
X
Pb
X
X
X
X
Pu
X
X
X
X
X
Rn
X
X
X
Sr
X
X
X
Th
X
X
X
3H
X
X
X
U
X
X
X
X
X
1 Contaminants that are primarily a health concern as a result of their radioactivity are identified
in this column. Some of these contaminants also exist as stable isotopes (e.g., cesium and
strontium).
2 The redox status column identifies contaminants (Cr, Pu, and U) that have variable oxidation
states within the pH and Eh limits commonly found in the environment and contaminants (Cd and
Pb) whose transport is affected by aqueous complexes or precipitates involving other redox-
sensitive constituents (e.g., dissolved sulfide).
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.
5.78
-------
Table 5.19. Distribution of dominant contaminant species at 3 pH values for an oxidizing
water described in Tables 5.1 and 5.2.1
pH 3
pH 7
pH 10
Element
Species
%
Species
%
Species
%
Cd
Cd2+
97
Cd2+
84
CdC03(aq)
96
CdHCa3
6
CdCOj(aq)
6
Cs
Cs+
100
Cs+
100
Cs+
100
Cr
HCrO,
99
CrO2,
78
CrO2;
99
HCrO,
22
Pb
Pb2+
96
PbCO;(aq)
75
PbC03(aq)
50
PbSO,(aq)
4
Pb2+
15
Pb(C03)2
38
PbHCO^
7
Pb(OH)2(aq)
9
PbOH+
3
Pb(OH)+
3
Pu
PuF2+
69
Pu(OH)2(C03)2
94
Pu(OH)2(C03)2
90
Pu02
24
Pu(OH)';(aq)
5
Pu(OH)';(aq)
10
Pu3+
5
Rn
Rn°
100
Rn°
100
Rn°
100
Sr
Sr2+
99
Sr2+
99
Sr2+
86
SrC03(aq)
12
Th
ThF2+
54
Th(HP04)23-
76
Th(0H)3C03
99
ThF3
42
Th(0H)3C03
22
3H
HTO
100
HTO
100
HTO
100
U
uo2f+
62
U02(C03)22
58
U02(C03)t
63
0.1 I^g/l
u
-------
Another objective of this report is to identify the important chemical, physical, and mineralogical
characteristics controlling sorption of these contaminants. These key aqueous- and solid-phase
parameters were used to assist in the selection of appropriate minimum and maximum Kd values. 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
1 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 5.20.
Sorption of all the contaminants, except tritium and radon, included in this study is influenced to some
degree by pH. The effect of pH on both adsorption and (co)precipitation is pervasive. The pH, per se,
typically has a small direct effect on contaminant adsorption. However, it has a profound effect on a
number of aqueous and solid phase properties that in turn have a direct effect on contaminant sorption.
The effects of pH on sorption are discussed in greater detail in Volume I. As discussed above, pH has
a profound effect on aqueous speciation (Table 5.19), which may affect adsorption. 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 polynuclear complexes,
oxidation state of contaminants and complexing/precipitating ligands, and H+ -competition for adsorption
sites.
The redox status of a system also influences the sorption of several contaminants included in this study
(Table 5.20). Like pH, redox has direct and indirect effects on contaminant (co)precipitation. The
direct effect occurs with contaminants like uranium and chromium where the oxidized species form
more soluble solid phases than the reduced species. Redox conditions also have a direct effect on the
sorption of plutonium, but the effects are quite complicated. 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) (SO4", sulfate) will convert into S(II) (S2", sulfide) and then the S(II)
may form sparingly soluble cadmium and lead precipitates. Thus, these 2 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).
Unlike most ancillary parameters, the effect of redox on sorption can be quite dramatic. 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
(Ames and Rai, 1978; Rai and Zachara, 1984).
5.80
-------
Table 5.20. Some of the more important aqueous- and solid-phase parameters
affecting contaminant sorption.1
Element
Important Aqueous- and Solid-Phase Parameters Influencing
Contaminant Sorption2
Cd
[Aluminum/Iron-Oxide Minerals], [Calcium], Cation Exchange Capacity,
[Clay Mineral], [Magnesium], [Organic Matter], pH, Redox, [Sulfide]
Cs
[Aluminum/Iron-Oxide Minerals], [Ammonium], Cation Exchange Capacity,
[Clay Mineral], [Mica-Like Clays], pH, [Potassium]
Cr
[Aluminum/Iron-Oxide Minerals], [Organic Matter], pH, Redox
Pb
[Aluminum/Iron-Oxide Minerals], [Carbonate, Fluoride, Sulfate,
Phosphate], [Clay Mineral], [Organic Matter], pH, Redox
Pu
[Aluminum/Iron-Oxide Minerals], [Carbonate, Fluoride, Sulfate,
Phosphate], [Clay Mineral], [Organic Matter], pH, Redox
Rn
None
Sr
Cation Exchange Capacity, [Calcium], [Carbonate], pH, [Stable Strontium]
Th
[Aluminum/Iron-Oxide Minerals], [Carbonate], [Organic Matter], pH
3H
None
U
[Aluminum/Iron-Oxide Minerals], [Carbonate, Fluoride, Sulfate,
Phosphate], [Clay Mineral], [Organic Matter], pH, Redox, [U]
1 For groundwaters with low ionic strength and low concentrations of contaminant,
chelating agents (e.g., EDTA), and natural organic matter.
2 Parameters listed in alphabetical order. Square brackets represent concentration.
5.81
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APPENDICES
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APPENDIX A
Acronyms, Abbreviations, Symbols, and Notation
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Appendix A
Acronyms, Abbreviations, Symbols, and Notation
A. 1.0 Acronyms And Abbreviations
AA
Atomic absorption
ASCII
American Standard Code for Information Interchange
ASTM
American Society for Testing and Materials
CCM
Constant capacitance (adsorption) model
CDTA
Trans- 1,2-diaminocyclohexane tetra-acetic acid
CEAM
Center for Exposure Assessment Modeling at EPA's Environmental Research
Laboratory in Athens, Georgia
CEC
Cation exchange capacity
CERCLA
Comprehensive Environmental Response, Compensation, and Liability Act
DLM
Diffuse (double) layer (adsorption) model
DDLM
Diffuse double layer (adsorption) model
DOE
U.S. Department of Energy
DTPA
Diethylenetriaminepentacetic acid
EDTA
Ethylenediaminetriacetic acid
EDX
Energy dispersive x-ray analysis
EPA
U.S. Environmental Protection Agency
EPRI
Electric Power Research Institute
HEDTA
N-(2-hydroxyethyl) ethylenedinitrilotriacetic acid
HLW
High level radioactive waste
IAEA
International Atomic Energy Agency
ICP
Inductively coupled plasma
ICP/MS
Inductively coupled plasma/mass spectroscopy
IEP (or iep)
Isoelectric point
LLNL
Lawrence Livermore National Laboratory, U.S. DOE
LLW
Low level radioactive waste
MCL
Maximum Contaminant Level
MEPAS
Multimedia Environmental Pollutant Assessment System
MS-DOS®
Microsoft® disk operating system (Microsoft and MS-DOS are register
trademarks of Microsoft Corporation.)
NPL
Superfund National Priorities List
NRC
U.S. Nuclear Regulatory Commission
NWWA
National Water Well Association
OERR
Office of Remedial and Emergency Response, U.S. EPA
ORIA
Office of Radiation and Indoor Air, U.S. EPA
OSWER
Office of Solid Waste and Emergency Response, U.S. EPA
A.2
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PC
Personal computers operating under the MS-DOS® and Microsoft® Windows
operating systems (Microsoft® Windows is a trademark of Microsoft
Corporation.)
PNL
Pacific Northwest Laboratory. In 1995, DOE formally changed the name of the
Pacific Northwest Laboratory to the Pacific Northwest National Laboratory.
PNNL
Pacific Northwest National Laboratory, U.S. DOE
PZC
Point of zero charge
RCRA
Resource Conservation and Recovery Act
SCM
Surface complexation model
SDMP
NRC's Site Decommissioning Management Plan
TDS
Total dissolved solids
TLM
Triple-layer adsorption model
UK
United Kingdom (UK)
UKDoE
United Kingdom Department of the Environment
UNSCEAR
United Nations Scientific Committee on the Effects of Atomic Radiation
A.3
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A.2.0 List of Symbols for the Elements and Corresponding Names
Symbol
Element
Symbol
Element
Symbol
Element
Ac
Actinium
Gd
Gadolinium
Po
Polonium
Ag
Silver
Ge
Germanium
Pr
Praseodymium
A1
Aluminum
H
Hydrogen
Pt
Platinum
Am
Americium
He
Helium
Pu
Plutonium
Ar
Argon
Hf
Hafnium
Ra
Radium
As
Arsenic
Hg
Mercury
Rb
Rubidium
At
Astatine
Ho
Holmium
Re
Rhenium
Au
Gold
I
Iodine
Rh
Rhodium
B
Boron
In
Indium
Rn
Radon
Ba
Barium
Ir
Iridium
Ru
Ruthenium
Be
Beryllium
K
Potassium
S
Sulfur
Bi
Bismuth
Kr
Krypton
Sb
Antimony
Bk
Berkelium
La
Lanthanum
Sc
Scandium
Br
Bromine
Li
Lithium
Se
Selenium
C
Carbon
Lu
Lutetium
Si
Silicon
Ca
Calcium
Lw
Lawrencium
Sm
Samarium
Cb
Columbium
Md
Mendelevium
Sn
Tin
Cd
Cadmium
Mg
Magnesium
Sr
Strontium
Ce
Cerium
Mn
Manganese
Ta
Tantalum
Cf
Californium
Mo
Molybdenum
Tb
Terbium
CI
Chlorine
N
Nitrogen
Tc
Technetium
Cm
Curium
Na
Sodium
Te
Tellurium
Co
Cobalt
Nb
Niobium
Th
Thorium
Cr
Chromium
Nd
Neodymium
Ti
Titanium
Cs
Cesium
Ne
Neon
T1
Thallium
Cu
Copper
Ni
Nickel
Tm
Thulium
Dy
Dysprosium
No
Nobelium
U
Uranium
Er
Erbium
Np
Neptunium
V
Vanadium
Es
Einsteinium
0
Oxygen
w
Tungsten
Eu
Europium
Os
Osmium
w
Wolfram
F
Fluorine
P
Phosphorus
Xe
Xenon
Fe
Iron
Pa
Protactinium
Y
Yttrium
Fm
Fermium
Pb
Lead
Yb
Ytterbium
Fr
Francium
Pd
Palladium
Zn
Zinc
Ga
Gallium
Pm
Promethium
Zr
Zirconium
A.4
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A.3.0 List of Symbols and Notation
Pb
Porous media bulk density (mass/length3)
A
Angstrom, 10"10 meters
ads
Adsorption or adsorbed
Ai
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
IAP
Ion activity product
IEP
Isoelectric point
Kd
Concentration-based partition (or distribution) coefficient
Kfi298
Equilibrium constant at 298 K
K T
A^r,T
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, 10"12 Curies
pE
Negative common logarithm of the free-electron activity
pH
Negative logarithm of the hydrogen ion activity
pHzpc
pH for zero point of charge
ppm
Parts per million
R
Ideal gas constant, 1.9872 cal/mol-K
A. 5
-------
Rf Retardation factor
s Solid phase species
sec Second
SI Saturation index, as defined by log (IAP/Kr X)
T Absolute temperature, usually in Kelvin unless otherwise specified
t Time
t1/2 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
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APPENDIX B
Definitions
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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% to -90%.
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
CaX(s) + 90Sr2+ (aq) = 90SrX(s) + Ca2+ (aq)
approximately stoichiometric and can be written, for example, as
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.
Clay Content - particle size fraction of soil that is less than 2 |im (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 isotopes 2H of hydrogen.
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+ + 2H20 = 2Pu3+ + PuOf +4H+
3PuC>2 + 4H+ = Pu3+ + 2PuOf +2H20.
Electron Activity - unity for the standard hydrogen electrode.
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
2 or more new substances. For example, the first hydrolysis reaction of U4+ can
be written as
U4+ + H20 = U0H3++H+.
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 35-65% clay and 65-35% 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.
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 1 center cation (as
compared to a polymeric species).
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.
Polynuclear Species - an aqueous species containing more than 1 central cation
moiety, e.g., (U02)2C03(0H)3 and Pb^OH)^.
Protium (H) - stable isotope !H of hydrogen.
Retrograde Solubility - solubility that decreases with increasing temperature, such
as those of calcite (CaC03) 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 of 'H (protium).
In natural water that produces 7.2 x 10"3 disintegrations per minute per milliliter
(dpm/ml) of tritium, 1 TU is approximately equal to 3.2 picocuries/milliliter
(pCi/ml).
-------
APPENDIX C
Partition Coefficients For Cadmium
-------
Appendix C
Partition Coefficients For Cadmium
C.1.0 Background
Cadmium Kd values and some important ancillary parameters that have been shown to influence
cadmium sorption were collected from the literature and tabulated. Data included in this data set
were from studies that reported Kd values and were conducted in systems consisting of
• Natural soils (as opposed to pure mineral phases)
• Low ionic strength solutions (<0.1 M)
• pH values between 4 and 10
• Solution cadmium concentration less than 10-5 M
• Low humic materials concentrations (<5 mg/1)
• No organic chelates (such as EDTA)
A total of 174 cadmium Kd values were found in the literature (see summary in Section C.3.0).
At the start of the literature search, attempts were made to identify Kd studies that included
ancillary data on aluminum/iron-oxide concentrations, calcium and magnesium solution
concentrations, pH, cation exchange capacity (CEC), clay content, redox status, organic matter
concentrations and sulfide concentrations. Upon reviewing the data and determining the
availability of cadmium Kd measurements having ancillary information, Kd values were collected
that included information on clay content, pH, CEC, total organic carbon (related to organic
matter), and dissolved cadmium concentrations. The selection of these parameters was based on
availability of data and the possibility that the parameter may impact cadmium Kd values. Of the
174 cadmium Kd values included in our tabulation, 62 values had associated clay content data,
174 values had associated pH data, 22 values had associated CEC data, 63 values had total
organic carbon data, 172 values had associated cadmium concentration data, and 16 had
associated aluminum/iron-oxide data. The descriptive statistics for this total set of cadmium Kd
values are listed in Table C.l.
C.2
-------
Table C.l. Descriptive statistics of the cadmium Kd data set for soils.
Cadmium
K,,
(ml/g)
Clay
Content
(wt.%)
pH
CEC
(meq/lOOg)
TOC
(mg/1)
Cd Cone.
(mg/1)
Fe Oxides
(wt.%)
Mean
226.7
14.2
5.88
21
5.5
3.67
1.32
Standard
Error
44.5
1.7
0.09
3
0.85
0.48
0.53
Median
121.8
10.24
5.83
23
2.0
0.01
0.38
Mode
80.0
6
6.8
2
0.4
0.01
0.19
Std. Dev
586.6
13.5
1.16
15
6.8
6.27
2.12
Sample
Variance
344086
182
1.34
245
45.9
39.4
4.51
Range
4359
86.2
6.20
58
32.4
34.9
8.28
Minimum
0.50
.9
3
2
0.2
0.01
0.01
Maximum
4360
87.1
9.2
60
32.6
35
8.29
No. Samples
174
62
174
22
63
172
16
C.2.0 Approach and Regression Models
C.2.1 Correlations with Cadmium Kd Values
Linear regression analyses were conducted between the ancillary parameters and cadmium Kd
values. The correlation coefficients from these analyses are presented in Table C.2. These
results were used for guidance for selecting appropriate independent variables to use in the
look-up table. The largest correlation coefficient was between pH and log(Kd). This value is
significant at the 0.001 level of probability. Attempts at improving this correlation coefficient
through the use of additional variables, i.e., using multiple-regression analysis, were not
successful. Multiple regression analyses were conducted with the following pairs of variables to
predict cadmium Kd values: total organic carbon and pH, clay content and pH, total organic
carbon and iron-oxides, and pH and CEC.
C.3
-------
Table C.2. Correlation coefficients (r) of the cadmium Kd data set for soils.
Cadmium
K,
l°g (Kd)
Clay
Content
pH
CEC
TOC
Cd Cone.
Cadmium
Kd
1
log (BCJ
0.69
1
Clay Cone.
-0.04
0.03
1
pH
0.50
0.75
0.06
1
CEC
0.40
0.41
0.62
0.35
1
TOC
0.20
0.06
0.13
-0.39
0.27
1
Cd Cone.
-0.02
-0.10
-0.39
0.22
-0.03
-0.09
1
Fe Oxide
Cone.
0.18
0.11
-0.06
0.16
0.19
0.18
0.01
C.2.2 Cadmium Kd Values as a Function of pH
The cadmium Kd values plotted as a function of pH are presented in Figure C. 1. A large amount
of scatter exists in these data. At any given pH, the range of Kd values may vary by 2 orders of
magnitude. This is not entirely satisfactory, but as explained above, using more than 1 variable
to help categorize the cadmium Kd values was not fruitful.
The look-up table (Table C.3) for cadmium Kd values was categorized by pH. The regression
equation for the line presented in Figure C.l is:
Cd Kd = -0.54 + 0.45(pH). (C.l)
The minimum and maximum values were estimated based on the scatter of data points observed
in Figure C.l.
C.4
-------
Figure C.l. Relation between cadmium Kd values and pH in soils.
Table C.3. Look-up table for estimated range of Kd values for cadmium based on pH.
[Tabulated values pertain to systems consisting of natural soils (as opposed
to pure mineral phases), low ionic strength (< 0.1 M), low humic material
concentrations (<5 mg/1), no organic chelates (such as EDTA), and
oxidizing conditions.]
pH
K„ (ml/g)
3-5
5-8
8-10
Minimum
1
8
50
Maximum
130
4,000
12,600
C.5
-------
C.3.0 Data Set for Soils
Table C.4 lists the available Kd values for cadmium identified for experiments conducted with
only soils. The Kd values are listed with ancillary parameters that included clay content, pH,
CEC, TOC, solution cadmium concentrations, and iron-oxide concentrations
Table C.4. Cadmium Kd data set for soils.
CdK,
(ml/g)
Clay
Cont.
(wt%)
PH
CEC
(meq/
100 g)
TOC
(wt%)
[Cd]
(mg/1)
Fe
Oxides
(wt.%)
Solution
Soil
Identification
Comments
Ref.a
52.5
54.7
4.8
30.2
1.54
1
0.33
0.005 M
CaN03
Alligator Ap
Converted
Freund. to Kd
Using lppm
1
288.4
8.3
5.7
2
0.61
1
0.1
0.005 M
CaN03
Cecil Ap
Converted
Freund. to Kd
Using lppm
1
13.9
51.2
5.4
2.4
0.26
1
0.08
0.005 M
CaN03
Cecil B
Converted
Freund. to Kd
Using lppm
1
186.6
0.9
5.9
22.54
6.62
1
1.68
0.005 M
CaNOj
Kula Apl
Converted
Freund. to Kd
Using lppm
1
52.7
17.6
3.9
26.9
11.6
1
1.19
0.005 M
CaN03
Lafitte Ap
Converted
Freund. to Kd
Using lppm
1
91.2
28.2
6
11
1.67
1
0.19
0.005 M
CaNOj
Molokai Ap
Converted
Freund. to Kd
Using lppm
1
28.8
2.8
6.9
4.1
0.21
1
0.06
0.005 M
CaN03
Norwood Ap
Converted
Freund. to Kd
Using lppm
1
97.9
6.2
6.6
8.6
0.83
1
0.3
0.005 M
CaNOj
Olivier Ap
Converted
Freund. to Kd
Using lppm
1
5.5
3.8
4.3
2.7
1.98
1
0
0.005 M
CaN03
Spodisol
Converted
Freund. to Kd
Using lppm
1
755.1
23.9
7.6
48.1
4.39
1
0.19
0.005 M
CaNOj
Webster Ap
Converted
Freund. to Kd
Using lppm
1
C.6
-------
CdKa
(ml/g)
Clay
Cont.
(wt%)
PH
CEC
(meq/
10° R)
TOC
(wt%)
[Cd]
(mg/1)
Fe
Oxides
(wt.%)
Solution
Soil
Identification
Comments
Ref.a
14.4
2.8
5.3
2
2.03
1
0.42
0.005 M
CaN03
Windsor Ap
Converted
Freund. to Kd
Using lppm
1
87.1
8.4
60
1.44
1
1.07
Water
Vertic
Torrifluvent
Converted
Freund. to Kd
Using lppm
2
33.88
5.2
33.8
32.6
1
Water
Organic
Converted
Freund. to Kd
Using lppm
2
20.42
5.8
23.8
3
1
8.29
Water
Boomer, Ultic
Haploxeralf
Converted
Freund. to Kd
Using lppm
2
10.47
6
25
3.2
1
1.07
Water
UlticPalexeralf
Converted
Freund. to Kd
Using lppm
2
80
8.2
8.2
0.21
35
0.01 M
NaCl
Gevulot
Calc. Fig 1.
3
200
7.8
15.4
0.83
25
0.01 M
NaCl
Bet Yizhaq
Calc. Fig 1.
3
133.3
8.3
18.9
0.23
30
0.01 M
NaCl
Gilat
Calc. Fig 1.
3
181.8
7.6
31.8
0.79
25
0.01 M
NaCl
Maaban
Michael
Calc. Fig 1.
3
266.7
7.9
37
0.86
15
0.01 M
NaCl
Hahoterim
Calc. Fig 1.
3
8
8
3.7
1.6
11.2
0.01 M
NaN03
Downer
Loamy Sand
4
17
8
4.8
1.6
11.2
0.01 M
NaN03
Downer
Loamy Sand
4
32
8
5.3
1.6
11.2
0.01 M
NaNOj
Downer
Loamy Sand
4
64
8
6
1.6
11.2
0.01 M
NaN03
Downer
Loamy Sand
4
92
8
6.2
1.6
11.2
0.01 M
NaNOj
Downer
Loamy Sand
4
110
8
6.8
1.6
11.2
0.01 M
NaN03
Downer
Loamy Sand
4
250
8
7.3
1.6
11.2
0.01 M
NaNOj
Downer
Loamy Sand
4
C.7
-------
CdKa
(ml/g)
Clay
Cont.
(wt%)
PH
CEC
(meq/
10° R)
TOC
(wt%)
[Cd]
(mg/1)
Fe
Oxides
(wt.%)
Solution
Soil
Identification
Comments
Ref.a
580
8
8.5
1.6
11.2
0.01 M
NaN03
Downer
Loamy Sand
4
0.5
6
3.1
0.4
11.2
0.01 M
NaN03
Freehold
Sandy Loam A
Horizon
4
3.3
6
3.8
0.4
11.2
0.01 M
NaN03
Freehold
Sandy Loam A
Horizon
4
7.5
6
4.5
0.4
11.2
0.01 M
NaNOj
Freehold
Sandy Loam A
Horizon
4
10
6
5.5
0.4
11.2
0.01 M
NaN03
Freehold
Sandy Loam A
Horizon
4
34
6
6.1
0.4
11.2
0.01 M
NaNOj
Freehold
Sandy Loam A
Horizon
4
45
6
6.8
0.4
11.2
0.01 M
NaN03
Freehold
Sandy Loam A
Horizon
4
80
6
7.5
0.4
11.2
0.01 M
NaNOj
Freehold
Sandy Loam A
Horizon
4
150
6
8
0.4
11.2
0.01 M
NaN03
Freehold
Sandy Loam A
Horizon
4
420
6
8.4
0.4
11.2
0.01 M
NaNOj
Freehold
Sandy Loam A
Horizon
4
900
6
9.1
0.4
11.2
0.01 M
NaN03
Freehold
Sandy Loam A
Horizon
4
2.1
13
3
16.8
11.2
0.01 M
NaNOj
Boonton Loam
4
10
13
3.7
16.8
11.2
0.01 M
NaN03
Boonton Loam
4
30
13
4.2
16.8
11.2
0.01 M
NaNOj
Boonton Loam
4
57
13
4.6
16.8
11.2
0.01 M
NaN03
Boonton Loam
4
C.8
-------
CdKa
(ml/g)
Clay
Cont.
(wt%)
PH
CEC
(meq/
10° R)
TOC
(wt%)
[Cd]
(mg/1)
Fe
Oxides
(wt.%)
Solution
Soil
Identification
Comments
Ref.a
101
13
5
16.8
11.2
0.01 M
NaN03
Boonton Loam
4
195
13
5.2
16.8
11.2
0.01 M
NaN03
Boonton Loam
4
420
13
5.8
16.8
11.2
0.01 M
NaN03
Boonton Loam
4
1,200
13
6.2
16.8
11.2
0.01 M
NaNOj
Boonton Loam
4
4,000
13
6.8
16.8
11.2
0.01 M
NaN03
Boonton Loam
4
1.2
16
3.3
9.8
11.2
0.01 M
NaNOj
Rockaway
Stony Loam
4
7.1
16
4.1
9.8
11.2
0.01 M
NaN03
Rockaway
Stony Loam
4
27
16
4.8
9.8
11.2
0.01 M
NaNOj
Rockaway
Stony Loam
4
53
16
5.1
9.8
11.2
0.01 M
NaN03
Rockaway
Stony Loam
4
170
16
5.6
9.8
11.2
0.01 M
NaNOj
Rockaway
Stony Loam
4
300
16
6.1
9.8
11.2
0.01 M
NaN03
Rockaway
Stony Loam
4
390
16
6.2
9.8
11.2
0.01 M
NaNOj
Rockaway
Stony Loam
4
910
16
6.5
9.8
11.2
0.01 M
NaN03
Rockaway
Stony Loam
4
1,070
16
6.8
9.8
11.2
0.01 M
NaNOj
Rockaway
Stony Loam
4
43
10
4.8
2.4
11.2
0.01 M
NaN03
Fill Material -
Delaware
River
4
67
10
5.7
2.4
11.2
0.01 M
NaNOj
Fill Material -
Delaware
River
4
130
10
6.3
2.4
11.2
0.01 M
NaN03
Fill Material -
Delaware
River
4
C.9
-------
CdKa
(ml/g)
Clay
Cont.
(wt%)
PH
CEC
(meq/
10° R)
TOC
(wt%)
[Cd]
(mg/1)
Fe
Oxides
(wt.%)
Solution
Soil
Identification
Comments
Ref.a
150
10
6.7
2.4
11.2
0.01 M
NaN03
Fill Material -
Delaware
River
4
370
10
7.3
2.4
11.2
0.01 M
NaN03
Fill Material -
Delaware
River
4
880
10
8
2.4
11.2
0.01 M
NaN03
Fill Material -
Delaware
River
4
1,950
10
9.2
2.4
11.2
0.01 M
NaNOj
Fill Material -
Delaware
River
4
1,000
12
8
1
3.7
Carbonate
Groundwate
r
Interbed
pH of
Groundwater
5
4,360
12.4
8
1
2.5
Carbonate
Groundwate
r
Alluvium
pH of
Groundwater
5
536.8
25.2
6.8
27.5
0.01 M
NaCl
Soil A
Desorption
6
440
25.2
6.8
27.5
0.01 M
NaCl
Soil A
Desorption
6
9
4.3
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
23.4
4.3
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
15.8
4.4
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
11.3
4.5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
31.2
4.5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
32.5
4.5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
23
4.5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
17.1
4.7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
13.1
4.8
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
C.10
-------
CdKa
(ml/g)
Clay
Cont.
(wt%)
PH
CEC
(meq/
10° R)
TOC
(wt%)
[Cd]
(mg/1)
Fe
Oxides
(wt.%)
Solution
Soil
Identification
Comments
Ref.a
24.9
4.6
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
26.8
4.7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
36.2
4.7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
32.9
4.7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
37.2
4.7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
29.2
4.8
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
28.3
4.8
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
22.6
4.9
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
37.4
4.9
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
40.9
4.9
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
63.5
4.7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
25.2
5.4
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
29.9
5.3
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
33.7
5.2
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
44.3
5.1
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
42.8
5.1
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
53.5
5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
56.2
4.9
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
68.7
5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
C.ll
-------
CdKa
(ml/g)
Clay
Cont.
(wt%)
PH
CEC
(meq/
10° R)
TOC
(wt%)
[Cd]
(mg/1)
Fe
Oxides
(wt.%)
Solution
Soil
Identification
Comments
Ref.a
82.3
5.1
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
75.7
5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
95.2
4.8
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
103
4.8
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
160
4.8
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
43.3
5.4
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
55.2
5.4
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
52.2
5.3
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
40.3
5.6
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
56.1
5.5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
67.5
5.5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
102.9
5.4
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
164.4
5.5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
163.8
5.3
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
202.1
5.2
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
172.4
5.2
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
149
5.2
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
72.8
5.6
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
81.6
5.7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
C.12
-------
CdKa
(ml/g)
Clay
Cont.
(wt%)
PH
CEC
(meq/
10° R)
TOC
(wt%)
[Cd]
(mg/1)
Fe
Oxides
(wt.%)
Solution
Soil
Identification
Comments
Ref.a
90
5.7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
94.3
5.6
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
48.1
6.2
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
56.5
6.4
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
81
6.5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
122.3
6.4
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
121.4
6.2
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
101.5
6
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
99.3
6
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
107.8
6
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
219.5
6.2
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
179.2
6.2
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
177
6.1
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
360.4
6
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
305.2
6
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
236.8
5.9
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
186.3
5.9
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
174.8
5.8
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
138.7
5.8
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
C.13
-------
CdKa
(ml/g)
Clay
Cont.
(wt%)
PH
CEC
(meq/
10° R)
TOC
(wt%)
[Cd]
(mg/1)
Fe
Oxides
(wt.%)
Solution
Soil
Identification
Comments
Ref.a
132.5
5.7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
375.6
5.9
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
403.3
5.8
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
510.8
5.8
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
225.9
5.7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
227.3
5.7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
248
5.7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
253.1
5.6
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
277.2
5.6
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
240.7
6.4
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
227.8
6.5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
281.1
6.6
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
551.2
6.2
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
519.8
6.2
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
418.7
6.2
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
353.7
6.2
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
400.8
6.4
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
609.2
6.3
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
545.7
6.3
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
C.14
-------
CdKa
(ml/g)
Clay
Cont.
(wt%)
PH
CEC
(meq/
10° R)
TOC
(wt%)
[Cd]
(mg/1)
Fe
Oxides
(wt.%)
Solution
Soil
Identification
Comments
Ref.a
515.9
6.4
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
545.7
6.4
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
760.9
6.4
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
665.7
6.5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
503.2
6.5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
515.2
7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
488.9
6.9
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
481
6.9
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
461.6
6.9
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
1,151
6.5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
868.7
6.6
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
637.2
6.7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
970.9
6.7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
950.5
6.8
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
886.2
6.9
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
1,106
6.9
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
970.9
7
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
2,248
7.1
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
1,909
7.2
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
C.15
-------
CdKa
(ml/g)
Clay
Cont.
(wt%)
PH
CEC
(meq/
10° R)
TOC
(wt%)
[Cd]
(mg/1)
Fe
Oxides
(wt.%)
Solution
Soil
Identification
Comments
Ref.a
1,411
7.3
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
1,383
7.4
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
2,337
7.5
0.01
0.001M
CaCl2
Agricultural
Danish Soil
Co = 0.7 to
12.6 ppb
7
a 1 = Buchter etal., 1989; 2 = Garcia-Miragaya, 1980; 3 = Navrot etal., 1978; 4 = Allen etal., 1995; 5 = Del Debbio,
1991; 6 = Madrid etal., 1992; 7 = Anderson and Christensen , 1988
C.4.0 References
Allen, G. E., Y. Chen, Y, Li, and C. P. Huang. 1995. "Soil Partition Coefficients for Cd by
Column Desorption and Comparison to Batch Adsorption Measurements." Environmental
Science and Technology, 29:1887-1891.
Anderson, P. R., and T. H. Christensen. 1988. "Distribution Coefficients of Cd, Co, Ni, and Zn
in Soils." Journal of Soil Science, 39:15-22.
Buchter, B., B. Davidoff, M. C. Amacher, C. Hinz, I. K. Iskandar, and H. M. Selim. 1989.
"Correlation of Freundlich Kd and n Retention Parameters with Soils and Element." Soil
Science, 148:370-379.
Del Debbio, J. A. 1991. "Sorption of Strontium, Selenium, Cadmium, and Mercury in Soil."
RadiochimicaActa, 52/53:181-186.
Garcia-Miragaya, J. 1980. "Specific Sorption of Trace Amounts of Cadmium by Soils."
Communications in Soil Science and Plant Analysis, 11:1157-1166.
Madrid, L., and E. Diz-Barrientos. 1992. "Influence of Carbonate on the Reaction of Heavy
Metals in Soils." Journal of Soil Science, 43:709-721.
Navrot, J., A. Singer, and A. Banin. 1978. "Adsorption of Cadmium and its Exchange
Characteristics in Some Israeli Soils." Journal of Soil Science, 29:205-511.
C.16
-------
APPENDIX D
Partition Coefficients For Cesium
-------
Appendix D
Partition Coefficients For Cesium
D.1.0 Background
Three generalized, simplifying assumptions were established for the selection of cesium Kd
values for the look-up table. These assumptions were based on the findings of the literature
reviewed we conducted on the geochemical processes affecting cesium sorption. The
assumptions are as follows:
• Cesium adsorption occurs entirely by cation exchange, except when mica-like minerals
are present. Cation exchange capacity (CEC), a parameter that is frequently not
measured, can be estimated by an empirical relationship with clay content and pH.
• Cesium adsorption onto mica-like minerals occurs much more readily than desorption.
Thus, Kd values, which are essentially always derived from adsorption studies, will
greatly overestimate the degree to which cesium will desorb from these surfaces.
• Cesium concentrations in groundwater plumes are low enough, less than approximately
10"7 M, such that cesium adsorption follows a linear isotherm.
These assumptions appear to be reasonable for a wide range of environmental conditions.
However, these simplifying assumptions are clearly compromised in systems with cesium
concentrations greater than approximately 10"7 M , ionic strengths greater than about 0.1 M, and
pH values greater than about 10.5. These assumptions will be discussed in more detail in the
following sections.
Based on the assumptions and limitation described above, cesium Kd values and some important
ancillary parameters that influence cation exchange were collected from the literature and
tabulated. Data included in this table were from studies that reported Kd values (not percent
adsorbed or Freundlich or Langmuir constants) and were conducted in systems consisting of:
• Low ionic strength (< 0.1 M)
• pH values between 4 and 10.5
• Dissolved cesium concentrations less than 10"7 M
• Low humic material concentrations (<5 mg/1)
• No organic chelates (e.g., EDTA)
The ancillary parameters included in these tables were clay content, mica content, pH, CEC,
surface area, and solution cesium concentrations. This cesium data set included 176 cesium Kd
values.
D.2
-------
Two separate data sets were compiled. The first one (see Section D.3) included both soils and
pure mineral phases. The lowest cesium Kd value was 0.6 ml/g for a measurement made on a
system containing a soil consisting primarily of quartz, kaolinite, and dolomite and an aqueous
phase consisting of groundwater with a relatively high ionic strength (I « 0.1 M) (Lieser el al.,
1986) (Table D. 1). The value is unexplainably much less than most other cesium Kd values
present in the data set. The largest cesium Kd values was 52,000 ml/g for a measurement made
on a pure vermiculite solid phase (Tamura, 1972). The average cesium Kd value was 2635 ±
530 ml/g.
Table D.l. Descriptive statistics of cesium Kd data set including soil and pure mineral
phases. [Data set is presented in Section D.3.]
Kd (ml/g)
Clay
(%)
Mica
(%)
pH
CEC
(meq/100 g)
Surface Area
(m2/g)
Mean
2,635
30
5.5
7.4
30.4
141.3
Standard Error
530
3.8
0.7
0.1
3.7
29.7
Median
247
42
4
8.2
4.8
31.2
Mode
40
42
4
8.2
1.8
17.7
Standard Deviation
7055
15
4.4
1.7
37.4
230.4
Sample Variance
49,781,885
226
20.0
2.8
1,396.9
53,106
Range
51,999
38
13
7.8
129.9
638
Minimum
0.6
4
2
2.4
0.00098
8
Maximum
52,000
42
15
10.2
130
646
No. Observations
177
15
41
139
103
60
Confidence Level
(95.0%)
1,046.6
8.3
1.4
0.3
7.3
59.5
D.3
-------
A second data set (see Section D.4) was created using only data generated from soil studies, that
is, data from pure mineral phases, and rocks, were eliminated from the data set. Descriptive
statistics of the soil-only data set are presented in Table D.2. Perhaps the most important finding
of this data set is the range and median1 of the 57 Kd values. Both statistics decreased
appreciably. In the soil-only data set, the median was 89 ml/g. The median is perhaps the single
central estimate of a cesium Kd value for this data set. The range of Kd values was from 7.1
ml/g, for a measurement made on a sandy carbonate soil (Routson et al., 1980), to 7610 ml/g for
a measurement made on another carbonate soil containing greater than 50 percent clay and silt
(Serne el al., 1993). Interestingly, these 2 soils were both collected from the U.S. Department of
Energy's Hanford Site in eastern Washington state.
Table D.2. Descriptive statistics of data set including soils only. [Data set is presented
in Section D.4.]
Cesium
Kd
(ml/g)
Clay
(%)
Mica
(%)
pH
CEC
(meq/lOOg)
Surface Area
(m2/g)
Mean
651
5
5.6
6.9
34
57.5
Standard Error
188
0.6
0.6
0.3
8.9
13.4
Median
89
5.0
4
6.7
20
60
Mode
22
NA
4
4.0
60
70
Standard Deviation
1423
1.0
4.3
1.9
29.5
44.6
Sample Variance
2026182
1.0
18.4
3.6
870
1986
Range
7602
2.0
13
7.8
57.4
123.4
Minimum
7.1
7.1
2
2.4
2.6
6.6
Maximum
7610
6.0
15
10.2
70.0
130
No. Observations
57
3
45
55
11
11
Confidence Level (95%)
378
2.5
1.29
0.5
19.8
30
1 The median is that value for which 50 percent of the observations, when arranged in order of
magnitude, lie on each side.
D.4
-------
The soil-only data set was frequently incomplete with regard to supporting data describing the
experimental conditions under which the cesium Kd values were measured (Table D.2). Quite
often the properties of the solid phase or the dissolved cesium concentration used in the Kd
experiments were not reported. For instance, there were only 3 cesium Kd values that had
accompanying clay content data, 11 cesium Kd values that had accompanying cation exchange
data, and 11 cesium Kd values that had accompanying surface area data (Table D.2).
Consequently, it was not possible to evaluate adequately the relationship between cesium Kd
values and these important, independent soil parameters. This is discussed in greater detail
below.
D.2.0 Approach and Regression Models
D.2.1 Correlations with Cesium Kd Values
A matrix of the correlation coefficients for the parameters included in the data set containing Kd
values determined in experiments with both soils and pure mineral phases is presented in
Table D.3. The correlation coefficients that are significant at or less than the 5 percent level of
probability (P < 0.05) are identified with a footnote. The parameter with the largest correlation
coefficient with cesium Kd was CEC (r = 0.52). Also significant was the correlation coefficient
between cesium Kd values and surface area (r = 0.42) and CEC and clay content (r = 0.64). The
poor correlation between cesium aqueous concentration ([Cs]aq) and cesium Kd values can be
attributed to the fact that the former parameter included concentration of the solution prior and
after contact with the soils. We report both under the same heading, because the authors
frequently neglected to indicate which they were reporting. More frequently, the spike
concentration (the cesium concentration prior to contact with the soil) was reported, and this
parameter by definition is not correlated to Kd values as well as the concentrations after contact
with soil (the denominator of the Kd term).
A matrix of the correlation coefficients for the parameters included in the data set containing Kd
values determined in experiments with only soils is presented in Table D.4. As mentioned above
(Table D.2), the reports in which soil was used for the Kd measurements tended to have little
supporting data about the aqueous and solid phases. Consequently, there was little information
for which to base correlations. This occasionally resulted in correlations that were not
scientifically meaningful. For example, the correlation between CEC and cesium Kd was -0.83,
for only 11 observations (10 degrees of freedom). The negative sign of this correlation
contradicts commonly accepted principles of surface chemistry.
D.5
-------
Table D.3. Correlation coefficients (r) of the cesium Kd value data set that
included soils and pure mineral phases. [Data set is presented in
Section D.3.]
Cesium
K„
Clay
Content
Mica
pH
CEC
Surface Area
Cesium Kd
1.00
Clay Content
0.05
1.00
Mica
0.29
0.00
1.00
pH
0.10
-0.11
0.08
1.00
CEC
0.52a
0.64a
NA
0.37
1.00
Surface Area
0.42a
0.35
NA
-0.11
0.47a
1.00
[Cs]aq
-0.07
0.85a
0.29
0.13
-0.17
-0.15
a Correlation coefficient is significant at the 5% level of significance (P < 0.05).
Table D.4. Correlation coefficients (r) of the soil-only data set. [Data set is
presented in Section D.4.]
Cesium
Kh
Clay
Content
Mica
pH
CEC
Surface Area
Cesium Kd
1.00
Clay Content
-0.21
1.00
Mica
0.27
0
1.00
pH
0.11
0.4
0.07
1.00
CEC
-0.83
NA
0.991
0.05
1.00
Surface Area
-0.31
NA
0.991
-0.03
0.37
1.00
[Cs]aq
0.18
NA
0.09
-0.04
0.00
0
1 Correlation coefficient is significant at >5% level of significance (P < 0.05).
D.6
-------
The high correlations between mica concentrations and CEC (r = 0.99) and mica concentrations
and surface area (r = 0.99) are somewhat misleading in the fact that both correlations represent
only 4 data points collected from 1 study site in Fontenay-aux-Roses in France (Legoux et al.,
1992).
D.2.2 Cesium Adsorption as a Function of CEC andpH
Akiba and Hashimoto (1990) showed a strong correlation between cesium Kd values and the
CEC of a large number of soils, minerals, and rock materials. The regression equation generated
from their study was:
A similar regression analysis using the entire data set (mineral, rocks, and soils) is presented in
Figure D.l.
log (Cs Kd) =1.2+1.0 log (CEC)
(D.l)
6
5 _ y = 2,09+ 0,73x,r = 0.60
o
0
o
-1
-1.5 -1 -0.5 0 0.5 1 1.5 2 2.5
log CEC (meq/100 g)
Figure D.l. Relation between cesium Kd values and CEC.
D.7
-------
By transposing the CEC and cesium Kd data into logarithms, the regression correlation slightly
increases from 0.52 (Table D.3) to 0.60 (Figure D. 1). However, a great amount of scatter in the
data can still be seen in the logarithmic transposed data. For instance, at log(CEC) of 0.25, the
cesium Kd values range over 4 orders of magnitude. It is important to note that the entire cesium
Kd data set only varies 5 orders of magnitude. Thus, the correlation with CEC, although the
strongest of all the independent variables examined, did not reduce greatly the variability of
possible cesium Kd values.
D.2.3 CEC as a Function of Clay Content and pH
Because CEC values are not always available to contaminant transport modelers, an attempt was
made to use independent variables more commonly available in the regression analysis.
Multiple regression analysis was conducted using clay content and pH as independent variables
to predict CEC values (Figure D.2). Clay content was highly correlated to CEC (r = 0.64). Soil
pH was not significantly correlated to either CEC or cesium Kd values.
50
40
60
O
O
O
W
O
10
1 1 o 1
1 I I
y = 4,1 + 0,44x, r
= 0.63
o
o
1 1 o
8
o
V
• i
- O O
° $ 00
o sfistfo
- As O
JPo O
fill o i i
- A
O o
o
' 1
10 20 30 40 50
Clay (%)
60 70
Figure D.2. Relation between CEC and clay content.
D.8
-------
D. 2.4 Cesium Adsorption onto Mica-Like Minerals
Cesium adsorption onto mica-like minerals has long been recognized as a non-reversible reaction
(Bruggenwert and Kamphorst, 1979; Comans et al., 1989; Cremers et al., 1988; Douglas, 1989;
Evans et al., 1983; Francis andBrinkley, 1976; Sawhney, 1972; Smith and Comans, 1996;
Tamura, 1972). This is an important property in adsorption reactions because 1 of the
assumptions in applying the Kd model to describe adsorption is that the rate at which adsorption
occurs is equal to the rate at which desorption occurs. This phenomena is referred to as an
adsorption hysteresis. Cesium adsorption onto mica-like minerals is appreciably faster than its
desorption. The reason for this is that the cesium ion fits perfectly into the hexagonal ring
formed on the tetrahedral sheet in the crystallographic structure of mica-like clays. This perfect
fit does not permit other cations that exist at much greater concentrations in nature to exchange
the cesium from these sites. This can be demonstrated using the data of Tamura (1972)
(Table D.5). He measured cesium Kd values for mica, vermiculite, and kaolinite using a water
and 0.1 M NaCl background solution. For mica, the Kd value remained about the same for both
solutions. For the vermiculite and kaolinite, the cesium Kd values greatly decreased when the
higher ionic strength solution was used. This indicates that the sodium, which existed at 11
orders of magnitude higher concentration than the cesium could out compete the adsorption of
cesium on the vermiculite and kaolinite but not on the mica. Another point of interest regarding
this data set is that the cesium Kd values do correlate with CEC of these different mineral phases
when water is the background solution. However, when the higher ionic strength solution is
used, the correlation with CEC no longer exists.
Comans et al. (1989) measured cesium Kd values of a mica (Fithian illite) by desorption and
adsorption experiments. Portions of their data are presented in Table D.6. Cesium Kd values
based on desorption experiments are appreciably greater than those measure in adsorption
experiments.
Table D.5. Effect of mineralogy on cesium exchange. [Data are from Tamura
(1972) who used an initial concentration of dissolved cesium of
1.67xl0"12 M.]
Mineral
CEC
in Water
in 0.1 M NaCl
Phases
(meq/100 g)
(ml/g)
(ml/g)
Mica
20
26,000
28,600
Vermiculite
127
52,000
2,700
Kaolinite
11.2
2,500
94
D.9
-------
Table D.6. Cesium Kd values measured on mica (Fithian illite) via adsorption
and desorption experiments. [Data are from Comans et al. (1989).]
Experimental Conditions
Adsorption
Cesium Kd
Desorption
Cesium Kd
K-saturated Mica, 7xl0"6 M Cs
2,890
5,200
K-saturated Mica, 2xl0"7 M Cs
9,000
11,300
Ca-saturated Mica, 7x10"6 M Cs
1,060
4,600
Ca-saturated Mica, 2x10"7 M Cs
600,000
1,050,000
Essentially all Kd values reported in the literature are measured using adsorption experiments.
Thus, in the case of soils containing mica-like soils, using adsorption Kd values will likely
overestimate the degree to which desorption will occur. To account for this difference in
adsorption and desorption, one could artificially increase the Kd values used in a transport code
when cesium is desorbing from contaminated soil.
D.2.5 Cesium Adsorption as a Function of Dissolved Cesium Concentrations
At very low concentrations, the adsorption isotherm for cesium is linear. The linear range varies
dependent on the adsorbing phase and on the background aqueous phase (Akiba et al., 1989;
Sposito, 1989). Table D.7 provides the linear range of some Freundlich adsorption isotherm data
reported in the literature. The upper limit of the linear range varies by several orders of
magnitude depending on the solid phase and aqueous chemistry. The lowest upper limit reported
in Table D.7 is 1 x 10"10 M cesium. This is in fact a rather high concentration when compared to
those found in groundwater plumes. For instance, the highest reported 137Cs concentration in the
groundwaters beneath the Hanford Site in 1994 was 1.94 x 10"13 M (or 2,310 pCi/1) for Well 299
E-28-23 (Hartman and Dresel, 1997). This is several orders of magnitude below the smallest
upper limit reported in Table D.7, suggesting that most far-field radioactive cesium adsorption
likely follows a linear isotherm. The simple Kd value describes a linear isotherm.
D.10
-------
Table D.7. Approximate upper limits of linear range of adsorption isotherms on various
solid phases.
Upper Limit of
Linear Range (M)
Solid Phase
Background
Aqueous Phase
Reference
1 x 10"7
Itado Tuff
Deionized Water
Akidae^a/., 1989
1 xlO"10
Sandstone
Deionized Water
Akidae^a/., 1989
5 x 10"5
Limestone
Deionized Water
Akidae^a/., 1989
1 x 10"10
Augite Andesite
Deionized Water
Akidae^a/., 1989
5 x 10"9
Olivine Basalt
Deionized Water
Akidae^a/., 1989
1 x 10"8
Rokko Granite
Deionized Water
Akidae^a/., 1989
5 x 10"8
Biotite
Deionized Water
Akidae^a/., 1989
5 x 10"7
Albite
Deionized Water
Akidae^a/., 1989
1 x 10"6
K-Feldspar
Deionized Water
Akidae^a/., 1989
1 x 10"1
Unwashed Kaolinite
Distilled Water/pH 10
Adeleye etal., 1994
<1 x 10"5
Ca Montmorillonite
Distilled Water/pH 10
Adeleye etal., 1994
<1 x 10"5
Na Montmorillonite
Distilled Water/pH 10
Adeleye etal., 1994
<1 x 10"5
Na Kaolinite
Distilled Water/pH 10
Adeleye etal., 1994
1 x 10"3
Na Montmorillonite
Distilled Water/pH 4
Adeleye etal., 1994
When a wider range of cesium concentrations are considered, cesium adsorption onto soils and
pure minerals has been reported to be almost without exception a non-linear relationship
(Adeleye et al., 1994; Akiba et al., 1989; Ames et al., 1982; Erten et al., 1988; Konishi etal.,
1988; Lieser and Staunton, 1994; Steinkopff, 1989; Torstenfelt et al.., 1982). Most investigators
have used a Freundlich equation to describe this relationship (Adeleye etal., 1994; Konishi et
al., 1988; Shiao etal., 1979; Staunton, 1994; Torstenfelt et al., 1982). The Freundlich equation
is
Csabsorbed — a (Cssolution) (D.2)
where Csabsorbed and Cssolution are the cesium concentrations adsorbed and in solution, respectively,
and a and b are fitting parameters. A short description of those Freundlich Equation reported in
the literature are presented in Table D.8. The descriptive statistics of the Freundlich Equations
D.ll
-------
reported in Table D.8 are described in Table D.9. A plot of available cesium adsorption versus
equilibrium cesium solution concentration is shown in Figure D.3.
10s
! 1
1
1 1 1
1 1
l <
104
r °
5?
1
103
r ^
! °
o
1
°:
*o
i
o
102
101
1
; °
1
1°
o
o
o
o
8
8
8^
10°
•
1
i i i
¦ i
1
0
0,2
0,4 0.6 0.8
1 1.2 1
,4 1,6
Solution Cs (jamol/l)
Figure D.3. Kd values calculated from an overall literature
Freundlich equation for cesium (Equation D.2).
D.12
-------
Table D.8. Freundlich equations identified in literature for cesium.
a 1
b 1
Range of Solution Cs
Concentration (M)
Experimental
Ref.2
1.7
0.677
Water/Batcombe Sediment
1
3,300
0.909
Water/Denchworth Sediment
1
260
0.841
Water/Tedburn Sediment
1
16
0.749
Water/Teigngrace Sediment
1
12.2
0.745
lxlO"8 to lxlO"12
Water/Batcombe Sediment
1
6,070
0.899
lxlO"8 to lxlO"12
Water/Denchworth Sediment
1
1,290
0.849
lxlO"8 to lxlO"12
Water/Tedburn Sediment
1
163
0.815
lxlO"8 to lxlO"12
Water/Teigngrace Sediment
1
1.23
0.657
lxlO"8 to lxlO"12
CaCl2/Batcombe Sediment
1
0.63
0.659
CaCl2/Batcombe Sediment
1
427
0.814
lxlO"8 to lxlO"12
CaCl2/Denchworth Sediment
1
1.5
0.599
CaCl2/Denchworth Sediment
1
48.1
0.754
lxlO"8 to lxlO"12
CaCl2/Tedburn Sediment
1
17
0.739
CaCl2/Tedburn Sediment
1
5.22
0.702
lxlO"8 to lxlO"12
CaCl2/Teigngrace Sediment
1
4.4
0.716
CaCl2/Teigngrace Sediment
1
0.22
1.1
lxlO"9 to 1.5xl0"2
Bentonite/Water
2
0.017
0.53
lxlO"9 to 1.5xl0"2
Bentonite/Water
2
0.13
1
lxlO"9 to 1.5xl0"2
B entonite/Groundwater
2
0.048
0.67
lxlO"9 to 1.5xl0"2
B entonite/Groundwater
2
5.10xl0"4
0.21
lxlO"9 to 1.5xl0"2
Takadata Loam/Water
2
3.00xl0"3
0.48
lxlO"9 to 1.5xl0"2
Takadata Loam/Groundwater
2
1.30xl0"5
0.013
lxlO"9 to 1.5xl0"2
Hachinohe Loam/W ater
2
2.30xl0"5
0.38
lxlO"9 to 1.5xl0"2
Hachinohe Loam/Groundwater
2
D.13
-------
a1
b 1
Range of Solution Cs
Concentration (M)
Experimental
Ref.2
2.70xl0"4
0.546
lxlO"8 to lxlO"2
Unwashed/Kaolinite/pH 2
3
5.20xl0"4
0.543
lxlO"8 to lxlO"2
Unwashed/Kaolinite/pH 4
3
2.04xl0"3
0.588
lxlO"8 to lxlO"2
Unwashed/Kaolinite/pH 10
3
2.27xl0"3
0.586
lxlO"8 to lxlO"2
Sodium/Kaolinite/pH 2
3
5.04xl0"2
0.723
lxlO"8 to lxlO"2
Sodium/Kaolinite/pH 4
3
3.49xl0"2
0.703
lxlO"8 to lxlO"2
Na/Kaolinite/pH 7
3
0.235
0.821
lxlO"8 to lxlO"2
Na/Kaolinite/pH 10
3
3.03xl0"2
0.804
lxlO"8 to lxlO"2
Ca/Kaolinite/pH 2
3
0.135
0.845
lxlO"8 to lxlO"2
Ca/Kaolinite/pH 4
3
0.247
0.881
lxlO"8 to lxlO"2
Ca/Kaolinite/pH 7
3
8.71xl0"3
0.694
lxlO"8 to lxlO"2
Ca/Kaolinite/pH 10
3
1.02xl0"4
0.503
lxlO"8 to lxlO"2
Na/Montmorillonite/pH 2
3
1.05xl0"2
0.709
lxlO"8 to lxlO"2
Na/Montmorillonite/pH 4
3
3.17xl0"2
0.755
lxlO"8 to lxlO"2
Na/Montmorillonite./pH 7
3
0.224
0.815
lxlO"8 to lxlO"2
Na/Montmorillonite/pH 10
3
0.241
0.839
lxlO"8 to lxlO"2
Ca/Montmorillonite/pH 2
3
0.481
0.897
lxlO"8 to lxlO"2
Ca/Montmorillonite/pH 4
3
1.84
0.938
lxlO"8 to lxlO"2
Ca/Montmorillonite/pH 7
3
0.274
0.82
lxlO"8 to lxlO"2
Ca/Montmorillonite/pH 10
3
3.40xl0"2
0.51
lxlO"7 to lxlO"3
Granite/pH 8.2
4
4.90xl0"2
0.5
lxlO"7 to lxlO"3
Granite/pH 8.2
4
4.00xl0"2
0.5
5
1 Parameters "a" and "b" are fitting parameters in the Freundlich equation.
2 References: 1 = Fukui, 1990; 2 = Konishi etal., 1988; 3 = Adeleye etal., 1994; 4 = Serne et
al., 1993; 5 = Shiao et al., 1979.
D.14
-------
Table D.9. Descriptive statistics of the cesium Freundlich equations (Table D.8)
reported in the literature.
Statistic
a
b
Mean
252
0.696
Standard Error
150.2
0.029
Median
0.222
0.720
Mode
NA
0.815
Standard Deviation
1019
0.198
Sample Variance
1038711
0.039
Range
6070
1.087
Minimum
0.000013
0.013
Maximum
6070
1.1
95% Confidence Level
302
0.059
Using the medians of the a and b parameters from the literature, we come up with the overall
equation:
CsadSorbed — 0.222(C
^solution)
(D 3)
This equation is plotted in Figure D.4. Using Csadsorbed and Cssolution from equation D.3, a Kd value
can be calculated according to equations D.4,
Kd ^ ^adsorbed/C ^solution.
(D.4)
Cesium Kd values calculated from Equations D.3 and D.4 are presented in Figure D.5.
D.15
-------
J I I I I I L
IQ-iS 10-14 10-13 jq-12 10-11 JQ-10 10-9 10-8
Solution Cs (mol/1)
Figure D.4. Generalized cesium Freundlich equation
(Equation D.3) derived from the literature.
1000
5B 100
¦
s
*o
Zfl
o 10
r "
1
1C
-10 jq-9 1q-8 JQ-7 10"« 10"S
Solution Cs (mol/1)
Figure D.5. Cesium Kd values calculated from generalized
Freundlich equation (Equations D.3 and D.4)
derived from the literature.
D.16
-------
D.2.6 Approach to Selecting Kd Values for Look-up Table
Linear regression analyses were conducted with data collected from the literature. These
analyses were used as guidance for selecting appropriate Kd values for the look-up table. The Kd
values used in the look-up tables could not be based entirely on statistical consideration because
the statistical analysis results were occasionally nonsensible. For example, the data showed a
negative correlation between pH and CEC, and pH and cesium Kd values. These trends
contradict well established principles of surface chemistry. Instead, the statistical analysis was
used to provide guidance as to the approximate range of values to use and to identify meaningful
trends between the cesium Kd values and the solid phase parameters. Thus, the Kd values
included in the look-up table were in part selected based on professional judgment. Again, only
low-ionic strength solutions, such as groundwaters, were considered; thus no solution variables
were included.
Two look-up tables containing cesium Kd values were created. The first table is for systems
containing low concentrations {i.e., less than about 5 percent of the clay-size fraction) of mica-
like minerals (Table D.10). The second table is for systems containing high concentrations of
mica-like minerals (Table D. 11). For both tables, the user will be able to reduce the range of
possible cesium Kd values with knowledge of either the CEC or the clay content.
The following steps were taken to assign values to each category in the look-up tables.
A relation between CEC and clay content was established using data presented in this section.
Three CEC and clay content categories were selected. The limits of these categories were
arbitrarily assigned. The central estimates for the <5 percent mica look-up table (Table D. 10)
were assigned using the CEC/cesium Kd equation in Figure D. 1. The central estimates for the >5
percent mica look-up table (Table D. 11) were assigned by multiplying the central estimates from
Table D.10 by a factor of 2.5. The 2.5 scaler was selected based on relationships existing in the
values in the data set and in Table D.6. Finally, the lower and upper limits for these central
estimates were estimated based on the assumption that there was 2.5 orders of magnitude
variability associated with the central estimates. The variability was based on visual inspection
of a number of figures containing the cesium Kd values, including Figure D. 1.
The calculations and equations used to estimate the central, minimum, and maximum estimates
used in the look-up tables are presented in Table D. 12.
D.17
-------
Table D.10. Estimated range of Kd values (ml/g) for cesium based on CEC or clay content for
systems containing <5% mica-like minerals in clay-size fraction and <10"9 M
aqueous cesium. [Table pertains to systems consisting of natural soils (as
opposed to pure mineral phases), low ionic strength (< 0.1 M), low humic
material concentrations (<5 mg/1), no organic chelates (such as EDTA), and
oxidizing conditions]
Kd (ml/g)
CEC (meq/100 g) / Clay Content (wt.%)
<3/<4
3-10/4-20
10-50/20-60
Central
200
500
1,500
Minimum
10
30
80
Maximum
3,500
9,000
26,700
Table D.ll. Estimated range of Kd values (ml/g) for cesium based on CEC or clay content for
systems containing >5% mica-like minerals in Clay-size fraction and <10"9 M
aqueous cesium. [Table pertains to systems consisting of natural soils (as
opposed to pure mineral phases), low ionic strength (< 0.1 M), low humic
material concentrations (<5 mg/1), no organic chelates (such as EDTA), and
oxidizing conditions.]
Kd (ml/g)
CEC (meq/100 g) / Clay Content (wt.%)
<3/<4
3-10/4-20
10-50/20-60
Central
500
1250
3750
Minimum
30
70
210
Maximum
9,000
22,000
66,700
D.18
-------
Table D.12. Calculations for values used in look-up table.
Mica
Concentration
in Clay Fraction
(%)
Clay
Content
(wt.%)
CE1
(ml/g)
Logarithm Scale
Base-10 Scale
Log CE
Lower Limit
(Log CE)/2
Lower Limit
10 (logCE)/2 (ml/g)
Upper Limit
J0 log CE + (log CE)/2 (ml/g)
<5
<4
200
2.301
1.151
14
2,828
<5
4-20
500
2.699
1.349
22
11,180
<5
20-60
1,500
3.176
1.588
39
58,095
>5
<4
500
2.699
1.349
22
11,180
>5
4-20
1,250
3.097
1.548
35
44,194
>5
20-60
3,750
3.574
1.787
61
229,640
1 CE = Central Estimate
D.19
-------
D.3.0 Kd Data Set for Soils and Pure Mineral Phases
Table D.13 lists the available cesium Kd values identified for experiments conducted with soils
and pure mineral phases.
Table D.13. Cesium Kd data base for soils and pure mineral phases
Cesium
Kd
(ml/g)
Clay
(wt.%)
Mica
(%)
PH
CECa
(meq/100 g)
SA1
(m2/g)
Aqueous Cs
G*M)
Background
Aqueous
Soil and Mineral
Phase ID and
Information
Ref2
247
6.2
1.90xl0"2
Gorleben
Groundwater
Gorleben Sediment
1
62
6.2
1.42x10"'
Gorleben Sediment
1
22
6.2
5.94X10"1
Gorleben Sediment
1
16
6.2
1.05
Gorleben Sediment
1
12
6.2
1.53
Gorleben Sediment
1
167
8.1
189
5.20xl0"3
Groundwater-1
SI: Quartz,
Kaolinite,
Plagioclase
2
1
7.8
113
5.20xl0"3
Groundwater-2
S2: Quartz,
Kaolinite, Dolomite
2
1500
9.3
60
70
1.00x10"'
Water pH 9.3
Bentonite
3
160
2.4
60
70
1.00x10"'
Groundwater
pH2.4
Bentonite
3
1100
9.3
60
70
1.00x10"'
Groundwater
pH 9.3
Bentonite
3
4100
6.1
20
130
1.00x10"'
Water pH 6.1
Takadate loam
3
1400
7.7
20
130
1.00x10"'
Groundwater
pH 7.7
Takadate loam
3
1100
6.6
70
60
1.00x10"'
Water pH 6.6
Hachinohe loam
3
280
8.3
70
60
1.00x10"'
Groundwater
pH 8.3
Hachinohe loam
3
237
8.2
2
22
l.OOxlO"3
ym-22
4
8220
8.2
109
103
l.OOxlO"3
ym-38
4
325
8.2
6
43
l.OOxlO"3
ym-45
4
22100
8.2
51
19
l.OOxlO"3
ym-48
4
35800
8.2
107
l.OOxlO"3
ym-49
4
42600
8.2
107
l.OOxlO"3
ym-49
4
205
8.2
4
l.OOxlO"3
ym-54
4
D.20
-------
Cesium
Kd
(ml/g)
Clay
(wt.%)
Mica
(%)
PH
CECa
(meq/100 g)
SA1
K/g)
Aqueous Cs
(HM)
Background
Aqueous
Soil and Mineral
Phase ID and
Information
Ref2
15200
8.4
31
l.OOxlO"3
low salts
JA-18
4
8440
8.3
31
l.OOxlO"3
hi salts
JA-18
4
143
8.2
8
l.OOxlO"3
low salts
JA-32
4
73
8.5
8
l.OOxlO"3
hi salts
JA-32
4
1390
8.4
100
l.OOxlO"3
low salts
JA-37
4
757
8.5
100
l.OOxlO"3
hi salts
JA-37
4
95
15
4
4.20xl04
0.005 MNa
Savannah River
5
120
15
5.5
4.20xl04
0.005 MNa
Savannah River
5
130
15
6.7
4.20xl04
0.005 MNa
Savannah River
5
130
15
7
4.20xl04
0.005 MNa
Savannah River
5
150
15
8.5
4.20xl04
0.005 MNa
Savannah River
5
160
15
10.2
4.20xl04
0.005 MNa
Savannah River
5
72
3
4
4.20xl04
0.005 MNa
4-Mile Creek
5
79
3
5.5
4.20xl04
0.005 MNa
4-Mile Creek
5
75
3
6.7
4.20xl04
0.005 MNa
4-Mile Creek
5
98
3
7
4.20xl04
0.005 MNa
4-Mile Creek
5
83
3
8.5
4.20xl04
0.005 MNa
4-Mile Creek
5
33
4
4
4.20xl04
0.005 MNa
Par Pond Soil
5
37
4
5.5
4.20xl04
0.005 MNa
Par Pond Soil
5
40
4
7
4.20xl04
0.005 MNa
Par Pond Soil
5
39
4
8.5
4.20xl04
0.005 MNa
Par Pond Soil
5
50
4
10.2
4.20xl04
0.005 MNa
Par Pond Soil
5
27
2
4
4.20xl04
0.005 MNa
Steel Creek Soil
5
25
2
5.5
4.20xl04
0.005 MNa
Steel Creek Soil
5
26
2
6.7
4.20xl04
0.005 MNa
Steel Creek Soil
5
26
2
7
4.20xl04
0.005 MNa
Steel Creek Soil
5
38
2
8.5
4.20xl04
0.005 MNa
Steel Creek Soil
5
39
2
10.2
4.20xl04
0.005 MNa
Steel Creek Soil
5
88
4
4
4.20xl04
0.005 MNa
Lower 3 Runs Soil
5
92
4
5.5
4.20xl04
0.005 MNa
Lower 3 Runs Soil
5
93
4
6.7
4.20xl04
0.005 MNa
Lower 3 Runs Soil
5
85
4
7
4.20xl04
0.005 MNa
Lower 3 Runs Soil
5
D.21
-------
Cesium
Kd
(ml/g)
Clay
(wt.%)
Mica
(%)
PH
CECa
(meq/100 g)
SA1
K/g)
Aqueous Cs
(HM)
Background
Aqueous
Soil and Mineral
Phase ID and
Information
Ref2
94
4
8.5
4.20xl04
0.005 MNa
Lower 3 Runs Soil
5
101
4
10.2
4.20xl04
0.005 MNa
Lower 3 Runs Soil
5
88
5
4
4.20xl04
0.005 MNa
Pen Branch Soil
5
89
5
5.5
4.20xl04
0.005 MNa
Pen Branch Soil
5
90
5
6.7
4.20xl04
0.005 MNa
Pen Branch Soil
5
84
5
7
4.20xl04
0.005 MNa
Pen Branch Soil
5
101
5
10.2
4.20xl04
0.005 MNa
Pen Branch Soil
5
22
2
4
4.20xl04
0.005 MNa
Upper 3 Runs Soil
5
31
2
5.5
4.20xl04
0.005 MNa
Upper 3 Runs Soil
5
37
2
6.7
4.20xl04
0.005 MNa
Upper 3 Runs Soil
5
40
2
7
4.20xl04
0.005 MNa
Upper 3 Runs Soil
5
78
2
10.2
4.20xl04
0.005 MNa
Upper 3 Runs Soil
5
27
8.25
1.83
17.7
2.72xl02
0.002 M
Groundwater
Umtanum Basalt
6
329
8.25
1.83
17.7
2.90X10"1
0.002 M
Groundwater
Umtanum Basalt
6
960
8.25
1.83
17.7
1.03xl0"3
0.002 M
Groundwater
Umtanum Basalt
6
1088
8.25
1.83
17.7
9.11xl0"6
0.002 M
Groundwater
Umtanum Basalt
6
1084
8.25
1.83
17.7
1.87xl0"6
0.002 M
Groundwater
Umtanum Basalt
6
28
8.6
1.83
17.7
2.63xl02
0.013 M
Groundwater
Umtanum Basalt
6
289
8.6
1.83
17.7
3.31x10"'
0.013 M
Groundwater
Umtanum Basalt
6
951
8.6
1.83
17.7
1.05xl0"3
0.013 M
Groundwater
Umtanum Basalt
6
1022
8.6
1.83
17.7
9.77xl0"6
0.013 M
Groundwater
Umtanum Basalt
6
1025
8.6
1.83
17.7
1.95xl0"6
0.013 M
Groundwater
Umtanum Basalt
6
18
8.2
1.5
10.3
3.61xl02
0.002 M
Groundwater
Flow E Basalt
6
189
8.2
1.5
10.3
5.00x10"'
0.002 M
Groundwater
Flow E Basalt
6
418
8.2
1.5
10.3
2.34x10"3
0.002 M
Groundwater
Flow E Basalt
6
D.22
-------
Cesium
Kd
(ml/g)
Clay
(wt.%)
Mica
(%)
PH
CECa
(meq/100 g)
SA1
K/g)
Aqueous Cs
(HM)
Background
Aqueous
Soil and Mineral
Phase ID and
Information
Ref2
450
8.2
1.5
10.3
2.17xl0"5
0.002 M
Groundwater
Flow E Basalt
6
487
8.2
1.5
10.3
3.98xl0"6
0.002 M
Groundwater
Flow E Basalt
6
20
8.7
1.5
10.3
3.39xl02
0.013 M
Groundwater
Flow E Basalt
6
214
8.7
1.5
10.3
4.47X10"1
0.013 M
Groundwater
Flow E Basalt
6
488
8.7
1.5
10.3
2.00xl0"3
0.013 M
Groundwater
Flow E Basalt
6
549
8.7
1.5
10.3
1.78xl0"5
0.013 M
Groundwater
Flow E Basalt
6
617
8.7
1.5
10.3
3.24xl0"6
0.013 M
Groundwater
Flow E Basalt
6
48
8.3
4.84
31.2
1.71xl02
0.002 M
Groundwater
Pomona Basalt
6
460
8.3
4.84
31.2
2.13x10"'
0.002 M
Groundwater
Pomona Basalt
6
1111
8.3
4.84
31.2
8.30xl04
0.002 M
Groundwater
Pomona Basalt
6
1466
8.3
4.84
31.2
6.37x10"6
0.002 M
Groundwater
Pomona Basalt
6
1281
8.3
4.84
31.2
1.39xl0"6
0.002 M
Groundwater
Pomona Basalt
6
56
8.55
4.84
31.2
1.51xl02
0.013 M
Groundwater
Pomona Basalt
6
389
8.55
4.84
31.2
2.57x10"'
0.013 M
Groundwater
Pomona Basalt
6
853
8.55
4.84
31.2
1.17xl0"3
0.013 M
Groundwater
Pomona Basalt
6
952
8.55
4.84
31.2
1.05xl0"5
0.013 M
Groundwater
Pomona Basalt
6
908
8.55
4.84
31.2
1.74xl0"6
0.013 M
Groundwater
Pomona Basalt
6
212
8.3
71
646
4.50x10'
0.002 M
Groundwater
Smectite
6
1080
8.3
71
646
9.17x10"'
0.002 M
Groundwater
Smectite
6
13042
8.3
71
646
7.66x10"5
0.002 M
Groundwater
Smectite
6
D.23
-------
Cesium
Kd
(ml/g)
Clay
(wt.%)
Mica
(%)
PH
CECa
(meq/100 g)
SA1
K/g)
Aqueous Cs
(HM)
Background
Aqueous
Soil and Mineral
Phase ID and
Information
Ref2
9794
8.3
71
646
l.OOxlO"6
0.002 M
Groundwater
Smectite
6
25000
8.3
71
646
7.00x10"8
0.002 M
Groundwater
Smectite
6
224
9.2
71
646
4.27X10"1
0.013 M
Groundwater
Smectite
6
2136
9.2
71
646
4.68xl0"2
0.013 M
Groundwater
Smectite
6
5882
9.2
71
646
1.70xl04
0.013 M
Groundwater
Smectite
6
8547
9.2
71
646
1.17xl0"6
0.013 M
Groundwater
Smectite
6
8333
9.2
71
646
2.40x10"7
0.013 M
Groundwater
Smectite
6
5000
24
4.4
82
6.80x10"2
lxlO"6 MKC1
Batcombe
7
5000
24
4.4
82
6.80x10"2
lxlO"5 MKC1
Batcombe
7
4700
24
4.4
82
6.80x10"2
lxl04 MKC1
Batcombe
7
2000
24
4.4
82
6.80x10"2
lxlO"3 MKC1
Batcombe
7
9000
42
6.2
72
6.80x10"2
lxlO"6 MKC1
Tedburn
7
8000
42
6.2
72
6.80x10"2
lxlO"5 MKC1
Tedburn
7
9000
42
6.2
72
6.80x10"2
lxlO4 MKC1
Tedburn
7
2000
42
6.2
72
6.80x10"2
lxlO"3 MKC1
Tedburn
7
1050
42
7.3
54
6.80x10"2
lxlO"6 MKC1
Teigngrace
7
1025
42
7.3
54
6.80x10"2
lxlO"5 MKC1
Teigngrace
7
1000
42
7.3
54
6.80x10"2
lxlO4 MKC1
Teigngrace
7
800
42
7.3
54
6.80x10"2
lxlO"3 MKC1
Teigngrace
7
11000
130
l.OOxlO"7
Water
Itago Tuff
8
10000
97
l.OOxlO"7
Water
Ohya Tuff
8
5000
2.4
l.OOxlO"7
Water
Sandstone
8
2000
1.9
l.OOxlO"7
Water
Shale
8
6000
1.9
l.OOxlO"7
Water
Augite Audesite
8
500
1.2
l.OOxlO"7
Water
Plagio Rhyolite
8
5800
0.75
l.OOxlO"7
Water
Olivine Basalt
8
900
0.54
l.OOxlO"7
Water
Ionada Granite
8
260
0.35
l.OOxlO"7
Water
Rokka Granite
8
D.24
-------
Cesium
Kd
(ml/g)
Clay
(wt.%)
Mica
(%)
PH
CECa
(meq/100 g)
SA1
K/g)
Aqueous Cs
(HM)
Background
Aqueous
Soil and Mineral
Phase ID and
Information
Ref2
80
0.033
l.OOxlO"7
Water
Limestone
8
2200
1.2
l.OOxlO"7
Water
Biotite
8
1800
0.93
l.OOxlO"7
Water
Chlorite
8
630
0.33
l.OOxlO"7
Water
Hornblende
8
420
0.11
l.OOxlO"7
Water
Grossular
8
460
0.0067
l.OOxlO"7
Water
Forsterite
8
30
0.0034
l.OOxlO"7
Water
K-feldspar
8
89
0.0032
l.OOxlO"7
Water
Albite
8
31
0.00098
l.OOxlO"7
Water
Quartz
8
1
0.15849
l.OOxlO"1
Calcite
9
3
0.19953
l.OOxlO"1
Apatite
9
6
1.58489
l.OOxlO"1
Hematite
9
13
1.77828
l.OOxlO"1
Orthoclase
9
16
5.62341
l.OOxlO"1
Serpentine
9
200
7.94328
l.OOxlO"1
Hornblende
9
631
39.8107
l.OOxlO"1
Biotite
9
794
63.0957
l.OOxlO"1
Muscovite
9
100
4.46684
l.OOxlO"1
Gneiss
9
16
6.30957
l.OOxlO"1
Diabase
9
158
10
l.OOxlO"1
Stripa Granite
9
562
11.2202
l.OOxlO"1
Finsjo Granite
9
900
5
l.OOxlO"1
Biotite
9
790
7
l.OOxlO"1
Biotite
9
700
9
l.OOxlO"1
Biotite
9
2
5
l.OOxlO"1
Hematite
9
4
7
l.OOxlO"1
Hematite
9
8
9
l.OOxlO"1
Hematite
9
40
5
l.OOxlO"1
Hornblende
9
100
7
l.OOxlO"1
Hornblende
9
240
9
l.OOxlO"1
Hornblende
9
3
5
l.OOxlO"1
Magnetite
9
5
7
l.OOxlO"1
Magnetite
9
D.25
-------
Cesium
Kd
(ml/g)
Clay
(wt.%)
Mica
(%)
PH
CECa
(meq/100 g)
SA1
K/g)
Aqueous Cs
(HM)
Background
Aqueous
Soil and Mineral
Phase ID and
Information
Ref2
9
9
l.OOxlO"1
Magnetite
9
700
5
l.OOxlO"1
Muscovite
9
810
7
l.OOxlO"1
Muscovite
9
840
9
l.OOxlO"1
Muscovite
9
7
5
l.OOxlO"1
Orthoclase
9
14
7
l.OOxlO"1
Orthoclase
9
7
9
l.OOxlO"1
Orthoclase
9
52000
127
1.67xl0"6
Deionized Water
Vermiculite
10
26000
20
1.67xl0"6
Deionized Water
Illite
10
2500
11.2
1.67xl0"6
Deionized Water
Kaolinite
10
2700
127
1.67xl0"6
0.1 NNaCl
Vermiculite
10
28600
20
1.67xl0"6
0.1 NNaCl
Illite
10
94
11.2
1.67xl0"6
0.1 NNaCl
Kaolinite
10
7
l.OOxlO"7
Groundwater
Hanford Vadose
Sediment
11
12
l.OOxlO"7
Groundwater
Hanford Vadose
Sediment
11
2190
4
9
7.7
8.40x10"3
Groundwater
Sediment CGS-1
12
7610
5
12
8.2
8.40xl0"3
Groundwater
Sediment TBS-1
12
620
6
9
7.9
8.40x10"3
Groundwater
Sediment Trench-8
12
1 CEC = cation exchange capacity; SA = surface area.
2 References: 1 = Lieser and Steinkopff, 1989; 2 = Lieser etal., 1986; 3 =Konishi etal., 1988; 4 = Vine etal., 1980;
5 = Elprince etal., 1977; 6 = Ames etal., 1982; 7 = Staunton, 1994; 8 = Akiba etal., 1989; 9 = Torstenfelt etal., 1982;
10 = Tamura, 1972; 11 = Routson etal., 1980; 12 = Sern eetal., 1993.
D.26
-------
D.4.0 Data Set for Soils
Table D.14 lists the available cesium Kd values identified for experiments conducted with only
soils.
Table D.14. Cesium Kd data set for soils only.
Cesium
K,
(ml/g)
Clay
(wt%)
Mica
(%)
PH
CEC(a)
(meq/100 g)
SA1
(m2/g)
Cs
(HM)
Aqueous Phase
Soil ID
and Information
Ref.2
247
6.2
1.90xl0"2
Gorleben
Groundwater
Gorleben Sediment
1
62
6.2
1.42x10"'
Gorleben Sediment
1
22
6.2
5.94X10"1
Gorleben Sediment
1
4100
6.1
20
130
1.00x10"'
Water pH 6.1
Takadate Loam
4
1400
7.7
20
130
1.00x10"'
Groundwater
pH 7.7
Takadate Loam
4
1100
6.6
70
60
1.00x10"'
Water pH 6.6
Hachinohe Loam
4
280
8.3
70
60
1.00x10"'
Groundwater
pH 8.3
Hachinohe loam
4
95
15
4
4.20xl0"4
0.005 MNa
Sav. River Site
Sediment
6
120
15
5.5
4.20xl0"4
0.005 MNa
Sav. River Site
Sediment
6
130
15
6.7
4.20xl0"4
0.005 MNa
Sav. River Site
Sediment
6
130
15
7
4.20xl0"4
0.005 MNa
Sav. River Site
Sediment
6
150
15
8.5
4.20xl0"4
0.005 MNa
Sav. River Site
Sediment
6
160
15
10.2
4.20xl0"4
0.005 MNa
Sav. River Site
Sediment
6
72
3
4
4.20xl0"4
0.005 MNa
4-Mile Creek Sediment
6
79
3
5.5
4.20xl0"4
0.005 MNa
4-Mile Creek Sediment
6
75
3
6.7
4.20xl0"4
0.005 MNa
4-Mile Creek
Sediment.
6
98
3
7
4.20xl0"4
0.005 MNa
4-Mile Creek
Sediment.
6
83
3
8.5
4.20xl0"4
0.005 MNa
4-Mile Creek
Sediment.
6
33
4
4
4.20xl0"4
0.005 MNa
Par Pond Soil
6
D.27
-------
Cesium
K,
(ml/g)
Clay
(wt%)
Mica
(%)
PH
CEC(a)
(meq/100 g)
SA1
(m2/g)
Cs
(HM)
Aqueous Phase
Soil ID
and Information
Ref.2
37
4
5.5
4.20xl0"4
0.005 MNa
Par Pond Soil
6
40
4
7
4.20xl0"4
0.005 MNa
Par Pond Soil
6
39
4
8.5
4.20xl0"4
0.005 MNa
Par Pond Soil
6
50
4
10.2
4.20xl0"4
0.005 MNa
Par Pond Soil
6
27
2
4
4.20xl0"4
0.005 MNa
Steel Creek Soil
6
25
2
5.5
4.20xl0"4
0.005 MNa
Steel Creek Soil
6
26
2
6.7
4.20xl0"4
0.005 MNa
Steel Creek Soil
6
26
2
7
4.20xl0"4
0.005 MNa
Steel Creek Soil
6
38
2
8.5
4.20xl0"4
0.005 MNa
Steel Creek Soil
6
39
2
10.2
4.20xl0"4
0.005 MNa
Steel Creek Soil
6
88
4
4
4.20xl0"4
0.005 MNa
Lower 3 Runs Soil
6
92
4
5.5
4.20xl0"4
0.005 MNa
Lower 3 Runs
Sediment
6
93
4
6.7
4.20xl0"4
0.005 MNa
Lower 3 Runs
Sediment
6
85
4
7
4.20xl0"4
0.005 MNa
Lower 3 Runs
Sediment
6
94
4
8.5
4.20xl0"4
0.005 MNa
Lower 3 Runs
Sediment
6
101
4
10.2
4.20xl0"4
0.005 MNa
Lower 3 Runs
Sediment
6
88
5
4
4.20xl0"4
0.005 MNa
Pen Branch Soil
6
89
5
5.5
4.20xl0"4
0.005 MNa
Pen Branch Soil
6
90
5
6.7
4.20xl0"4
0.005 MNa
Pen Branch Soil
6
84
5
7
4.20xl0"4
0.005 MNa
Pen Branch Soil
6
101
5
10.2
4.20xl0"4
0.005 MNa
Pen Branch Soil
6
22
2
4
4.20xl0"4
0.005 MNa
Upper 3 Runs Soil
6
31
2
5.5
4.20xl0"4
0.005 MNa
Upper 3 Runs Soil
6
37
2
6.7
4.20xl0"4
0.005 MNa
Upper 3 Runs Soil
6
40
2
7
4.20xl0"4
0.005 MNa
Upper 3 Runs Soil
6
78
2
10.2
4.20xl0"4
0.005 MNa
Upper 3 Runs Soil
6
7
l.OOxlO"7
Groundwater
Hanford Vadose
Sediment
8
12
l.OOxlO"7
Groundwater
Hanford Vadose
Sediment
8
D.28
-------
Cesium
K,
(ml/g)
Clay
(wt%)
Mica
(%)
PH
CEC(a)
(meq/100 g)
SA1
(m2/g)
Cs
(HM)
Aqueous Phase
Soil ID
and Information
Ref.2
3,000
6
7.6
3
8.6
l.OOxlO"1
Groundwater
Sediment A
10
4,800
7.5
5.9
4.3
12.2
l.OOxlO"1
Groundwater
Sediment B
10
3,100
8
6.6
4.7
14.7
l.OOxlO"1
Groundwater
Sediment C
10
3,000
5
8
2.6
6.6
l.OOxlO"1
Groundwater
Sediment D
10
2,190
4
9
7.7
8.40xl0"3
Groundwater
Sediment CGS-1
11
7,610
5
12
8.2
8.40xl0"3
Groundwater
Sediment TBS-1
11
620
6
9
7.9
8.40xl0"3
Groundwater
Sediment Trench-8
11
1 CEC = cation exchange capacity; SA = surface area.
2 1 = Lieser and Steinkopff, 1989; 4 = Konishi etal., 1988; 6 = Elprince etal., 1977; 8 = Routson etal., 1980; 10 = Legoux et
al, 1992; 11 = Sern eetal., 1993.
D.5.0 References
Adeleye, S. A., P. G. Clay, and M. O. A. Oladipo. 1994. "Sorption of Caesium, Strontium and
Europium Ions on Clay Minerals." Journal of Materials Science, 29:954-958.
Akiba, D., and H. Hashimoto. 1990. "Distribution Coefficient of Strontium on Variety of
Minerals and Rocks." Journal of Nuclear Science and Technology, 21:215-219.
Akiba, D., H. Hashimoto, and T. Kanno. 1989. "Distribution Coefficient of Cesium and Cation
Exchange Capacity of Minerals and Rocks." Journal of Nuclear Science and Technology,
26:1130-1135.
Ames, L., and D. Rai. 1978. Radionuclide Interactions with Soil and Rock Media. Volume 1:
Processes Influencing Radionuclide Mobility and Retention, Element Chemistry and
Geochemistry, Conclusions and Evaluation. PB-292 460, Pacific Northwest Laboratory,
Richland, Washington.
Ames, L. L., J. E. McGarrah, B. A. Walker, and P. F. Salter. 1982. " Sorption of Uranium and
Cesium by Hanford Basalts and Associated Secondary Smectite." Chemical Geology,
35:205-225.
Comans, R. N. J., J. J. Middelburg, J. Zonderhuis, J. R. W. Woittiez, G. J. De Lange, H. A. Das,
C. H. Van Der Weijden. 1989. "Mobilization of Radiocaesium in Pore Water in Lake
Sediments." Nature, 367-369.
Cremers, A., A. Elsen. P. De Preter, and A. Maes. 1988. "Quantitative Analysis of
Radiocaesium Retention in Soils." Nature, 335:247-249.
D.29
-------
Bruggenwert, M. G. M., and A. Kamphorst. 1979. "Survey of Experimental Information on
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Behavior of Plutonium and Other Actinides in the Environment." In Actinides in the
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D.31
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D.32
-------
APPENDIX E
Partition Coefficients For Chromium(VI)
-------
Appendix E
Partition Coefficients For Chromium(VI)
E.1.0 Background
The review of chromium Kd data obtained for a number of soils (summarized in Table E.l)
indicated that a number of factors influence the adsorption behavior of chromium. These factors
and their effects on chromium adsorption on soils and sediments were used as the basis for
generating a look-up table. These factors are:
• Concentrations of Cr(III) in soil solutions are typically controlled by
dissolution/precipitation reactions therefore, adsorption reactions are not significant in
soil Cr(III) chemistry.
• Increasing pH decreases adsorption (decrease in Kd) of Cr(VI) on minerals and soils. The
data are quantified for only a limited number of soils.
• The redox state of the soil affects chromium adsorption. Ferrous iron associated with
iron oxide/hydroxide minerals in soils can reduce Cr(VI) which results in precipitation
(higher Kd). Soils containing Mn oxides oxidize Cr(III) into Cr(VI) form thus resulting
in lower Kd values. The relation between oxide/hydroxide contents of iron and
manganese and their effects on Kd have not been adequately quantified except for a few
soils.
• The presence of competing anions reduce Cr(VI) adsorption. The inhibiting effect varies
in the order HPO4", H2PC>4 »S04_ CO3/HCO3 CI", N03". These effects have been
quantified as a function of pH for only 2 soils.
The factors which influence chromium adsorption were identified from the following sources of
data. Experimental data for Cr(VI) adsorption onto iron oxyhydroxide and aluminum hydroxide
minerals (Davis andLeckie, 1980; Griffin etal., 1977; Leckie etal., 1980; Rai etal., 1986)
indicate that adsorption increases with decreasing pH over the pH range 4 to 10. Such
adsorption behavior is explained on the basis that these oxides show a decrease in the number of
positively charged surface sites with increasing pH. Rai etal. (1986) investigated the adsorption
behavior of Cr(VI) on amorphous iron oxide surfaces. The experiments were conducted with
initial concentrations of 5xl0"6 M Cr(VI). The results showed very high Kd values (478,630
ml/g) at lower pH values (5.65), and lower Kd values (6,607 ml/g) at higher pH values (7.80). In
the presence of competing anions (S04: 2.5xl0"3 M, solution in equilibrium with 3.5xl0"3 atm
C02), at the same pH values, the observed Kd values were 18,620 ml/g and 132 ml/g respectively
leading to the conclusion that depending on concentration competing anions reduce Cr(VI)
adsorption by at least an order of magnitude. Column experiments on 3 different soils conducted
by Selim and Amacher (1988) confirmed the influence of soil pH on Cr(VI) adsorption. Cecil,
E.2
-------
Windsor, and Olivier soils with pH values of 5.1, 5.4, and 6.4 exhibited chromium Kd values in
the range -9-100 ml/g, 2-10 ml/g, and -1-3 ml/g respectively. Adsorption of Cr(VI) on
4 different subsoils was studied by Rai et al. (1988). The authors interpreted the results of these
experiments using surface complexation models. Using their adsorption data, we calculated the
Kd values for these soils. The data showed that 3 of the 4 soils studied exhibited decreasing Kd
values with increasing pH. The Kd values for these soils were close to 1 ml/g at higher pH
values (>8). At lower pH values (about 4.5) the Kd values were about 2 to 3 orders of magnitude
greater than the values observed at higher pH values One of the soils with a very high natural
pH value (10.5) however did not show any adsorption affinity (Kd < 1 ml/g) for Cr(VI).
The data regarding the effects of soil organic matter on Cr(VI) adsorption are rather sparse. In
1 study, Stollenwerk and Grove (1985) evaluated the effects of soil organic matter on adsorption
of Cr(VI). Their results indicated that organic matter did not influence Cr(VI) adsorption
properties. In another study, the Cr(VI) adsorption properties of an organic soil was examined
by Wong et al. (1983). The chromium adsorption measurements on bottom, middle, and top
layers of this soil produced Kd values of 346, 865, and 2,905 ml/g respectively. Also, another Kd
measurement using an organic-rich fine sandy soil from the same area yielded a value of 1,729
ml/g.
A series of column (lysimeter) measurements involving Cr(VI) adsorption on 4 different layers
of a sandy soil yielded average Kd values that ranged from 6 to 263 ml/g (Sheppard et al., 1987).
These measurements showed that coarse-textured soils tend to have lower Kd values as compared
to fine-textured soils such as loam (Kd - 1,000 ml/g, Sheppard and Sheppard, 1987).
Stollenwerk and Grove (1985) examined Cr(VI) adsorption on an alluvium from an aquifer in
Telluride, Colorado. A Kd value of 5 ml/g was obtained for Cr(VI) adsorption on this alluvium.
Removing organic matter from the soil did not significantly affect the Kd value. However,
removing iron oxide and hydroxide coatings resulted in a Kd value of about 0.25 leading the
authors to conclude that a major fraction of Cr(VI) adsorption capacity of this soil is due to its
iron oxide and hydroxide content. Desorption experiments conducted on Cr adsorbed soil aged
for 1.5 yrs indicated that over this time period, a fraction of Cr(VI) had been reduced to Cr(III)
by ferrous iron and had probably coprecipitated with iron hydroxides.
Studies by Stollenwerk and Grove (1985) and Sheppard et al. (1987) using soils showed that Kd
decreases as a function of increasing equilibrium concentration of Cr(VI). Another study
conducted by Rai et al. (1988) on 4 different soils confirmed that Kd values decrease with
increasing equilibrium Cr(VI) concentration.
Other studies also show that iron and manganese oxide contents of soils significantly affect the
adsorption of Cr(VI) on soils (Korte et al., 1976). However, these investigators did not publish
either Kd values or any correlative relationships between Kd and the oxide contents. The
adsorption data obtained by Rai et al. (1988) also showed that quantities of sodium dithionite-
citrate-bicarbonate (DCB) extractable iron content of soils is a good indicator of a soil's ability
to reduce Cr(VI) to Cr(III) oxidation state. The reduced Cr has been shown to coprecipitate with
ferric hydroxide. Therefore, observed removal of Cr(VI) from solution when contacted with
E.3
-------
chromium-reductive soils may stem from both adsorption and precipitation reaction. Similarly,
Rai et al. (1988) also showed that certain soils containing manganese oxides may oxidize Cr(III)
into Cr(VI). Depending on solution concentrations, the oxidized form (VI) of chromium may
also precipitate in the form of Ba(S,Cr)04 Such complex geochemical behavior chromium in
soils implies that depending on the properties of a soil, the measured Kd values may reflect both
adsorption and precipitation reactions.
An evaluation of competing anions indicated that Cr(VI) adsorption was inhibited to the greatest
extent by HPO4" and H2PC>4 ions and to a very small extent by CI" and N03" ions. The data
indicate that Cr(VI) adsorption was inhibited by anions in order of HPO4", H2PC>4 » SO4" » CI",
NOj (Leckie et al., 1980; MacNaughton, 1977; Rai et al., 1986; Rai et al., 1988; Stollenwerk and
Grove, 1985).
E.4
-------
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E.2.0 Approach
The approach used to develop the look-up table was to identify the key parameters that control Cr(VI)
adsorption reactions. From the data of Rai etal. (1988) and other studies of Cr(VI) adsorption on soils pH
was identified as a key parameter. The data show (Table E.2) that the Kd values are significantly higher at
lower pH values and decline with increasing pH. Also, Kd values for soils show a wider range at lower pH,
but values for all soils converge as pH value approaches about 8. Another parameter which seems to
influence soil adsorption of Cr(VI) is the capacity of soils to reduce Cr(VI) to Cr(III). Leckie el al. (1980)
and Rai et al. (1988) showed that iron oxides in the soil reduce Cr(VI) to Cr(III) and precipitate Cr(III) as a
(Fe,Cr)(OH)3 mineral. Also, studies conducted by Rai el al. (1988) show that DCB extractable iron content
is a good indicator as to whether a soil can reduce significant quantities of Cr(VI) which results in higher Kd
values. It is important to note the total iron oxide content is a poor indicator of a soil's Cr(VI) reducing
capacity and that DCB extractable iron better represents the fraction of iron content that would reduce
Cr(VI) to Cr(III). The data indicated that Holton/Cloudland soil with the highest concentrations of DCB
extractable iron (0.435 mmol/g) exhibited higher Kd values than other soils which did not show an
observable Cr(VI) reduction tendency.
Based on this information, 4 ranges of pH, which encompass the pH range of most natural soils, were
selected for the look-up table (Table E.3). Within each pH range, 3 ranges of DCB extractable iron content
were selected to represent the categories of soils that definitely reduce (>0.3 mmol/g), probably reduce (0.26
to 0.29 mmol/g), and do not reduce (<2.5 mmol/g) Cr(VI) to Cr(III) form. The range of Kd values to be
expected within each of the 12 categories was estimated from the data listed in Table E.2. The variations of
Kd values as a function of pH and DCB extractable iron as independent variables based on experimental data
(Table E.2) is also shown as a 3-dimensional graph (Figure E.l). The graph indicates that soils with lower
pH values and higher DCB extractable iron contents exhibit greater adsorption (higher Kd) of Cr(VI). At
higher pH values (>7), Cr(VI) adsorption tends to be very low (very low Kd values) irrespective of DCB
extractable iron content. Similarly, soils which contain very low DCB extractable iron, adsorb very little
Cr(VI) (very low Kd values) irrespective of soil pH values.
Additionally, Cr(VI) adsorption studies show that the presence of competing anions such as HPO4", H2PC>4,
SO4", CO3", and HCOj will reduce the Kd values as compared to a noncompetitive adsorption process. The
only available data set that can be used to assess the competing anion effect was developed by Rai et al.
(1988). However, they used fixed concentrations of competing anions namely SO4", CO3", and HCOj (fixed
through a single selected partial pressure of C02) concentrations (Tables E.4 and E.5). Among these
competing anions, SO4" at about 3 orders of magnitude higher concentrations (2 x 10"3 M or 191.5 mg/1) than
Cr(VI) concentration depressed Cr(VI) Kd values roughly by an order of magnitude as compared to
noncompetitive adsorption. Therefore, the look-up table was developed on the assumption that Kd values of
Cr(VI) would be reduced as soluble SO4" concentrations increase from 0 to 2xl0"3 M (or 191.5 mg/1).
E.7
-------
| Ocala Soil |
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6.67 |
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6.36 |
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6.07 |
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3.02 |
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3.03 |
3.03
3.02 |
3.02
3.03 |
3.03
3.03 |
3.02
3.03 |
3.02
3.04 |
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8.94 |
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E.9
-------
Figure E.l. Variation of Kd for Cr(VI) as a function of pH and DCB extractable iron
content without the presence of competing anions.
E.3.0 Data Set for Soils
The data set used to develop the look-up table is from the adsorption data collected by Rai et al. (1988). The
adsorption data for Cr(VI) as a function of pH developed for 4 well-characterized soils were used to
calculate the Kd values (Table E.2). All 4 soil samples were obtained from subsurface horizons and
characterized as to their pH, texture, CEC, organic and inorganic carbon contents, surface areas, extractable
(hydroxylamine hydrochloride, and DCB) iron, manganese, aluminum, and silica, KOH extractable
aluminum and silica, and clay mineralogy. Additionally, Cr oxidizing and reducing properties of these soils
were also determined (Rai etal., 1988). Effects of competing anions such as sulfate and carbonate on Cr(VI)
adsorption were determined for 2 of the soils (Cecil/Pacolet, and Kehoma). The Kd values from competitive
anion experiments were calculated (Tables E.4 and E.5) and used in developing the look-up table
(Table E.3).
E.10
-------
Table E.4. Data from Rai el al. (1988) on effects of competing anions on Cr(VI)
adsorption on Cecil/Pacolet soil.
CifVI)1
Cr(VI) + Sulfate1
Cr(YI) + Carbonate1
PH
-log C
(mol/m3)
-logS
(mol/kg)
K,
(ml/g)
PH
-log C
(mol/m3)
-log S
(mol/kg)
K,
(ml/g)
PH
-log C
(mol/m3)
-logS
(mol/kg)
K,
(ml/g)
9.26
3.05
5.66
2
8.92
3.05
6.27
1
9.62
3.05
6.88
0
9.29
3.05
5.88
1
8.38
3.07
5.71
2
9.15
3.05
6.79
0
8.57
3.11
5.34
6
8.38
3.04
5.70
2
9.01
3.06
6.35
1
7.80
3.30
5.00
20
7.70
3.12
5.28
7
7.92
3.06
6.12
1
7.41
3.44
4.89
35
7.67
3.12
5.28
7
7.95
3.06
6.10
1
7.38
3.46
4.88
38
7.37
3.19
5.11
12
7.53
3.08
5.85
2
6.99
3.66
4.81
71
7.24
3.23
5.09
14
7.52
3.07
6.06
1
6.94
3.65
4.81
69
6.85
3.34
4.95
24
7.19
3.12
5.55
4
6.67
3.79
4.78
102
6.76
3.37
4.96
26
7.31
3.10
5.67
3
6.49
3.79
4.78
102
6.58
3.43
4.92
32
7.22
3.12
5.55
4
6.19
3.99
4.75
174
6.56
3.34
4.95
25
6.99
3.13
5.48
4
6.16
3.94
4.75
155
6.15
3.55
4.85
50
6.70
3.22
5.21
10
5.89
4.08
4.74
219
6.15
3.51
4.88
43
6.68
3.21
5.24
9
5.84
4.06
4.74
209
5.75
3.58
4.82
58
5.84
3.65
4.87
60
5.46
4.19
4.73
288
5.79
3.56
4.86
51
6.08
3.54
4.91
43
5.49
4.21
4.73
302
5.35
3.60
4.83
59
5.12
4.11
4.78
214
4.98
4.33
4.72
407
5.33
3.59
4.84
57
5.12
4.14
4.78
229
4.98
4.32
4.72
398
4.68
3.55
4.86
49
4.76
4.20
4.78
263
4.49
4.52
4.71
646
4.69
3.47
4.86
41
4.75
4.11
4.78
214
4.49
4.39
4.72
468
4.33
4.39
4.76
427
4.34
4.37
4.77
398
1 Cr(VI) concentration: 10"6M, Sulfate Concentration: 10"2 7M, C02: 10"16atm.
E. 11
-------
Table E.5. Data from Rai et al. (1988) on effects of competing anions on
Cr(VI) adsorption on Kenoma soil.
CifVI)1
Cr(VI) + Sulfate + Carbonate1
PH
-log C
(mol/m3)
-log S
(mol/kg)
K,
(ml/g)
PH
-log C
(mol/m3)
-logS
(mol/kg)
K,
(ml/g)
8.42
3.03
6.25
1
7.49
3.06
6.22
1
7.71
3.05
5.84
2
7.42
3.06
6.35
1
7.70
3.04
5.97
1
7.3
3.07
5.98
1
7.35
3.09
5.54
4
7.38
3.08
5.9
2
7.40
3.08
5.59
3
7.08
3.08
5.83
2
7.20
3.03
5.36
5
6.93
3.1
5.64
3
7.16
3.13
5.37
6
6.49
3.15
5.43
5
6.89
3.16
5.27
8
6.52
3.16
5.39
6
6.92
3.15
5.29
7
6.32
3.17
5.33
7
6.70
3.23
5.13
13
6.32
3.18
5.31
7
6.47
3.26
5.09
15
5.97
3.23
5.21
10
6.02
3.36
4.98
24
5.97
3.21
5.25
9
6.02
3.35
4.99
23
5.7
3.23
5.2
11
5.61
3.39
4.95
28
5.69
3.24
5.18
11
5.62
3.40
4.95
28
5.54
3.24
5.19
11
5.52
3.25
5.18
12
5.03
3.18
5.32
7
5.02
3.21
5.26
9
Cr( VI) concentration: 10"6M, Sulfate Concentration: 10"2 7M, C02: 10"16atm.
E. 12
-------
E.4.0 References
Davis, J. A. and J. O. Leckie. 1980. "Surface Ionization and Complexation at the Oxide/Water Interface. 3.
Adsorption of Anions." Journal of Colloid Interfacial Science, 74:32-43.
Griffin, R. A., A. K. Au, and R. R. Frost. 1977. "Effect of pH on adsorption of Chromium form Landfill-
Leachate by Clay Minerals." Journal of Environmental Science Health, 12:431-449.
Korte N. E., J. Skopp, W. H. Fuller, E. E. Niebla and B. A. Alesii. 1976. "Trace Element Movement in
Soils: Influence of Soil Physical and Chemical Properties." Soil Science, 122:350-359.
Leckie, J. O., M. M. Benjamin, K. Hayes, G. Kaufman, and S. Altman. 1980. Adsorption/Coprecipitation of
Trace Elements from Water with Iron Oxyhydroxides. EPRI-RP-910. Electric Power Research Institute,
Palo Alto, California.
MacNaughton, M. G. 1977. "Adsorption of Chromium (VI) at the Oxide-Water Interface." In Biological
Implications of Metals in the Environment, H. Drucker and R. F. Wildung (eds.), pp. 244-253, CONF-
750929, National Technical Information Service, Springfield, Virginia.
Rai, D., J. M. Zachara, L. E. Eary, C. C. Ainsworth, J. E. Amonette, C. E. Cowan, R. W. Szelmeczka, C. T.
Resch, R. L. Schmidt, D. C. Girvin, and S. C. Smith. 1988. Chromium reactions in Geological
Materials. EPRI-EA-5741. Electric Power Research Institute, Palo Alto, California.
Rai, D., J. M. Zachara, L. E. Eary, D. C. Girvin, D. A. Moore, C. T. Resch, B. M. Sass, and R. L. Schmidt.
1986. Geochemical Behavior of Chromium Species. EPRI-EA-4544. Electric Power Research Institute,
Palo Alto, California.
Ramirez, L. M., J. B. Rodriguez and F. Barba. 1985. "Heavy Metal Concentration in Sludge-Soil Systems
as a result of Water Infiltration." In Tropical Hydrology and Caribbean Island Water Resources
Congress, F. Quinones and A. N. Sanchez (eds.), pp. 20-25, American Water Resources Association,
Bethesda, Maryland.
Rhoades, J. D. 1996. "Salinity: electrical Conductivity and Total Dissolved Solids." In Methods of Soil
Analysis, Part 3, Chemical Methods, J. M. Bigham (ed.), pp. 417-436. Soil Science Society of America
Inc. Madison, Wisconsin.
Richards, L. A. 1954. Diagnosis and Improvement of Saline and Alkali Soils. Agricultural Handbook 60,
U. S. Department of Agriculture, Washington, D.C.
Selim, H. M. and M C. Amcher. 1988. "A Second-Order Kinetic Approach for Modeling Solute Retention
and transport in Soils." Water Resources Research, 24:2061-2075.
E.13
-------
Sheppard, M. I., D. H. Thibault, and J. H. Mitchell. 1987. "Element Leaching and Capillary Rise
in Sandy Soil Cores: Experimental Results." Journal of Environmental Quality, 16:273-284.
Sheppard, M. I., and S. C. Sheppard. 1987. "A Solute Transport Model Evaluated on Two Experimental
Systems." Ecological Modeling, 37:191-206.
Stollenwerk, K. G., and D. B. Grove. 1985. "Adsorption and Desorption of Hexavalent Chromium in an
Alluvial Aquifer Near Telluride, Colorado." Journal of Environmental Quality, 14:150-155.
Wong, K. V., S. Sengupta, D. Dasgupta, E. L. Daly, N. Nemerow, and H. P. Gerrish. 1983. "Heavy Metal
Migration in Soil-Leachate Systems." Biocycle, 24:30-33.
E. 14
-------
APPENDIX F
Partition Coefficients For Lead
-------
Appendix F
Partition Coefficients For Lead
F.1.0 Background
The review of lead Kd data reported in the literature for a number of soils led to the following
important conclusions regarding the factors which influence lead adsorption on minerals, soils,
and sediments. These principles were used to evaluate available quantitative data and generate a
look-up table. These conclusions are:
• Lead may precipitate in soils if soluble concentrations exceed about 4 mg/1 at pH 4 and
about 0.2 mg/1 at pH 8. In the presence of phosphate and chloride, these solubility limits
may be as low as 0.3 mg/1 at pH 4 and 0.001 mg/1 at pH 8. Therefore, in experiments in
which concentrations of lead exceed these values, the calculated Kd values may reflect
precipitation reactions rather than adsorption reactions.
• Anionic constituents such as phosphate, chloride, and carbonate are known to influence
lead reactions in soils either by precipitation of minerals of limited solubility or by
reducing adsorption through complex formation.
• A number of adsorption studies indicate that within the pH range of soils (4 to 11), lead
adsorption increases with increasing pH.
• Adsorption of lead increases with increasing organic matter content of soils.
• Increasing equilibrium solution concentrations correlates with decreasing lead adsorption
(decrease in Kd).
Lead adsorption behavior on soils and soil constituents (clays, oxides, hydroxides,
oxyhydroxides, and organic matter) has been studied extensively. However, calculations by
Rickard and Nriagu (1978) show that the solution lead concentrations used in a number of
adsorption studies may be high enough to induce precipitation. For instance, their calculations
show that lead may precipitate in soils if soluble concentrations exceed about 4 mg/1 at pH 4 and
about 0.2 mg/1 at pH 8. In the presence of phosphate and chloride, these solubility limits may be
as low as 0.3 mg/1 at pH 4 and 0.001 mg/1 at pH 8. Therefore, in experiments in which
concentrations of lead exceed these values, the calculated Kd values may reflect precipitation
reactions rather than adsorption reactions.
Based on lead adsorption behavior of 12 soils from Italy, Soldatini etal. (1976) concluded that
soil organic matter and clay content were 2 major factors which influence lead adsorption. In
these experiments, the maximum adsorption appeared to exceed the cation exchange capacity
F.2
-------
(CEC) of the soils. Such an anomaly may have resulted from precipitation reactions brought
about by high initial lead concentrations used in these experiments (20 to 830 mg/1).
Lead adsorption characteristics of 7 alkaline soils from India were determined by Singh and
Sekhon (1977). The authors concluded that soil clay, organic matter, and the calcium carbonate
influenced lead adsorption by these soils. However, the initial lead concentrations used in these
experiments ranged from 5 to 100 mg/1, indicating that in these alkaline soils the dominant lead
removal mechanism was quite possibly precipitation.
In another adsorption study, Abd-Elfattah and Wada (1981) measured the lead adsorption
behavior of 7 Japanese soils. They concluded that soil mineral components which influenced
lead adsorption ranged in the order: iron oxides>halloysite>imogolite, allophane>humus,
kaolinite>montmorillonite. These data may not be reliable because high lead concentrations (up
to 2,900 mg/1) used in these experiments may have resulted in precipitation reactions
dominating the experimental system.
Anionic constituents, such as phosphate, chloride, and carbonate, are known to influence lead
reactions in soils either by precipitation of minerals of limited solubility or by reducing
adsorption through complex formation (Rickard and Nriagu, 1978). A recent study by Bargar el
al. (1998) showed that chloride solutions could induce precipitation of lead as solid PbOHCl.
Presence of synthetic chelating ligands such as ethylenediaminetetraacetic acid (EDTA) has been
shown to reduce lead adsorption on soils (Peters and Shem, 1992). These investigators showed
that the presence of strongly chelating EDTA in concentrations as low as 0.01 M reduced Kd for
lead by about 3 orders of magnitude. By comparison quantitative data is lacking on the effects
of more common inorganic ligands (phosphate, chloride, and carbonate) on lead adsorption on
soils.
A number of adsorption studies indicate that within the pH range of soils (4 to 11), lead
adsorption increases with increasing pH (Bittel and Miller, 1974; Braids et al., 1972; Griffin and
Shimp, 1976; Haji-Djafari et al., 1981; Hildebrand and Blum, 1974; Overstreet and
Krishnamurthy, 1950; Scrudato andEstes, 1975; Zimdahl andHassett, 1977). Griffin and Shimp
(1976) also noted that clay minerals adsorbing increasing amounts of lead with increasing pH
may also be attributed to the formation of lead carbonate precipitates which was observed when
the solution pH values exceeded 5 or 6.
Solid organic matter such as humic material in soils and sediments are known to adsorb lead
(Rickard and Nriagu, 1978; Zimdahl and Hassett, 1977). Additionally, soluble organic matter
such as fulvates and amino acids are known to chelate soluble lead and affect its adsorption on
soils (Rickard and Nriagu, 1978). Gerritse et al. (1982) examined the lead adsorption properties
of soils as a function of organic matter content of soils. Initial lead concentrations used in these
experiments ranged from 0.001 to 0.1 mg/1. Based on adsorption data, the investigators
expressed Kd value for a soil as a function of organic matter content (as wt.%) and the
distribution coefficient of the organic matter. The data also indicated that irrespective of soil
organic matter content, lead adsorption increased with increasing soil pH (from 4 to 8). In
F.3
-------
certain soils, lead is also known to form methyl- lead complexes (Rickard and Nriagu, 1978).
However, quantitative relationship between the redox status of soils and its effect on overall
lead adsorption due to methylation of lead species is not known.
Tso (1970), and Sheppard etal. (1989) studied the retention of 210Pb in soils and its uptake by
plants. These investigators found that lead in trace concentrations was strongly retained on soils
(high Kd values). Lead adsorption by a subsurface soil sample from Hanford, Washington was
investigated by Rhoads el al. (1992). Adsorption data from these experiments showed that Kd
values increased with decreasing lead concentrations in solution (from 0.2 mg/1 to 0.0062 mg/1).
At a fixed pH of 8.35, the authors found that Kd values were log-linearly correlated with
equilibrium concentrations of lead in solution. Calculations showed that if lead concentrations
exceeded about 0.207 mg/1, lead-hydroxycarbonate (hydrocerussite) would probably precipitate
in this soil.
The Kd data described above are listed in Table F.l.
F.2.0 Approach
The initial step in developing a look-up table consisted of identifying the key parameters which
were correlated with lead adsorption (Kd values) on soils and sediments. Data sets developed by
Gerritse et al. (1982) and Rhoads el al. (1992) containing both soil pH and equilibrium lead
concentrations as independent variables were selected to develop regression relationships with
Kd as the dependent variable. From these data it was found that a polynomial relationship
existed between Kd values and soil pH measurements. This relationship (Figure F.l) with a
correlation coefficient of 0.971 (r2) could be expressed as:
The relationship between equilibrium concentrations of lead and Kd values for a Hanford soil at a
fixed pH was expressed by Rhoads et al. (1992) as:
where C is the equilibrium concentration of lead in |ig/l. The look-up table (Table F.2) was
developed from using the relationships F.l and F.2. Four equilibrium concentration and 3 pH
categories were used to estimate the maximum and minimum Kd values in each category. The
relationship between the Kd values and the 2 independent variables (pH and the equilibrium
concentration) is shown as a 3-dimensional surface (Figure F.2). This graph illustrates that the
highest Kd values are encountered under conditions of high pH values and very low equilibrium
lead concentrations and in contrast, the lowest Kd values are encountered under lower pH and
higher lead concentrations. The Kd values listed in the look-up table encompasses the ranges of
pH and lead concentrations normally encountered in surface and subsurface soils and sediments.
Kd (ml/g) = 1639 - 902.4(pH) + 150.4(pH)2
(F.l)
Kd (ml/g) = 9,550 C0335
(F.2)
F.4
-------
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F.5
-------
Figure F.l. Correlative relationship between Kd and pH.
F.6
-------
Figure F.2. Variation of Kd as a function of pH and the equilibrium lead
concentrations.
F.7
-------
F.3.0 Data Set for Soils
The data sets developed by Gerritse et al. (1982) and Rhoads et al. (1992) were used to
develop the look-up table (Table F.2). Gerritse et al. (1982) developed adsorption data for
2 well-characterized soils using a range of lead concentrations ( 0.001 to 0.1 mg/1) which
precluded the possibility of precipitation reactions. Similarly, adsorption data developed by
Rhoads et al. (1992) encompassed a range of lead concentrations from 0.0001 to 0.2 mg/1 at a
fixed pH value. Both these data sets were used for estimating the range of Kd values for the
range of pH and lead concentration values found in soils.
Table F.2. Estimated range of Kd values for lead as a function of soil pH, and
equilibrium lead concentrations.
Equilibrium Lead
Concentration (|j,g/l)
K„ (ml/g)
Soil pH
4.0 - 6.3
6.4 - 8.7
8.8-11.0
0.1 -0.9
Minimum
940
4,360
11,520
Maximum
8,650
23,270
44,580
1.0-9.9
Minimum
420
1,950
5,160
Maximum
4,000
10,760
20,620
10-99.9
Minimum
190
900
2,380
Maximum
1,850
4,970
9,530
100-200
Minimum
150
710
1,880
Maximum
860
2,300
4,410
F.8
-------
F.4.0 References
Abd-Elfattah, A., and K. Wada. 1981. "Adsorption of Lead, Copper, Zinc, Cobalt, and
Cadmium by Soils that Differ in Cation-Exchange Material." Journal of Soil Science,
32:71-283.
Bargar, J. R., G. E. Brown, Jr., and G. A. Parks. 1998. "Surface Complexation of Pb(II) at
Oxide-Water Interface: III. XAFS Determination of Pb(II) and Pb(II)-Chloro Adsorption
Complexes on Goethite and Alumina." Geochimica et Cosmochimica Acta, 62(2): 193-207.
Bittel, J. R., and R. J. Miller. 1974. "Lead, Cadmium, and Calcium Selectivity Coefficients on
Montmorillonite, Illite, and Kaolinite." Journal of Environmental Quality, 3:250-253.
Braids, O. C., F. J. Drone, R. Gadde, H. A. Laitenen, and J. E. Bittel. 1972. "Movement of Lead
in Soil-Water System." In Environmental Pollution of Lead and Other Metals, pp 164-238,
University of Illinois, Urbana, Illinois.
Chow, T. J. 1978. "Lead in Natural Waters." In The Biogeochemistry of Lead in the
Environment. Part A. Ecological Cycles., J. O. Nriagu (ed.), pp. 185-218, Elsevier/North
Holland, New York, New York.
Forbes, E. A., A. M. Posner, and J. P. Quirk. 1976. "The Specific Adsorption of Cd, Co, Cu,
Pb, and Zn on Goethite." Journal of Soil Science, 27:154-166.
Gerritse, R. G., R. Vriesema, J. W. Dalenberg, and H. P. De Roos. 1982. "Effect of Sewage
Sludge on Trace Element Mobility in Soils." Journal of Environmental Quality,
11:359-364.
Grasselly, G., and M. Hetenyi. 1971. "The Role of Manganese Minerals in the Migration of
Elements." Society of Mining Geology of Japan, Special Issue 3:474-477.
Griffin, R. A., and N. F. Shimp. 1976. "Effect of pH on Exchange-Adsorption or Precipitation
of Lead from Landfill Leachates by Clay Minerals." Environmental Science and
Technology, 10:1256-1261.
Haji-Djafari, S., P. E. Antommaria, and H. L. Crouse. 1981. "Attenuation of Radionuclides and
Toxic Elements by In Situ Soils at a Uranium Tailings Pond in central Wyoming." In
Permeability and Groundwater Contaminant Transport, T. F. Zimmie, and C. O. Riggs
(eds.), pp 221-242. ASTM STP 746. American Society of Testing Materials. Washington,
DC.
Hildebrand, E. E., and W. E. Blum. 1974. "Lead Fixation by Clay Minerals."
Naturewissenschaften, 61:169-170.
F.9
-------
Leckie, J. O., M. M. Benjamin, K. Hayes, G. Kaufman, and S. Altman. 1980.
Adsorption/Coprecipitation of Trace Elements from Water with Iron Oxyhydroxides.
EPRI-RP-910, Electric Power Research Institute, Palo Alto, California.
Overstreet, R., and C. Krishnamurthy. 1950. "An Experimental Evaluation of Ion-exchange
Relationships." Soil Science, 69:41-50.
Peters, R. W., and L. Shem. 1992. "Adsorption/Desorption Characteristics of Lead on Various
Types of Soil." Environmental Progress, 11:234-240.
Rhoads, K., B. N. Bjornstad, R. E. Lewis, S. S. Teel, K. J. Cantrell, R. J. Serne, J. L. Smoot, C.
T. Kincaid, and S. K. Wurstner. 1992. Estimation of the Release andMigration of Lead
Through Soils and Groundwater at the Hanford Site 218-E-12B Burial Ground. Volume 1:
Final Report. PNL-8356 Volume 1, Pacific Northwest Laboratory, Richland, Washington.
Rhoades, J. D. 1996. "Salinity: electrical Conductivity and Total Dissolved Solids." In Methods
of Soil Analysis, Part 3, Chemical Methods, J. M. Bigham (ed.), pp. 417-436. Soil Science
Society of America Inc., Madison, Wisconsin.
Richards, L. A. 1954. Diagnosis and Improvement of Saline and Alkali Soils. Agricultural
Handbook 60, U. S. Department of Agriculture, Washington, D.C.
Rickard, D. T., and J. E. Nriagu. 1978. "Aqueous Environmental Chemistry of Lead." In The
Biogeochemistry of Lead in the Environment. Part A. Ecological Cycles, J. O. Nriagu (ed.),
pp. 291-284, Elsevier/North Holland, New York, New York.
Scrudato, R. J., and E. L. Estes. 1975. "Clay-Lead Sorption Studies." Environmental Geology,
1:167-170.
Sheppard, S. C., W. G. Evenden, and R. J. Pollock. 1989. "Uptake of Natural Radionuclides by
Field and Garden Crops." Canadian Journal of Soil Science, 69:751-767.
Singh, B, and G. S. Sekhon. 1977. "Adsorption, Desorption and Solubility Relationships of
Lead and Cadmium in Some Alkaline Soils." Journal of Soil Science, 28:271-275.
Soldatini, G. F., R. Riffaldi, and R. Levi-Minzi. 1976. "Lead adsorption by Soils." Water, Air
and Soil Pollution, 6:111-128.
Tso, T.C. 1970. "Limited Removal of 210Po and 210Pb from Soil and Fertilizer Leaching."
Agronomy Journal, 62:663 -664.
Zimdahl, R. L., and J. J. Hassett. 1977. "Lead in Soil." In Lead in the Environment. W. R.
Boggess andB. G. Wixson (eds.), pp. 93-98. NSF/RA-770214. National Science
Foundation, Washington, D.C.
F.10
-------
APPENDIX G
Partition Coefficients For Plutonium
-------
Appendix G
Partition Coefficients For Plutonium
G.1.0 Background
A number of studies have focussed on the adsorption behavior of plutonium on minerals, soils,
and other geological materials. A review data from diverse literature sources indicated that Kd
values for plutonium typically range over 4 orders of magnitude (Thibault et al., 1990). Also,
from these data a number of factors which influence the adsorption behavior of plutonium have
been identified. These factors and their effects on plutonium adsorption on soils and sediments
were used as the basis for generating a look-up table. These factors are:
• Typically, in many experiments, the oxidation state of plutonium in solution was not
determined or controlled therefore it would be inappropriate to compare the Kd data
obtained from different investigations.
• In natural systems with organic carbon concentrations exceeding -10 mg/kg, plutonium
exists mainly in trivalent and tetravalent redox states. If initial plutonium concentrations
exceed ~10"7 M, the measured Kd values would reflect mainly precipitation reactions and
not adsorption reactions.
• Adsorption data show that the presence of ligands influence plutonium adsorption onto
soils. Increasing concentrations of ligands decrease plutonium adsorption.
• If no complexing ligands are present plutonium adsorption increases with increasing pH
(between 5.5 and 9.0).
• Plutonium is known to adsorb onto soil components such as aluminum and iron oxides,
hydroxides, oxyhydroxides, and clay minerals. However, the relationship between the
amounts of these components in soils and the measured adsorption of plutonium has not
been quantified.
Because plutonium in nature can exist in multiple oxidation states (III, IV, V, and VI), soil redox
potential would influence the plutonium redox state and its adsorption on soils. However, our
literature review found no plutonium adsorption studies which included soil redox potential as a
variable. Studies conducted by Nelson et al. (1987) and Choppin and Morse (1987) indicated
that the oxidation state of dissolved plutonium under natural conditions depended on the
colloidal organic carbon content in the system. Additionally, Nelson et al (1987) also showed
that plutonium precipitation occurred if the solution concentration exceeded 10"7 M.
A number of investigators have examined potential adsorption of plutonium on minerals, soils,
and other geological substrates. Earlier experiments conducted by Evans (1956), Tamura
(1972), Van Dalen et al. (1975) showed that plutonium adsorption onto mineral surfaces was
-------
influenced significantly by the type of mineral, the pH and mineral particle size. The reported
values ranged from zero for quartz (Tamura, 1972) to 4,990 ml/g for montmorillonite (Evans,
1956). [The Kd for glauconite tabulated by Evans (1956) was listed as "infinite"(certainly greater
than 5,000 ml/g), because the concentration of dissolved plutonium measured in the Kd
defemination was below detection.] These Kd values are only qualitative because, the initial
concentrations of plutonium used in these experiments were apparently high enough to induce
precipitation of plutonium solid phases therefore, the observed phenomena was likely due to
mainly precipitation and not adsorption. Second, the redox status of plutonium was unknown in
these experiments thus these reported Kd values cannot be Kd readily compared to values derived
from other experiments.
The importance of the plutonium redox status on adsorption was demonstrated by Bondietti et al.
(1975) who reported about 2 orders of magnitude difference in Kd values between hexavalent
(250 ml/g) and tetravalent (21,000 ml/g) plutonium species adsorbing on to montmorillonite.
Bondietti et al. (1975) also demonstrated that natural dissolved organic matter (fulvic acid)
reduces plutonium from hexavalent to tetravalent state thus potentially affecting plutonium
adsorption in natural systems. Some of the earlier adsorption experiments also demonstrated that
complexation of plutonium by various ligands significantly influences its adsorption behavior.
Increasing concentrations of acetate (Rhodes, 1957) and oxalate (Bensen, 1960) ligands resulted
in decreasing adsorption of plutonium. Adsorption experiments conducted more recently
(Sanchez et al., 1985) indicate that increasing concentrations of carbonate ligand also depresses
the plutonium adsorption on various mineral surfaces.
Even though the adsorption behavior of plutonium on soil minerals such as glauconite (Evans,
1956), montmorillonite (Billon, 1982; Bondietti et al., 1975), attapulgite (Billon, 1982), and
oxides, hydroxides, and oxyhydroxides (Evans, 1956; Charyulu et al., 1991; Sanchez et al.,
1985; Tamura, 1972; Ticknor, 1993; Van Dalen et al., 1975) has been studied, correlative
relationships between the type and quantities of soil minerals in soils and the overall plutonium
adsorption behavior of the soils have not been established.
Adsorption experiments conducted by Billon (1982) indicated Kd values for Pu(IV) ranging from
about 32,000 to 320,000 ml/g (depending on pH) for bentonite or attapulgite as adsorbents.
Because of relatively high initial concentrations of plutonium [1.7xl0"6 to 4xlO"6M of Pu(IV)]
used in these experiments, it is likely that precipitation and not adsorption resulted in very high
Kd values. Additional experiments conducted with Pu(VI) species on bentonite substrate
resulted in Kd values ranging from about 100 to 63,100 ml/g when pH was varied from 3.1 to
7.52. The validity of these data are questionable because of high initital concentrations of
plutonium used in these experiments may have induced precipitation of plutonium.
Experiments conducted by Ticknor (1993) showed that plutonium sorbed on goethite and
hematite from slightly basic solutions [(pH: 7.5) containing high dissolved salts, but extremely
low bicarbonate concentrations (8.2 x 10"6 to 2.9 x 10"4M)] resulted in distribution coefficients,
Kd, ranging from 170 to 1,400 ml/g. According to Pius et al. (1995), significant removal of
Pu(IV) from solutions containing 0.1 to 1 M concentrations of sodium carbonate was observed
with alumina, silica gel, and hydrous titanium oxide as substrates. These investigators also noted
that the presence of carbonate lowered the sorption distribution coefficient for these adsorbents.
-------
However, even at 0.5 M carbonate, the coefficients were 60 ml/g, 1,300 ml/g, and 15,000 ml/g,
respectively, for alumina, silica gel, and hydrous titanium oxide. In another study using
bicarbonate solutions, the distribution coefficient for Pu(IV) sorption on alumina was lowered to
about 30 ml/g at 0.5 M bicarbonate (Charyulu el al., 1991). However, one should note that the
initial concentrations of Pu(IV) used by these investigators ranged from 8.4 x 10"6 to 4.2 x 10"5
M, which means that the solutions were probably supersaturated with respect to Pu02xH20 solid
phase. Because of the experimental conditions used by Pius et al. (1995) and Charyulu et al.
(1991), the principal mechanism of plutonium removal from solution could have been
precipitation as easily as adsorption.
Barney et al. (1992) measured adsorption of plutonium from carbonate-free wastewater solutions
onto commercial alumina adsorbents over a pH range of 5.5 to 9.0. Plutonium adsorption Kd
values increased from about 10 ml/g at a pH of 5.5 to about 50,000 ml/g at a pH of 9.0. The
slopes of the Kd compared to the pH curves were close to 1, which indicated that 1 hydrogen ion
is released to the solution for each plutonium ion that is adsorbed on the alumina surface. This
behavior is typical of adsorption reactions of multivalent hydrolyzable metal ions with oxide
surfaces. Changing the initial concentration of plutonium from about 10"9 to 10"10 M did not
affect the Kd values, which showed that plutonium precipitation was not significant in these tests.
Also, the initial plutonium concentrations were below the measured solubility limits of
plutonium hydroxide. This experiment demonstrated that in carbonate-free systems, plutonium
would be adsorbed on alumina substrates.
Another study of adsorption of Pu(IV) and Pu(V) on goethite was conducted by Sanchez et al.
(1985). The experimental conditions used by these investigators were evaluated for assessing
whether the reaction being studied was indeed adsorption. The initial plutonium concentrations
used in their experiments were 10"10 and 10"11 moles per liter. These concentrations are well
below the equilibrium saturation levels for Pu02xH20. The equilibrating solutions used in these
experiments contained salts such as NaN03, NaCl, Na2S04, and NaHC03 and did not contain any
ionic constituents that may have potentially formed solid solution precipitates. Therefore, it is
reasonably certain that the dominant reaction being studied was adsorption and not precipitation
of pure or solid solution phases.
The Pu(IV) and (V) adsorption data obtained in 0.1 M NaN03 electrolyte medium by Sanchez el
al. (1985) indicated isotherms typical of metal and/or metal-like complex specie adsorption on
substrate (Benjamin and Leckie, 1981). This indicated that Pu(IV) and Pu(V) adsorbed onto the
ionized hydroxyl sites in the form of free ions and their hydrolytic species with metal ion and the
metal-ion part of the complexes adsorbing onto the surface. The adsorption isotherms obtained
at the higher initial concentration (10"10 M) of total soluble Pu(IV) and Pu(V) showed that the
adsorption edges (pH value at which 50 percent adsorption occurs) increased towards a higher
pH value, which is typical of the metal-like adsorption behavior of adsorbing species (Benjamin
and Leckie, 1981). These data also showed that the adsorption edges for Pu(V) was shifted
about 2 pH units higher as compared to the adsorption edges observed for Pu(V), indicating that
plutonium in the higher oxidation state (pentavalent) had lower adsorbing affinity as compared
with tetravalent plutonium. This difference in adsorption was attributed to the fact that Pu(V)
hydrolyzes less strongly than Pu(IV),
-------
The Pu(IV) and Pu(V) adsorption data obtained in 0.1 M NaN03 media represents conditions
where only free cations and the respective hydrolytic species are the adsorbing species.
Extensive experimental observations have shown that, when present, strong complexing agents
have a significant effect on the metal ion adsorption (Benjamin and Leckie, 1981). This
modified adsorption behavior in the presence of complex-forming ligands is characterized by
Benjamin and Leckie as ligand-like adsorption. Sanchez et al. (1985) also conducted
experiments to examine the effect of dissolved carbonate (from 10 to 1,000 meq/1) on the
adsorption of Pu(IV) and Pu(V) on goethite. Their adsorption data showed that at a fixed pH
value of 8.6, increasing carbonate concentration beyond 100 meq/1 greatly decreased the
adsorption of plutonium in both oxidation states. These data demonstrated that practically no
Pu(IV) or Pu(V) adsorption occurred on goethite when the total carbonate concentration
approached 1,000 meq/1 (0.5 M C03). However, data collected by Glover et al. (1976) showed
that, at very low concentrations of dissolved carbonate {i.e., 0.1-6 meq/1) typically encountered
in soils, adsorption of Pu(IV) increased with increasing dissolved carbonate concentration.
These results indicate that Pu(IV) in these soils may adsorb in the form of PuHCO^ species.
Such complete suppression of Pu(IV) and Pu(V) adsorption was attributed to the presence of
anionic plutonium-hydroxy carbonate species in solution and to the fact that goethite at this pH
contains mainly negatively charged sites that have negligible affinity to adsorb anionic species.
This adsorption behavior of Pu(IV) and Pu(V) in the presence of carbonate ions that form strong
hydroxy carbonate complexes is typical of ligand-like adsorption of metal ions described by
Benjamin and Leckie (1981). Ligand-like adsorption is described as adsorption of a metal-
ligand complex that is analogous to adsorption of the free ligand species. Also, the metal-ligand
complexes may not adsorb at all if these complexes are highly stable. These data clearly
demonstrate that increasing total carbonate and hydroxyl solution concentrations significantly
decrease Pu(IV) and Pu(V) on iron oxyhydroxide surfaces.
Similar suppression of adsorption of higher valence state actinides in the presence of carbonate
and hydroxyl ions has been observed by a number of investigators. Some of these studies
include adsorption of U(VI) on goethite (Hsi and Langmuir, 1985; Koehler et al., 1992; Tripathi,
1984), ferrihydrite (Payne et al., 1992), and clinoptilolite (Pabalan and Turner, 1992), and Np(V)
adsorption on ferrihydrite, hematite, and kaolinite (Koehler etal., 1992).
Some of the early plutonium adsorption experiments on soils were conducted by Rhodes (1957)
and Prout (1958). Rhodes (1957) conducted plutonium adsorption experiments using a
calcareous subsurface soil from Hanford as the adsorbent. The data indicated that adsorption
varied as a function of pH ranging from 18 ml/g under highly acidic conditions to >1980 ml/g at
highly alkaline conditions. These data are unreliable because initial plutonium concentration of
6.8xl0"7 M used in these experiments may have resulted in precipitation of plutonium solid
phases. Prout (1958) studied adsorption of plutonium in +3, +4, and +6 redox states on a
Savannah River Plant soil as a function of pH. The calculated Kd ranged from <10 to >10,000
ml/g, -100 to -10,000 ml/g, and <10 to -3,000 ml/g for Pu(III), Pu(IV), and Pu(VI) respectively.
Maximum Kd values were observed between pH values of about 6.5 and 8.5. Because the initial
concentrations of plutonium used in these experiments were about lxlO"6 M, precipitation
reaction may have accounted for the observed removal of plutonium from solution phase.
-------
Bondietti et al. (1975) conducted Pu(IV) adsorption studies with the clay fraction isolated from a
silt loam soil as the adsorbent. The Kd values from these experiments were reported be as high
as 1.04xl06 and 1.68x10s ml/g . Experiments conducted by Dahlman et al (1976) also showed
exceedingly high Kd value (3x10s ml/g) for Pu(IV) adsorption on clay fraction from a silt loam
soil. In view of this anomalously high Kd value, the authors concluded that actual mechanism of
plutonium removal from solution phase may have been the precipitation reaction.
Nishita et al. (1976) extracted plutonium from a contaminated clay loam soil with solutions
ranging in pH from 1.21 to 13.25. The solution pH in these experiments were adjusted with nitric
acid and sodium hydroxide. The calculated Kd from these experiments varied from 3.02 to 3,086
ml/g, with highest Kd values noted within the pH range of 4.7 to 7.1. In another set of
experiments Nishita (1978) extracted plutonium from the same clay loam soil with acetate (a
ligand which forms complexes with plutonium) containing extraction solutions. The pH values
for these set of extractions ranged from 2.81 to 11.19. The calculated Kd values in this
experiment ranged from 37 to 2,857 ml/g with highest Kd values being observed between pH
values 8.6 to 9.7.
Plutonium adsorption on 14 soil samples obtained from 7 different U.S. Department of Energy
(DOE) sites were studied by Glover et al (1976). Initial concentrations of plutonium in these
experiments were 10"8,10"7, and 10"6M, respectively. The observed Kd values ranged from 30 to
14,000 ml/g. It is likely that removal of plutonium observed under higher initial concentrations
(10"7, and 10"6M) may have been due to precipitation reactions and not from adsorption
reactions.
Rodgers (1976) conducted plutonium adsorption studies on clay and silt fractions from a glacial
till soil from DOE's Mound Facility in Ohio. He noted that Kd values ranged from about 50 to
166,700 ml/g. The highest Kd values were observed between pH values of 5 to 6.
The effects of strong chelating agents such as ethylenediaminetetraacetic acid (EDTA) and
diethylenetriaminepentacetic acid (DTPA) on Pu(IV) adsorption by 3 different soils were
investigated by Relyea and Brown (1978). The soils used for the adsorption were a sand (Fuquay
from South Carolina), a loamy sand (Burbank from Washington), and a silt loam (Muscatine
from Illinois) with initial concentrations of Pu(IV) fixed at about 5xl0"8M. Without the
chelating ligands, the Kd values were 316, 6,000, and 8,000 ml/g for the sand, the loamy sand,
and the silt loam respectively. When 10"3 M of EDTA was present in the matrix solution, the
measured Kd values were 120, 94.5 and 338 ml/g for the sand, the loamy sand, and the silt loam
respectively. These significant reductions in adsorption were attributed to the limited affinity of
Pu-EDTA complexes to adsorb onto the soil mineral surfaces. Increasing the EDTA
concentration by an order of magnitude resulted in reductions in Kd values from about 1 order
(for silt loam) to 2 orders (for sand) of magnitude. Using a stronger chelating agent (10"3 M
DTP A) resulted in very low Kd values (0.12 ml/g for sand, 1.06 ml/g for loamy sand, and 0.24
ml/g for silt loam) which were about 3 to 4 orders of magnitude smaller as compared to the
values from chelate-free systems. The results obtained from desorption experiments (using
EDTA and DTPA ligands) showed that the Kd values were 1 to 2 orders of magnitude higher
than the values calculated from adsorption experiments leading to the conclusion that some
fraction of plutonium in soil was specifically adsorbed (not exchangeable). These data showed
-------
that Pu(IV) adsorption on soils would be significantly reduced if the equilibrating solutions
contain strong chelating ligands, such as EDTA and DTPA.
The reduction of plutonium adsorption on soils by strong synthetic chelating agents was also
confirmed by experiments conducted by Delegard et al. (1984). These investigators conducted
tests to identify tank waste components that could significantly affect sorption of plutonium on
3 typical shallow sediments from the the DOE Hanford Site. They found that sorption was
decreased by the chelating agents, 0.05 M EDTA and 0.1 M HEDTA
(N-2-hydroxyethylethylenediaminetriacetate) but not by low concentrations of carbonate
(0.05 M). Delegard's data also showed that roughly a twofold increase in ionic strength caused
an order of magnitude decrease in plutonium adsorption.
Based on an adsorption study of plutonium on basalt interbed sediments from the vicinity of
Hanford site, Barney (1984) reported a Kd value of about 500 ml/g. This relatively lower Kd
value may have resulted from the relatively enhanced concentration of 215 mg/1 of carbonate
(a complex forming ligand) which was present in the groundwater used in the experiments.
Later, sorption of plutonium in +4, +5, and +6 redox states on a Hanford Site shallow sediment
was studied by Barney (1992) to elucidate any differences in rate and amount of adsorption of
plutonium in different redox states. The initial plutonium concentrations used in these
experiments varied between about 10"11 to 10"9 M with synthetic ground water as a background
electrolyte. The data indicated that the Kd values ranged from 2,100 to 11,600, 2,700 to 4,600,
and 1,000 to 4,600 ml/g for plutonium in +4, +5, and +6 redox states, respectively. The data also
indicated that Pu(V) and Pu(VI) upon adsorption was reduced to the tetravalent state. In these
experiments, the Kd data obtained at lower initial concentrations (~lxl0"n M) of plutonium are
reliable because the dominant plutonium removal mechanism from solution was adsorption.
Using batch equilibration techniques, Bell and Bates (1988) measured Kd values for plutonium
which ranged from 32 to 7,600 ml/g. The soils used in these experiments were obtained from the
Sellafield and Drigg sites in England and their texture ranged from clay to sand. Ground water
spiked with about 2. lxlO"8 M of plutonium was used in these adsorption experiments. The data
also showed that the adsorption of plutonium on these soils varied as a function of pH, with
maximum adsorption occuring at a pH value of about 6.
A number of studies indicate that Kd values for plutonium adsorption on river, oceanic, and lake
sediments range from about lxlO3 to 1x10s ml/g. Duursma and coworkers calculated that Kd for
marine sediments was about lxlO4 ml/g (Duursma and Eisma, 1973; Duursma and Gross, 1971;
Duursma and Parsi, 1974). Studies by Mo and Lowman (1975) on plutonium-contaminated
calcareous sediments in aerated and anoxic seawater medium yielded Kd values from 1.64xl04 to
3.85x 10s ml/g. Based on distribution of plutonium between solution and suspended particle
phases in sea water, Nelson et al. (1987) calculated that for plutonium in oxidized states (V, VI),
the Kd was ~2.5xl03ml/g, and ~2.8xl06 ml/g for plutonium in reduced states (III, IV). Based on
a number of observations of lake and sea water samples, Nelson et al (1987) reported that Kd
values for lake particulates ranged from 3,000 to 4xl05ml/g, and for oceanic particulates ranged
from 1x10s to 4x10s ml/g.
G.2.0 Data Set for Soils
-------
The most detailed data set on plutonium Kd measurements were obtained by Glover el al. (1976).
These data set were based on 17 soil samples from 9 different sites that included 7 DOE sites.
The characterization of the soil included measurements of CEC, electrical conductivity, pH and
soluble carbonate of the soil extracts, inorganic and organic carbon content, and the soil texture
(wt.% of sand, silt, and clay content). The textures of these soils ranged from clay to fine sand.
Three different initial concentrations of plutonium (10"8, 10"7, and 10"6M) were used in these
experiments. This data set is the most extensive as far as the determination of a number of soil
properties therefore, it can be examined for correlative relationships between Kd values and the
measured soil parameters. The data set generated at initial plutonium concentrations of 10"8 M
were chosen for statistical analyses because the data sets obtained at higher initial concentrations
of plutonium may have been affected by precipitation reactions (Table G. 1).
G.3.0 Approach and Regression Models
The most detailed data set on plutonium Kd measurements were obtained by Glover el al. (1976).
This data set was based on 17 soil samples from 9 different sites that included 7 DOE sites. The
characterization of the soil included measurements of CEC, electrical conductivity, pH and
soluble carbonate of the soil extracts, inorganic and organic carbon content, and the soil texture
(wt.% of sand, silt, and clay content). The textures of these soils ranged from clay to fine sand.
Three different initial concentrations of plutonium (10"8, 10"7, and 10"6M) were used in these
experiments. This data set is the most extensive as far as the determination of a number of soil
properties therefore, it can be examined for correlative relationships between Kd values and the
measured soil parameters. The data set generated at an initial plutonium concentration of 10"8 M
was chosen for statistical analyses because the data sets obtained at higher initial concentrations
of plutonium may have been confounded by precipitation reactions
In developing regression models, initially it is assumed that all variables are influential.
However, based on theoretical considerations or prior experience with similar models, one
usually knows that some variables are more important than others. As a first step, all the
variables are plotted in a pairwise fashion to ascertain any statistical relationship that may exist
between these variables. This is typically accomplished by the use of scatter diagrams in which
the relationship of each variable with other variables is examined in a pair-wise fashion and
displayed as a series of 2-dimensional graphs. This was accomplished by using the Statistica™
software. The variables graphed included the distribution coefficient (Kd in ml/g), pH, CEC (in
meq/lOOg), electrical conductivity of soil extract (EC in mmhos/cm), dissolved carbonate
concentration in soil extract (DCARB in meq/1), inorganic carbon content (IC as percent
CaC03), organic carbon content (OC as wt.%), and the clay content (CLAY as wt.%).
-------
Table G.l. Plutonium adsorption data for soil samples. [Data taken from results
reported by Glover et al. (1976) for measurements conducted at an initial
plutonium concentrations of 10"8 M.]
Soil
Sample
K,
(ml/g)
PH
CEC1
(meq/100 g)
EC1
(mmhos/cm)
DCARB1
(meq/1)
IC %'
CaC03
OC1
(%
mass)
CLAY1
(%
mass)
CO-A
2,200
5.7
20.0
3.6
5.97
0.4
2.4
36
CO-B
200
5.6
17.5
0.4
0.97
0.3
3.4
22
CO-C
1,900
7.9
29.6
0.4
1.98
2.4
0.7
64
ID-A
1,700
7.8
15.5
0.5
2.71
17.2
0.8
34
ID-B
320
8.3
13.8
0.8
2.51
7.9
0.2
32
ID-C
690
8.0
8.2
1.0
2.52
5.2
0.3
23
ID-D
2,100
7.5
17.5
1.2
4.90
0.0
0.1
3
WA-A
100
8.0
6.4
0.9
2.60
0.6
0.3
14
WA-B
430
8.2
5.8
0.4
2.30
0.0
0.1
14
SC
280
5.4
2.9
0.4
0.50
0.2
0.7
20
NY
810
5.4
16.0
1.2
1.40
0.0
2.7
36
NM
100
6.4
7.0
1.7
2.80
0.2
0.7
18
AR-A
710
6.2
34.4
0.5
0.10
0.9
3.2
56
AR-B
80
4.8
3.8
0.4
0.10
0.7
0.6
9
AR-C
430
2.3
16.2
0.3
0.10
0.6
2.3
37
IL
230
3.6
17.4
0.5
0.10
0.7
3.6
16
1 CEC: Cation exchange capacity; EC: Electrical conductivity; DCARB: Dissolved
carbonate; IC: Inorganic carbon; OC: Organic carbon; CLAY: Soil clay content.
-------
The scatterplots are typically displayed in a matrix format with columns and rows representing
the dependent and independent variables respectively. For instance, the first row of plots shows
the relationship between Kd as a dependent variable and other variables each in turn as selected
as independent variables. Additionally, histograms displayed in each row illustrate the value
distribution of each variable when it is being considered as the dependent variable.
The scatter matrix (Figure G.l) shows that regression relationships may exist between Kd and
CEC, DCARB, and CLAY. Other relationships may exist between the CEC and CLAY,
DCARB, and between PH, EC and DCARB. These relationships affirm that the CEC of soils
depends mainly on the clay content. Similarly, the electrical conductivity of a soil solution
depends on total concentrations of soluble ions and increasing dissolved carbonate concentration
would contribute towards increasing EC. Also the pH of a soil solution would reflect the
carbonate content of a soil with soils containing solid carbonate tending towards a pH value of
-8.3.
While a scatter diagram is a useful tool to initially assess the pairwise relationships between a
number of variables, this concept cannot be extended to analyze multiple regression relationships
(Montgomery and Peck, 1982). These authors point out that if there is 1 dominant regressive
relationship, the corresponding scatter diagram would reveal this correlation. They also indicate
however, that if several regressive relationships exist between a dependent variable and other
independent variables, or when correlative relationships exist between independent variables
themselves, the scatter diagrams cannot be used to assess multiple regressive relationships.
Typically, in regression model building, significant variables have to be selected out of a number
of available variables. Montgomery and Peck (1982) indicate that regression model building
involves 2 conflicting objectives. First, the models have to include as many independent
variables as possible so that the influence of these variables on the predicted dependent variable
is not ignored. Second, the regression model should include a minimum number of independent
variables as possible so that the variance of predicted dependent variable is minimized.
Variable selection was conducted by using forward stepwise and backward stepwise elimination
methods (Montgomery and Peck, 1982). In the forward stepwise method, each independent
variable is added in a stepwise fashion until an appropriate model is obtained. The backward
stepwise elimination method starts off by including all independent variables and in each step
deletes (selects out) the least significant variables resulting in a final model which includes only
the most influential independent variables.
-------
KD
DilQl= ==~
°s
° o o
° » 0 *> O 0®
o°° °
0 0 0
%© °°°
% °
0 o°
IK 0
»o • °
0 0 0
8°o ¦%,
*•
°o 0
%
° %
0 0 0
° 00
°o°
0 0 0
°°
O °o O o o e
0 °
e
0
0
0
o
o
°o° ° O °°
8 o° °
o
o
° ° £ °o„
= o ° 81
CEC
Mm MM E3D
0
0
"Bo 8
8 * °
0
0
0 °° rf> 0 °
=0 *
0
0
* 0
s °
0
0
0 0 0 0 00
0
0
00 0rf£
_ 00 0
0 0
o
o
0 0 »°
OGOO ° o o
0
o
° °0o
o ° o GO O OO
0
0
00 000
(DO %0 0 O
EC
Oo„
0
0
° 0 °
800 00 0
0
0
°o e 0
°o 0 0
o
0
° Oq O 0
o^o o
o
0
° 3®
3>
o o o ° o
0
0
rfb OO
°0
% 00 0
0
0
(f«» 0
& °
DCARB
Flra.—.Uln
0
0
f, ¦ ¦ 0
\
0
0
0 0
cr 000
0
0
0°o cP
4> 0
0 o° 0 0
o
0
o
o
Ocoo °0 00
o
0
o
o
o o o oo % o
0
0
0
000 ° 0
0
0
0
<§> 0 0 0 0
0
0
0
0
°ooo oQ) 0 0
IC
0
0
0
qO (§ O 0 °00
0
0
0
0
0 CQlD 0 © 0
S o
o 0
O O o o
o Oo o o
° o o
o
o o o o
o°o
0
8 0
Q) 0 O 0
ffo 0 0
O O
0 0
e 0
0
o
o° ° O „°
§»°°
0
o
° °° o°o
• 0"°
0
0
0*00
°y s
0
0
°Oo § 0
•
0
0
0 0
-------
Model Type
Forecasting Equation
R2
Linear Regression
Forward Stepwise
Kd = 284.6 (DCARB) + 27.8 (CLAY) - 594.2
0.7305
Linear Regression
Forward Stepwise
Kd = 488.3 (DCARB) + 29.9 (CLAY) - 119.1 (pH) - 356.8 (EC)
0.8930
Linear Regression
Backward Stepwise
Kd = 284.6 (DCARB) + 27.8 (CLAY) - 594.2
0.7305
Linear Regression
Backward Stepwise
Kd = 351.4 (DCARB)
0.7113
Piecewise Linear
Regression
Kd = 25.7 (DCARB) + 12.14 (CLAY)+ 2.41 for Kd values <767.5
Kd = 286.0 (DCARB) + 21.3(CLAY) - 81.2 for Kd values >767.5
0.9730
Polynomial
Kd = -156.0 (DCARB) + 15.2 (CLAY) +16.1 (DCARB)2 - 0.04 (CLAY)2 + 11.3 (DCARB)(CLAY) - 87.0
0.9222
Polynomial
Kd = -171.1(DCARB) + 10.5 (CLAY) +17.2(DCARB)2 + 0.02 (CLAY)2 + 11.6 (DCARB)(CLAY)
0.9219
Polynomial
Kd = -106.1 (DCARB) + 11.2 (CLAY) + 12.5 (DCARB)(CLAY) - 72.4
0.9194
Polynomial
Kd = -137.9 (DCARB) + 9.3 (CLAY) + 13.4 (DCARB)(CLAY)
0.9190
Table G.3. Estimated range of Kd values for plutonium as a function of the
soluble carbonate and soil clay content values.
K '
Clay Content (wt.%)
0-30
31-50
51-70
Soluble Carbonate
(meq/1)
Soluble Carbonate
(meq/1)
Soluble Carbonate
(meq/1)
0.1-2
3-4
5-6
0.1-2
3-4
5-6
0.1-2
3-4
5-6
Minimum
5
80
130
380
1,440
2,010
620
1,860
2,440
Maximum
420
470
520
1,560
2,130
2,700
1,980
2,550
3,130
-------
Figure G.2. Variation of Kd for plutonium as a function of clay content and
dissolved carbonate concentrations.
-------
G.4.0 References
Barney, G. S. 1984. "Radionuclide Sorption and Desorption Reactions with Interbed Materials
from the Columbia River Basalt Formation." In Geochemical Behavior of Radioactive
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Barney, G. S. 1992. Adsorption of Plutonium on Shallow Sediments at the Hanford Site,
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Bell, J., and T. H. Bates. 1988. "Distribution coefficients of Radionuclides between Soils and
Groundwaters and their Dependence on Various test Parameters." Science of Total
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Benjamin, M. M., and J. O. Leckie. 1981. "Conceptual Model for Metal-Ligand-Surface
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Billon, A. 1982. "Fixation D'elements Transuraniens a Differents Degres D'oxydation Sur Les
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Chemical Society, Washington, D.C.
-------
Delegard, C. H., G. S. Barney, and S. A. Gallagher. 1984. "Effects of Hanford High-Level
Waste Components on the Solubility and Sorption of Cobalt, Strontium, Neptunium,
Plutonium, and Americium. " In Geochemical Behavior of Disposed Radioactive Waste,
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Duursma, E. K., and M. G. Gross. 1971. "Marine Sediments and Radioactivity." In
Radioactivity in the Marine Environment, pp. 147-160, National Academy of Sciences,
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Duursma, E. K., and D. Eisma. 1973. "Theoretical, Experimental and Field Studies Concerning
Reactions of Radioisotopes with Sediments and Suspended Particles of the Sea. Part C:
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Duursma, E. K., and P. Parsi. 1974. "Distribution Coefficients of Plutonium between Sediment
and Seawater." In Activities of the Int. Laboratory of Marine Radioactivity, pp. 94-96,
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Evans, E. J. 1956. Plutonium Retention in Chalk River Soil. CRHP-660. Chalk River
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Glover, P. A., F. J. Miner, and W. O. Polzer. 1976. "Plutonium and Americium Behavior in the
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Proceedings of Actinide-Sediment Reactions Working Meeting, Seattle, Washington.
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Hsi, C. K. D., and D. Langmuir. 1985. "Adsorption of Uranyl onto Ferric Oxyhydroxides:
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Koehler M., E.Wieland, and J. O. Leckie. 1992. "Metal-Ligand Interactions during Adsorption
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Washington, D.C.
Montgomery, D. C., and E. A. Peck. 1982. Introduction to Linear Regression Analysis. John
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Pabalan, R. T., andD. R. Turner. 1992. Sorption Modelingfor HLWPerformance Assessment.
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-------
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-------
APPENDIX H
Partition Coefficients For Strontium
-------
Appendix H
Partition Coefficients For Strontium
H.1.0 Background
Two simplifying assumptions underlying the selection of strontium Kd values included in the
look-up table were made. These assumptions are that the adsorption of strontium adsorption
occurs by cation exchange and follows a linear isotherm. These assumptions appear to be
reasonable for a wide range of environmental conditions. However, these simplifying
assumptions are compromised in systems with strontium concentrations greater than about
10"4 M, humic substance concentrations greater than about 5 mg/1, ionic strengths greater than
about 0.1 M, and pH levels greater than approximately 12.
Based on these assumptions and limitations, strontium Kd values and some important ancillary
parameters that influence cation exchange were collected from the literature and tabulated in
Section H.3. The tabulated data were from studies that reported Kd values (not percent adsorbed
or Freundlich or Langmuir constants) and were conducted in systems consisting of
• Natural soils (as opposed to pure mineral phases)
• Low ionic strength (< 0.1 M)
• pH values between 4 and 10
• Strontium concentrations less than 10"4 M
• Low humic material concentrations (<5 mg/L)
• No organic chelates (such as EDTA)
The ancillary parameters included clay content, pH, CEC, surface area, solution calcium
concentrations, and solution strontium concentrations. The table in Section H.3 describes
63 strontium Kd values. Strontium Kd values for soils as well as pure mineral phases are
tabulated in Section H.4. This table contains 166 entries, but was not used to provide guidance
regarding the selection of Kd values to be included in the look-up table.
Statistical analysis were conducted with the data collected from the literature. These analyses
were used as guidance for selecting appropriate Kd values for the look-up table. The Kd values
used in the look-up tables could not be based entirely on statistical consideration because the
statistical analysis results were occasionally nonsensible. For instance, negative Kd values were
predicted by 1 regression analysis. Thus, the Kd values included in the look-up table were not
selected purely by objective reasoning. Instead, the statistical analysis was used as a tool to
provide guidance for the selection of the approximate range of values to use and to identify
meaningful trends between the strontium Kd values and the soil parameters.
The descriptive statistics of the strontium Kd data set for soil data only (entire data set presented
in Section H.3) is presented in Table H. 1. The 63 strontium Kd values in this data set ranged
H.2
-------
from 1.6 ml/g for a measurement made on a sandy soil dominated by quartz (Lieser etal., 1986)
to 10,200 ml/g for a measurement made on a tuff1 soil collected at Yucca Mountain, Nevada
(Sample YM-38; Vine etal., 1980). The average strontium Kd value was 355 ±184 ml/g. The
median2 strontium Kd value was 15.0 ml/g. This is perhaps the single central estimate of a
strontium Kd value for this data set.
Table H.l. Descriptive statistics of strontium Kd data set for soils.
SrKd
(ml/g)
Clay
Content
(wt.%)
pH
CEC
(meq/100 g)
Surface
Area
(m2/g)
Ca
(mg/1)
Mean
355
7.1
6.8
4.97
1.4
56
Standard Error
183
1.1
0.21
1.21
0
23
Median
15
5
6.7
0.9
1.4
0
Mode
21
5
6.2
2
1.4
0
Standard Deviation
1,458
7.85
1.35
9.66
0.00
134
Kurtosis
34
10.7
-0.5
11.6
-3
3.4
Minimum
1.6
0.5
3.6
0.05
1.4
0.00
Maximum
10,200
42.4
9.2
54
1.4
400
Number of
Observations
63
48
42
63
7.00
32
1 Tuff is a general name applied to material dominated by pyroclastic rocks composed of
particles fragmented and ejected during volcanic eruptions.
2 The median is that value for which 50 percent of the observations, when arranged in order of
magnitude, lie on each side.
H.3
-------
H.2.0 Approach and Regression Models
H.2.1 Correlations with Strontium Kd Values
A matrix of the correlation coefficients of the strontium Kd values and soil parameters are
presented in Table H.2. The correlation coefficients significant at or less than the 5 percent level
of probability (P < 0.05) are identified in Table H.2. The highest correlation coefficient with
strontium Kd values was with CEC (r = 0.84). Also significant are the correlation coefficients
between strontium Kd values and clay content (r = 0.82) and CEC and clay content (r = 0.91)
(Table H.2).
H.2.2 Strontium Kd Values as a Function of CEC andpH
The CEC and strontium Kd data are presented in Figure H. 1. It should be noted that a
logarithmic scale was used for the y-axis to assist in the visualization of the data and is not meant
to suggest any particular model. A great deal of scatter exists in this data, especially in the lower
CEC range where more data exist. For example, between the narrow CEC range of 5.5 to
6.0 meq/100 g, 9 strontium Kd values are reported ( Keren and O'Connor, 1983; McHenry, 1958;
Serne et al., 1993). The strontium Kd values range from 3 ml/g for a surface noncalcareous
sandy loam collected from New Mexico (Keren and O'Connor, 1983) to 70 ml/g for a carbonate
surface soil collected from Washington (McHenry, 1958). Thus, over an order of magnitude
variability in strontium Kd values may be expected at a given CEC level.
Table H.2. Correlation coefficients (r) of the strontium Kd data set for soils.
Strontium
K,
Clay
Content
pH
CEC
Surface
Area
Ca Cone.
Strontium Kd
1.00
Clay Content
0.821
1.00
pH
0.28
0.03
1.00
CEC
p
00
4^
0.911
0.281
1.00
Surface Area
0.00
-1.00
0.00
1.001
1.00
Ca Cone.
-0.17
0.00
-0.20
0.03
1.00
1 Correlation coefficients significant at or less than the 5% level of probability (P < 0.05).
H.4
-------
Figure H.l. Relation between strontium Kd values and
CEC in soils.
Another important issue regarding this data set is that 83 percent of the observations exists at
CEC values less than 15 meq/100 g. The few Kd values associated with CEC values greater than
15 meq/100 g may have had a disproportionally large influence on the regression equation
calculation (Neter and Wasserman, 1974). Consequently, estimates of strontium Kd values using
these data for low CEC soils, such as sandy aquifers, may be especially inaccurate.
The regression equation for the data in Figure H. 1 is presented as Equation 1 in Table H.3. Also
presented in Table H.3 are the 95 percent confidence limits of the calculated regression
coefficients, the y-intercepts, and slopes. These coefficients, when used to calculate Kd values,
suggest a Kd range at a given CEC by slightly over an order of magnitude. The lower 95 percent
confidence limit coefficients can provide guidance in selecting lower (or conservative) Kd
values.
The large negative intercept in Equation 1 compromises its value for predicting strontium Kd
values in low CEC soils, a potentially critical region of the data, because many aquifers matrix
have low CEC values. At CEC values less than 2.2 meq/100 g, Equation 1 yields negative
H.5
-------
strontium Kd values, which are clearly unrealistic.1 To provide a better estimate of strontium Kd
values at low CEC values, 2 approaches were evaluated. First, the data in Figure H. 1 was
reanalyzed such that the intercept of the regression equation was set to zero, i.e., the regression
equation was forced through the origin. The statistics of the resulting regression analysis are
presented as Equation 2 in Table H.3. The coefficient of determination (R2) for Equation 2
slightly decreased compared to Equation 1 to 0.67 and remained highly significant (F= 2xl0"16).
However, the large value for the slope resulted in unrealistically high strontium Kd values. For
example at 1 meq/100 g, Equation 2 yields a strontium Kd value of 114 ml/g, which is much
greater than the actual data presented in Figure H. 1.
The second approach to improving the prediction of strontium Kd values at low CEC was to limit
the data included in the regression analysis to those with CEC less than 15 meq/100 g. These
data are redrawn in Figure H.2. The accompanying regression statistics with the y-intercept
calculated and forced through the origin are presented in Table H.3 as Equations 3 and 4,
respectively. The regression equations are markedly different from there respective equations
describing the entire data set, Equations 1 and 2. Not surprisingly, the equations calculate
strontium Kd more similar to those in this reduced data set. Although the coefficients of
determination for Equations 3 and 4 decreased compared to those of Equations 1 and 2, they
likely represent these low CEC data more accurately.
Including both CEC and pH as independent variables further improved the predictive capability
of the equation for the full data set as well as the data set for soils with CEC less than 15
meq/100 g (Equations 5 and 6 in Table H.3). Multiple regression analyses with additional
parameters did not significantly improve the model (results not presented).
H.2.3 Strontium Kd Values as a Function of Clay Content and pH
Because CEC data are not always available to contaminant transport modelers, an attempt was
made to use independent variables in the regression analysis that are more commonly available
to modelers. Multiple regression analysis was conducted using clay content and pH as
independent variables to predict CEC (Equations 7 and 8 in Table H.3) and strontium Kd values
(Equations 9 and 10 in Table H.3; Figures H.3 and H.4). The values of pH and clay content
were highly correlated to soil CEC for the entire data set (R2 = 0.86) and for those data limited to
CEC less than 15 meq/100 g (R2 = 0.57). Thus, it is not surprising that clay content and pH were
correlated to strontium Kd values for both the entire data set and for those associated with CEC
less than 15 meq/100 g.
1 A negative Kd value is physically possible and is indicative of the phenomena referred to as
anion exclusion or negative adsorption. It is typically and commonly associated with anions
being repelled by the negative charge of permanently charged minerals.
H.6
-------
140
120
15) "100
1 80
I
i-H
CO
60
40
20
0
o o
4 §
A.O U
jo° i
o @0
o
J fi_L
4 6 8 10 12 14
CEC (meq/100 g)
Figure H.2. Relation between strontium Kd values for soils with
CEC values less than 15 meq/100 g.
1000
60
100
H5
W 10
: 1
o
1 :
O :
O
i °
:
b
¦
i J> 8
- T 0
i a O
¦
* ol
*
>o
SOD
0
1 1
•
3
i
i
10 20 30
Clay (%, wt.)
40
50
Figure H.3. Relation between strontium Kd values and
soil clay contents.
H.7
-------
Table H.3. Simple and multiple regression analysis results involving strontium Kd values,
cation exchange capacity (CEC; meq/100 g), pH, and clay content (percent).
95% Confidence Limits1
#
Equation
n2
Data
Range3
Intercept
Slope First
Independent
Parameter
Slope Second
Independent
Parameter
R24
F Value5
Lower
Upper
Lower
Upper
Lower
Upper
1
Kd = -272 + 126(CEC)
63
All
-501
-43
105
147
...
...
0.70
lxlO"17
2
Kd= 114(CEC)
63
All
—
...
95
134
...
...
0.67
2xl0"16
3
Kd = 10.0 + 4.05(CEC)
57
CEC<15
3.32
16.6
2.13
5.96
...
...
0.25
9xl0"5
4
Kd = 5.85 (CEC)
57
CEC<15
—
...
4.25
7.44
...
...
0.12
7xl0"3
5
Kd = -42 +14(CEC) +
2.33(pH)
27
All
-176
91
11.3
18.3
-17.7
22.4
0.77
3xl0"8
6
Kd = 3.53 (CEC) +
1.67(pH)
25
CEC<15
...
...
0.62
6.46
-0.50
3.85
0.34
9xl0"3
7
CEC = -4.45 +
0.70(clay) + 0.60(pH)
27
All
-10.6
1.67
0.59
0.82
-0.30
1.50
0.86
4x10"
8
CEC = 0.40(clay) +
0.19(pH)
25
CEC<15
...
...
0.24
0.56
-0.01
0.40
0.55
lxlO"4
9
Kd = -108 + 10.5(clay) +
11.2(pH)
27
All
-270
53.3
7.32
13.6
-12.5
34.9
0.67
2xl0"6
10
Kd = 3.54(clay) +
1.67(pH)
25
CEC<15
...
...
0.62
6.46
-0.50
3.85
0.34
9xl0"3
11
Clay = 3.36 +
1.12(CEC)
48
All
2.30
4.41
0.97
1.26
...
...
0.84
lxlO"19
12
Clay = 1.34(CEC)
48
All
...
...
1.16
1.51
...
...
0.69
2xl0"13
1 The 95% confidence limits provides the range within which one can be 95% confident that the statistical parameter
exist.
2 The number of observations in the data set.
3 All available observations were included in regression analysis except when noted.
4 R2 is the coefficient of determination and represents the proportion of the total treatment sum of squares accounted for
by regression (1.00 is a perfect match between the regression equation and the data set).
5 The F factor is a measure of the statistical significance of the regression analysis. The acceptable level of significance
is not standardize and varies with the use of the data and the discipline. Frequently, a regression analysis with a F value
of less than 0.05 is considered to describe a significant relationship.
H.8
-------
W
l
U-
Xfl
100
10
.o o oo°o
o
So
o
o
8 o
<*?S
10
pH
Figure H.4. Relation between strontium Kd values and soil pH.
H.2.4 Approach
Two strontium Kd look up tables were created. The first table requires knowledge of the CEC
and pH of the system in order to select the appropriate strontium Kd value (Table H.4). The
second table requires knowledge of the clay content and pH to select the appropriate strontium
Kd value (Table H.5).
A full factorial table was created that included 3 pH categories and 3 CEC categories. This
resulted in 9 cells. Each cell contained a range for the estimated minimum- and maximum Kd
values. A 2 step process was used in selecting the appropriate Kd values for each cell. For the
first step, the appropriate equations in Table H.3 were used to calculate Kd values. The lower
and upper 95 percent confidence limit coefficients were used to provide guidance regarding the
minimum and maximum Kd values. For the 2 lowest CEC categories, Equation 6 in Table H.3
was used. For the highest CEC category, Equation 5 was used. For the second step, these
calculated values were adjusted by "eye balling the data" to agree with the data in Figures
H.2-H.4. It is important to note that some of the look-up table categories did not have any actual
observations, e.g., pH <5 and CEC = 10 to 50 meq/100 g. For these categories, the regression
analysis and the values in adjacent categories were used to assist in the Kd selection process.
H.9
-------
Table H.4. Look-up table for estimated range of Kd values for strontium based on CEC
and pH. [Tabulated values pertain to systems consisting of natural soils (as
opposed to pure mineral phases), low ionic strength (< 0.1 M), low humic
material concentrations (<5 mg/1), no organic chelates (such as EDTA), and
oxidizing conditions.]
K„ (ml/g)
CEC (meq/100 g)
3
3-10
10-50
pH
pH
pH
<5
5-8
8-10
< 5
5-8
8-10
<5
5-8
8-10
Minimum
1
2
3
10
15
20
100
200
300
Maximum
40
60
120
150
200
300
1,500
1,600
1,700
Table H.5. Look-up table for estimated range of Kd values for strontium based on clay
content and pH. [Tabulated values pertain to systems consisting of natural
soils (as opposed to pure mineral phases), low ionic strength (< 0.1 M), low
humic material concentrations (<5 mg/1), no organic chelates (such as
EDTA), and oxidizing conditions.]
Krt (ml/g)
Clay Content (wt.%)
<4%
4 - 20%
20 - 60%
pH
pH
pH
<5
5-8
8-10
< 5
5-8
8-10
<5
5-8
8-10
Minimum
1
2
3
10
15
20
100
200
300
Maximum
40
60
120
150
200
300
1,500
1,600
1,700
H.10
-------
A second look-up table (Table H.5) was created from the first look-up table in which clay
content replaced CEC as an independent variable. This second table was created because it is
likely that clay content data will be more readily available for modelers than CEC data. To
accomplish this, clay contents associated with the CEC values used to delineate the different
categories were calculated using regression equations; Equation 11 was used for the high
category (10 to 50 meq/100 g) and Equation 10 was used for the 2 lower CEC categories. The
results of these calculations are presented in Table H.6. It should be noted that, by using either
Equation 11 or 12, the calculated clay content at 15 meq/100 g of soil equaled 20 percent clay.
Table H.6. Calculations of clay contents using regression equations containing
cation exchange capacity as a independent variable.
Equation1
Y-Intercept
Slope
CEC
(meq/100 g)
Clay Content
(%)
12
...
1.34
3
4
12
...
1.34
15
20
11
3.36
1.1.2
15
20
11
3.36
1.12
50
59
1 Number of equation in Table H.3.
H.l 1
-------
H.3.0 Kd Data Set for Soils
Table H.7 lists the available Kd values identified for experiments conducted with only soils. The Kd
values are listed with ancillary parameters that included clay content, pH, CEC, surface area,
solution calcium concentrations, and solution strontium concentrations.
Table H.7. Strontium Kd data set for soils.
SrK,
(ml/g)
Clay
Content
(%)
PH
CEC
(meq/
100 g)
Surface
Area
(m2/g)
[Ca]
ppm
[Sr]
Background
Solution
Soil
ID
Reference1, Comments
21
0.8
5.2
0.9
1.4
0
*
NaCIO,
Soil A
1, * = 4.4e2Bq/ml 85-Sr in
2.4xl0"8 M SrCl2
19
0.8
5.6
0.9
1.4
0
*
NaCIO,
Soil A
1, * = 4.4e2Bq/ml 85-Sr in
2.4xl0"8 M SrCl2
22
0.8
6.2
0.9
1.4
0
*
NaCIO,
Soil A
1, * = 4.4e2Bq/ml 85-Sr in
2.4xl0"8 M SrCl2
26
0.8
6.45
0.9
1.4
0
*
NaCIO,
Soil A
1, * = 4.4e2Bq/ml 85-Sr in
2.4xl0"8 M SrCl2
24
0.8
6.6
0.9
1.4
0
*
NaCIO,
Soil A
1, * = 4.4e2Bq/ml 85-Sr in
2.4xl0"8 M SrCl2
30
0.8
8.4
0.9
1.4
0
*
NaCIO,
Soil A
1, * = 4.4e2Bq/ml 85-Sr in
2.4xl0"8 M SrCl2
43
0.8
9.2
0.9
1.4
0
*
NaCIO,
Soil A
1, * = 4.4e2Bq/ml 85-Sr in
2.4xl0"8 M SrCl2
21.4
5
0.47
Groundwater
2
25
5
0.83
Groundwater
2, CEC was estimated by
adding exch. Ca,Mg,K
12.7
5
0.39
Groundwater
2, GW = 7.4Ca, 1.7Mg,
2.2Na,5.6Cl, 18ppmS04
7.9
5
0.46
Groundwater
2, Aquifer sediments
15.6
5
0.81
Groundwater
Chalk River Nat'l Lab,
Ottawa, Canada
9.4
5
0.21
Groundwater
2, Described as sand texture
7.6
5
0.25
Groundwater
2, Assumed 5% clay, mean
[clay] in sandy soils
6.4
5
0.24
Groundwater
2
7.7
5
0.26
Groundwater
2
28.1
5
0.76
Groundwater
2
7.63
5
0.26
Groundwater
2
H.12
-------
SrK,
(ml/g)
Clay
Content
(%)
PH
CEC
(meq/
100 g)
Surface
Area
(m2/g)
[Ca]
ppm
[Sr]
Background
Solution
Soil
ID
Reference1, Comments
11.4
5
0.41
Groundwater
2
20.1
5
0.44
Groundwater
2
13
5
0.25
Groundwater
2
9.8
5
0.29
Groundwater
2
11
5
0.22
Groundwater
2
13
5
0.39
Groundwater
2
7.8
5
0.2
Groundwater
2
3.8
5
0.1
Groundwater
2
3
5
0.1
Groundwater
2
2.5
5
0.13
Groundwater
2
4
10
4
5.5
0
lxlO"8M
O.OlMNaCl
Puye
soil-Na
3
15
10
5
5.5
0
lxlO"8M
O.OlMNaCl
Puye
soil-Na
3, Noncalcareous soils
21
10
6
5.5
0
lxlO"8M
O.OlMNaCl
Puye
soil-Na
3
24
10
7.4
5.5
0
lxlO"8M
O.OlMNaCl
Puye
soil-Na
3
3
10
3.6
5.5
400
lxlO"8M
0.01M CaCl
Puye
soil-Ca
3
4.5
10
5.2
5.5
400
lxlO"8M
0.01M CaCl
Puye
soil-Ca
3
5.2
10
6.8
5.5
400
lxlO"8M
O.OlMCaCl
Puye
soil-Ca
3
5.7
10
7.9
5.5
400
lxlO"8M
0.01M CaCl
Puye
soil-Ca
3
3.5
5.2
2
0
lxlO"10M
NaOH/HCl
Hanford
soil
4
4.6
5.6
2
0
lxlO"10M
NaOH/HCl
Hanford
soil
4, Carbonate system
5.8
5.8
2
0
lxlO"10M
NaOH/HCl
Hanford
soil
4
6.1
5.9
2
0
lxlO"10M
NaOH/HCl
Hanford
soil
4
8.3
6
2
0
lxlO"10M
NaOH/HCl
Hanford
soil
4
H.13
-------
SrK,
(ml/g)
Clay
Content
(%)
PH
CEC
(meq/
100 g)
Surface
Area
(m2/g)
[Ca]
ppm
[Sr]
Background
Solution
Soil
ID
Reference1, Comments
17
7.4
2
0
lxlO"10M
NaOH/HCl
Hanford
soil
4
21
7.6
2
0
lxlO"10M
NaOH/HCl
Hanford
soil
4
27
7.8
2
0
lxlO"10M
NaOH/HCl
Hanford
soil
4
47
8.4
2
0
lxlO"10M
NaOH/HCl
Hanford
soil
4
81
9.1
2
0
lxlO"10M
NaOH/HCl
Hanford
soil
4
19.1
4
7.66
10.4
129
100
l_iCi/l
Hanford
Groundwater
cgs-1
5
21.5
6
7.87
5.9
58.5
100
l_iCi/l
Hanford
Groundwater
trench-8
5, Groundwater pH =8.3
23.2
5
8.17
4.57
35.1
100
l_iCi/l
Hanford
Groundwater
tbs-1
5, Hanford, Richland,
Washington surface and
subsurface sediments
48.5
8.24
3
3.8xlO"8M
Yucca
Groundwater
YM-22
6, Los Alamos, New Mexico
10,200
8.17
54
3.8xlO"8M
Yucca
Groundwater
YM-38
6, Yucca Mountain tuff
sediments
2,500
8.13
21
3.8xlO"8M
Yucca
Groundwater
YM48
6, Approximate initial pH,
final pH are presented
3,790
8.24
27
3.8xlO"8M
Yucca
Groundwater
YM-49
6, Final pH 8.1-8.5
3,820
8.24
27
3.8xlO"8M
Yucca
Groundwater
YM-50
6, Sediments = 106-500 |_im
fractions
1.6
0.5
6.2
0.05
10xl0"6M
Groundwater
Sediments
7
2.6
3
6.2
0.3
10xl0"6M
Groundwater
Sediments
7, Added kaolinite to sand
3.4
5
6.2
0.5
10xl0"6M
Groundwater
Sediments
7, CEC estimated based on
kaolinite =10 meq/100 g
4.6
8
6.2
0.8
10xl0"6M
Groundwater
Sediments
7
6.7
13
6.2
1.3
10xl0"6M
Groundwater
Sediments
7
400
42.4
7.2
34
0
Water
Ringhold
Soil
8, soil from Richland,
Washington
H.14
-------
SrK,
(ml/g)
Clay
Content
(%)
PH
CEC
(meq/
100 g)
Surface
Area
(m2/g)
[Ca]
ppm
[Sr]
Background
Solution
Soil
ID
Reference1, Comments
135
26.9
8.3
13.6
0
Water
Bowdoin
Soil
8, soil from Montana
600
33.5
6.5
26.3
0
Water
Hall soil
8, soil from Nebraska
70
3.5
8.3
5.8
0
Water
Composite
Soil
8, soil from Hanford Site,
Richland, Washington
1 References: 1 = Ohnuki, 1994,2 = Patterson and Spoel, 1981; 3 = Keren and O'Connor, 1983; 4 = Rhodes and Nelson, 1957; 5 = Serne
etal., 1993; 6 = Vine etal., 1980; 7 = Lieser and Steinkopff, 1989; 8 = McHenry, 1958
H.l 5
-------
H.4.0 Kd Data Set for Pure Mineral Phases and Soils
Table H.8 lists the available Kd values identified for experiments conducted with pure mineral
phases as well as soils. The Kd values are listed with ancillary parameters that included clay
content, pH, CEC, surface area, solution calcium concentrations, and solution strontium
concentrations.
Table H.8. Strontium Kd data set for pure mineral phases and soils.
SrK,
(ml/g)
Clay
Content
(%)
PH
CEC
(meq/
100 g)
Surface
Area
(m2/g)
[Ca]
(ppm)
[Sr]
Background
Solution
Soil ID
Reference1
and Comments
21
0.8
5.2
0.9
1.4
0
*
NaCIO,
Soil A
1, Ohnuki, 1994
19
0.8
5.6
0.9
1.4
0
*
NaCIO,
Soil A
1, * = 4.4x102 Bq/ml 85-
Sr in 2.4xlO"8M SrCl2
22
0.8
6.2
0.9
1.4
0
*
NaCIO,
Soil A
1, * = 4.4x102 Bq/ml 85-
Sr in 2.4xlO"8M SrCl2
26
0.8
6.45
0.9
1.4
0
*
NaCIO,
Soil A
1, * = 4.4x102 Bq/ml 85-
Sr in 2.4xlO"8M SrCl2
24
0.8
6.6
0.9
1.4
0
*
NaCIO,
Soil A
1, * = 4.4x102 Bq/ml 85-
Sr in 2.4xlO"8M SrCl2
30
0.8
8.4
0.9
1.4
0
*
NaCIO,
Soil A
1, * = 4.4x102 Bq/ml 85-
Sr in 2.4xlO"8M SrCl2
43
0.8
9.2
0.9
1.4
0
*
NaCIO,
Soil A
1, * = 4.4x102 Bq/ml 85-
Sr in 2.4xlO"8M SrCl2
0
5.5
*
Quartz
1, * =4.4xl02 Bq/ml 85-
Sr in 2.4xlO"8M SrCl2
290
5.5
3.3
26.4
0
*
Kaolinite
1, * = 4.4x102 Bq/ml 85-
Sr in 2.4xlO"8M SrCl2
140
5.5
3.6
43.9
0
*
Halloysite
1, * = 4.4xl02 Bq/ml 85-
Sr in 2.4xlO"8M SrCl2
17
5.5
0.6
1.4
0
*
Chlorite
1, * = 4.4xl02 Bq/ml 85-
Sr in 2.4xlO"8M SrCl2
37
5.5
1.9
2.2
0
*
Sericite
1, * =4.4xl02 Bq/ml 85-
Sr in 2.4xlO"8M SrCl2
8
5.5
0.5
0.7
0
*
Oligoclase
1, * = 4.4xl02 Bq/ml 85-
Sr in 2.4xlO"8M SrCl2
6
5.5
0.5
0
*
Hornblend
1, * = 4.4xl02 Bq/ml 85-
Sr in 2.4xlO"8M SrCl2
H.16
-------
SrK,
(ml/g)
Clay
Content
(%)
PH
CEC
(meq/
100 g)
Surface
Area
(m2/g)
[Ca]
(ppm)
[Sr]
Background
Solution
Soil ID
Reference1
and Comments
16
5.5
0.7
0
*
Pyroxene
1, * = 4.4xl02 Bq/ml 85-
Sr in 2.4xlO"8M SrCl2
110
5.5
8.5
19.3
0
*
Mn02
1, * = 4.4xl02 Bq/ml 85-
Sr in 2.4xlO"8M SrCl2
7.7
5.8
24
113 |_iCi/l
Groundwater
AA 45/1
2 Jackson and Inch, 1989
9.9
6.1
25
105 |_iCi/l
Groundwater
AA45/3
2, Kd = -,38Ca + 0.82. r2
= 0.19
12.6
6.1
23
105 |_iCi/l
Groundwater
AA45/4
2, Ca not important to Sr
Kd
13.7
5.8
22
123 |_iCi/l
Groundwater
AA45/5
2
10.1
6
24
99 |_iCi/l
Groundwater
AA45/7
2
15.8
5.8
21
143 |_iCi/l
Groundwater
AA38/1
2
13.8
5.8
27
113 |_iCi/l
Groundwater
AA38/2
2
11
5.9
21
114 |_iCi/l
Groundwater
AA38/3
2
14.2
5.6
21
124 |_iCi/l
Groundwater
AA38/4
2
6
5.8
24
115 |aCi/l
Groundwater
AA38/5
2
7.5
5.9
21
117 |_iCi/l
Groundwater
AA38/6
2
6.9
5.9
17
108 |aCi/l
Groundwater
AA38/8
2
8.3
6.1
24
68 |aCi/l
Groundwater
AA27/1
2
8
6.2
21
71 |_iCi/l
Groundwater
AA27/2
2
6.7
6.2
28
72 |aCi/l
Groundwater
AA27/3
2
6.8
6.2
84 |aCi/l
Groundwater
AA27/4
2
4.9
6.2
18
84 |aCi/l
Groundwater
AA27/5
2
5.1
6.2
19
87 |aCi/l
Groundwater
AA27/6
2
8.5
6.2
17
88 |_iCi/l
Groundwater
AA27/7
2
8.8
6.2
18
90 |aCi/l
Groundwater
AA27/8
2
5.6
6.3
20
77 |aCi/l
Groundwater
AA34/1
2
5.3
6.4
16
79 |aCi/l
Groundwater
AA34/2
2
7.2
6.4
18
65 |aCi/l
Groundwater
AA34/3
2
5.1
6.3
18
72 |_iCi/l
Groundwater
AA34/4
2
6.5
6.4
17
75 |aCi/l
Groundwater
AA34/5
2
6
6.2
14
79 |aCi/l
Groundwater
AA34/6
2
H.17
-------
SrK,
(ml/g)
Clay
Content
(%)
PH
CEC
(meq/
100 g)
Surface
Area
(m2/g)
[Ca]
(ppm)
[Sr]
Background
Solution
Soil ID
Reference1
and Comments
6.5
6.2
15
107 |_iCi/l
Groundwater
AA34/7
2
7.6
6.2
17
107 |_iCi/l
Groundwater
AA34/8
2
21.4
0.47
Groundwater
3 Patterson and Spoel,
1981
25
0.83
Groundwater
3, CEC was approximated
by adding exch. Ca,Mg,K
12.7
0.39
Groundwater
3, Groundwater =7.4 ppm
Ca, 1.7 ppm Mg, 2.2 ppm
Na, 5.6 ppm CI, 18 ppm
so4
7.9
0.46
Groundwater
3
15.6
0.81
Groundwater
3
9.4
0.21
Groundwater
3
7.6
0.25
Groundwater
3
6.4
0.24
Groundwater
3
7.7
0.26
Groundwater
3
28.1
0.76
Groundwater
3
7.63
0.26
Groundwater
3
11.4
0.41
Groundwater
3
20.1
0.44
Groundwater
3
13
0.25
Groundwater
3
9.8
0.29
Groundwater
3
11
0.22
Groundwater
3
13
0.39
Groundwater
3
7.8
0.2
Groundwater
3
3.8
0.1
Groundwater
3
3
0.1
Groundwater
3
2.5
0.13
Groundwater
3
4
10
4
5.5
0
lxlO"8M
.OlMNaCl
Puye
soil-Na
4
15
10
5
5.5
0
lxlO"8M
.OlMNaCl
4, Noncalcareous soils
21
10
6
5.5
0
lxlO"8M
.OlMNaCl
4
24
10
7.4
5.5
0
lxlO"8M
.OlMNaCl
4
H.l 8
-------
SrK,
(ml/g)
Clay
Content
(%)
PH
CEC
(meq/
100 g)
Surface
Area
(m2/g)
[Ca]
(ppm)
[Sr]
Background
Solution
Soil ID
Reference1
and Comments
3
10
3.6
5.5
400
lxlO"8M
,01MCaCl2
Puye
soil-Ca
4
4.5
10
5.2
5.5
400
lxlO"8M
,01MCaCl2
4
5.2
10
6.8
5.5
400
lxlO"8M
,01MCaCl2
4
5.7
10
7.9
5.5
400
lxlO"8M
,01MCaCl2
4
7.2
3
0
0.1 ppm
2,000 ppm
Na
Hanford Soil
5
12.7
5
0
0.1 ppm
2,000 ppm
Na
Hanford Soil
5
14.9
7
0
0.1 ppm
2,000 ppm
Na
Hanford Soil
5
12.9
9
0
0.1 ppm
2,000 ppm
Na
Hanford Soil
5
25.1
11
0
0.1 ppm
2,000 ppm
Na
Hanford Soil
5
40.6
0.98
C-27
6
48.6
0.96
C-27
6
35
0.88
C-97
6
39.2
0.8
C-55
6
25.2
0.73
C-81
6
16.4
0.39
C-62
6
10.3
0.36
C-71
6
8.2
0.32
C-85
6
7.6
0.25
C-77
6
7.8
0.51
MK-4
6
11.2
0.38
TK3
6
10.5
0.34
RK2
6
3.7
0.34
NK2
6
3.5
5.2
2
0
lxlO"10M
NaOH/HCl
Hanford soil
7
4.6
5.6
2
0
lxlO"10M
NaOH/HCl
Hanford soil
7
5.8
5.8
2
0
lxlO"10M
NaOH/HCl
Hanford soil
7
6.1
5.9
2
0
lxlO"10M
NaOH/HCl
Hanford soil
7
8.3
6
2
0
lxlO"10M
NaOH/HCl
Hanford soil
7
H.19
-------
SrK,
(ml/g)
Clay
Content
(%)
PH
CEC
(meq/
100 g)
Surface
Area
(m2/g)
[Ca]
(ppm)
[Sr]
Background
Solution
Soil ID
Reference1
and Comments
17
7.4
2
0
lxlO"10M
NaOH/HCl
Hanford soil
7
21
7.6
2
0
lxlO"10M
NaOH/HCl
Hanford soil
7
27
7.8
2
0
lxlO"10M
NaOH/HCl
Hanford soil
7
47
8.4
2
0
lxlO"10M
NaOH/HCl
Hanford soil
7
81
9.1
2
0
lxlO"10M
NaOH/HCl
Hanford soil
7
140
70
2.4
70
0
lxlO"8M
Water
Bentonite
8
160
70
2.4
70
lxlO"8M
Groundwater
Bentonite
8
1500
70
9.3
70
0
lxlO"8M
Water
Bentonite
8
1100
70
9.3
70
lxlO"8M
Groundwater
Bentonite
8
1800
10
6.1
130
0
lxlO"8M
Water
Takadate Loam
8, hydrohalloysite=10%,
70% silt
950
10
8
130
lxlO"8M
Groundwater
Takadate Loam
8, hydrohalloysite=10%,
70% silt
550
10
6.5
60
0
lxlO"8M
Water
Hachinohe
Loam
8, hydrohalloysite = 10%,
90%) silt
260
10
8.2
60
lxlO"8M
Groundwater
Hachinohe
Loam
8, hydrohalloysite = 10%,
90%o silt
19.1
4
7.66
10.4
129
100 |_iCi/l
Hanford
Groundwater
cgs-1
9
21.5
6
7.87
5.9
58.5
100 nCi/1
Hanford
Groundwater
trench-8
9, Groundwater pH = 8.3
23.2
5
8.17
4.57
35.1
100 nCi/1
Hanford
Groundwater
tbs-1
9
48.5
0
8.24
3
3.8xlO"8M
Yucca
Groundwater
YM-22
10, Los Alamos, New
Mexico
10200
0
8.17
54
3.8xlO"8M
Yucca
Groundwater
YM-38
10, Yucca Mt tuff
sediments
2500
0
8.13
21
3.8xlO"8M
Yucca
Groundwater
YM48
10, Approximate initial
pH, final pH are
presented
3790
0
8.24
27
3.8xl0"8 M
Yucca
Groundwater
YM-49
10, Final pH 8.1-8.5
3820
0
8.24
27
3.8xlO"8M
Yucca
Groundwater
YM-50
10, Sediments = 106-500
l_im fractions
27000
0
8.4
31
10
3.8xlO"8M
Yucca
Groundwater
JA-18
10
H.20
-------
SrK,
(ml/g)
Clay
Content
(%)
PH
CEC
(meq/
100 g)
Surface
Area
(m2/g)
[Ca]
(ppm)
[Sr]
Background
Solution
Soil ID
Reference1
and Comments
4850
0
8.63
31
50
3.8xlO"8M
Yucca
Groundwater
JA-19
10
85
0
8.25
8
10
3.8xlO"8M
Yucca
Groundwater
JA-32
10
17.7
0
8.5
8
50
3.8xlO"8M
Yucca
Groundwater
JA-33
10
385
0
8.39
105
10
3.8xlO"8M
Yucca
Groundwater
JA-37
10
149
0
8.45
105
50
3.8xlO"8M
Yucca
Groundwater
JA-38
10
25000
12
10 nCi/ml
kaolinite
13
530
12
10 nCi/ml
chlorite
13
71,000
12
10 nCi/ml
FeOOH
13
1.6
0.5
6.2
0.05
10xl0"6M
Groundwater
Sediments
14
2.6
3
6.2
0.3
10xl0"6M
Groundwater
Sediments
14, Added Kaolinite to
sand
3.4
5
6.2
0.5
10xl0"6M
Groundwater
Sediments
14, CEC estimated based
on kaolinite =10
meq/100 g
4.6
8
6.2
0.8
10xl0"6M
Groundwater
Sediments
14
6.7
13
6.2
1.3
10xl0"6M
Groundwater
Sediments
14
17,000
97
lxlO"10M
Ohya tuff
14, Akiba and
Hashimoto, 1990
150
3.4
lxlO"10M
Pyrophyllite
14, log Kd = log CEC +
constant: for trace [Sr]
780
2.4
lxlO"10M
Sandstone
14, pH not held constant,
ranged from 6 to 9.
95
1.9
lxlO"10M
Shale
14, lg solid: 50ml
sol'n,centrifuged,32-
60mesh
440
1.9
lxlO"10M
Augite
Andesite
14, CEC of Cs and Kd of
Sr
39
1.2
lxlO"10M
Plagiorhyolite
14
380
0.75
lxlO"10M
Olivine Basalt
14
50
0.57
lxlO"10M
Vitric Massive
Tuff
14
82
0.54
lxlO"10M
Inada granite
14
H.21
-------
SrK,
(ml/g)
Clay
Content
(%)
PH
CEC
(meq/
100 g)
Surface
Area
(m2/g)
[Ca]
(ppm)
[Sr]
Background
Solution
Soil ID
Reference1
and Comments
22
0.35
lxlO"10M
Rokko Granite
14
1.3
0.033
lxlO"10M
Limestone
14
2,000
2
lxlO"10M
Muscovite
14
140
0.93
lxlO"10M
Chlorite
14
40
0.36
lxlO"10M
Hedenbergite
14
20
0.33
lxlO"10M
Hornblende
14
71
0.11
lxlO"10M
Grossular
14
150
0.07
lxlO"10M
Microcline
14
0.92
0.067
lxlO"10M
Forsterite
14
14
0.034
lxlO"10M
K-Feldspar
14
30
0.032
lxlO"10M
Albite
14
3
0.022
lxlO"10M
Epidote
14
23
0.0098
lxlO"10M
Quartz
14
400
42.4
7.2
34
0
Water
Ringhold Soil
11, Soil from Richland
WA
135
26.9
8.3
13.6
0
Water
Bowdoin Soil
11, from Montana
600
33.5
6.5
26.3
0
Water
Hall Soil
11, from Nebraska
70
3.5
8.3
5.8
0
Water
Composite Soil
11, from Hanford Site
2.4
4
Groundwater
Eolian Sand
12
4.7
5
Eolian Sand
12, Belgian soils
6
7
Eolian Sand
12, Composition of
Groundwater was not
given
2.3
4
Mol White
Sand
12, Compared static vs.
dynamic Kd
5.5
5
Mol White
Sand
12
4.8
7
Mol White
Sand
12
H.22
-------
SrK,
(ml/g)
Clay
Content
(%)
PH
CEC
(meq/
100 g)
Surface
Area
(m2/g)
[Ca]
(ppm)
[Sr]
Background
Solution
Soil ID
Reference1
and Comments
2.6
4
Mol Lignitic
Sand
12
5.3
5
Mol Lignitic
Sand
12
7.2
7
Mol Lignitic
Sand
12
I References: 1 = Ohnuki, 1994; 2 = Jackson and Inch ,1989; 3 =Patterson and Spoel ,1981; 4 = Keren and O'Connor, 1983; 5 Nelson,
1959; 6 = Inch and Killey, 1987; 7 = Rhodes and Nelson, 1957; 8 = Konishi etal., 1988; 9 = Serne etal., 1993; 10 = Vine etal., 1980;
II =McHenry, 1958;12 = Baetsle etal., 1964; 13 = Ohnuki, 1991; 14 = Lieser and Steinkopff, 1989
H.23
-------
H.5.0 References
Adeleye, S. A., P. G. Clay, and M. O. A. Oladipo. 1994. "Sorption of Caesium, Strontium and
Europium Ions on Clay Minerals." Journal of Materials Science, 29:954-958.
Akiba, D., and H. Hashimoto. 1990. "Distribution Coefficient of Strontium on Variety of
Minerals and Rocks." Journal of Nuclear Science and Technology, 21:215-219.
Ames, L., and D. Rai. 1978. Radionuclide Interactions with Soil and Rock Media. Volume 1:
Processes Influencing Radionuclide Mobility and Retention, Element Chemistry and
Geochemistry, Conclusions and Evaluation. PB-292 460, Pacific Northwest National
Laboratory, Richland, Washington.
Baetsle, L. H., P. Dejonghe, W. Maes, E. S. Simpson, J. Souffriau, and P. Staner. 1964.
Underground Radionuclide Movement. EURAEC-703, European Atomic Energy
Commission, Vienna, Austria.
Cantrell, K., P. F. Martin, and J. E. Szecsody. 1994. "Clinoptilolite as an In-Situ Permeable
Barrier to Strontium Migration in Ground Water." In In-Situ Remediation: Scientific Basis
for Current and Future Technologies. Part 2., G. W. Gee and N. Richard Wing (eds.).
pp. 839-850. Battelle Press, Columbus, Ohio.
Cui, D., and R. E. Eriksen. 1995. "Reversibility of Strontium Sorption on Fracture Fillings." In
Scientific Basis for Nuclear Waste Management XVIII, T. Murakami and R. C. Ewing (eds.),
pp. 1045-1052. Material Research Society Symposium Proceedings, Volume 353, Materials
Research Society, Pittsburgh, Pennsylvania.
Del Debbio, J. A. 1991. "Sorption of Strontium, Selenium, Cadmium, and Mercury in Soil."
RadiochimicaActa, 52/53:181-186.
Faure, G., and J. L. Powell. 1972. Strontium Isotope Geology. Springer-Verlag, Berlin,
Germany.
Inch, K. J., and R. W. D. Killey. 1987. "Surface Area and Radionuclide Sorption in
Contaminated Aquifers." Water Pollution Research Journal of Canada, 22:85-98.
Jackson, R. E., and K. J. Inch. 1989. "The In-Situ Adsorption of 90Sr in a Sand Aquifer at the
Chalk River Nuclear Laboratories." Journal of Contaminant Hydrology, 4:27-50.
Keren, R., and G. A. O'Connor. 1983. "Strontium Adsorption by Noncalcareous Soils -
Exchangeable Ions and Solution Composition Effects." Soil Science, 135:308-315.
H.24
-------
Konishi, M., K. Yamamoto, T. Yanagi, and Y. Okajima. 1988. "Sorption Behavior of Cesium,
Strontium and Americium Ions on Clay Materials." Journal of Nuclear Science and
Technology, 25:929-933.
Lefevre, R., M. Sardin, and D. Schweich. 1993. "Migration of Strontium in Clayey and
Calcareous Sandy Soil: Precipitation and Ion Exchange." Journal of Contaminant
Hydrology, 13:215-229.
Lieser, K. H., B. Gleitsmann, and Th. Steinkopff 1986. "Sorption of Trace Elements or
Radionuclides in Natural Systems Containing Groundwater and Sediments." Radiochimica
Acta, 40:33-37.
Lieser, K. H., and Th. Steinkopff. 1989. "Sorption Equilibria of Radionuclides or Trace
Elements in Multicomponent Systems." Radiochimica Acta, 47:55-61.
McHenry, J. R. 1958. "Ion Exchange Properties of Strontium in a Calcareous Soil " Soil
Science Society of America, Proceedings, 22:514-518.
Nelson, J. L. 1959. Recent Studies atHanford on Soil and Mineral Reactions in Waste
Disposal. HW-SA-2273, Westinghouse Hanford Company, Richland, Washington.
Hem, J. D. 1985. Study and Interpretation of the Chemical Characteristics of Natural Water.
Water Supply Paper 2254. Distribution Branch, Text Products Section, U.S. Geological
Survey, Alexandria, Virginia.
Neter, J. and W. Wasserman. 1974. Applied Linear Statistical Models. Richard D. Irwin, Inc.,
Homewood, Illinois.
Ohnuki, T. 1991. "Characteristics of Migration of 85Sr and 137Cs in Alkaline Solution Through
Sandy Soil." Material Research Society Proceedings, 212:609-616.
Ohnuki, T. 1994. "Sorption Characteristics of Strontium on Sandy Soils and Their
Components." Radiochimica Acta, 64:237-245.
Patterson, R. J., and T. Spoel. 1981. "Laboratory Measurements of the Strontium Distribution
Coefficient for Sediments From a Shallow Sand Aquifer." Water Resources Research,
17:513-520.
Petersen, L. W., P. Moldrup, O. H. Jacobsen, and D. E. Rolston. 1996. "Relations Between
Specific Surface Area and Soils Physical and Chemical Properties." Soil Science, 161:9-21.
Rhodes, D. W., and J. L. Nelson. 1957. Disposal of Radioactive Liquid Wastes From the
Uranium Recovery Plant. HW-54721, Westinghouse Hanford Company, Richland,
Washington.
H.25
-------
Satmark, B., and Y. Albinsson. 1991. "Sorption of Fission Products on Colloids Made of
Naturally Occurring Minerals and the Stability of these Colloids." Radiochimica Acta,
58/59:155-161.
Serne, R. J., J. L. Conca, V. L. LeGore, K. J. Cantrell, C. W. Lindenmeier, J. A. Campbell, J. E.
Amonette, and M. I. Wood. 1993. Solid-Waste Leach Characteristics and Contaminant-
Sediment Interactions. Volume 1: Batch Leach and Adsorption Tests and Sediment
Characterization. PNL-8889, Pacific Northwest National Laboratory, Richland,
Washington.
Serne, R. J., and V. L. LeGore. Strontium-90 Adsorption-Desorption Properties and Sediment
Characterization at the 100 N-Area. PNL-10899, Pacific Northwest National Laboratory,
Richland, Washington.
Sposito, G. 1984. The Surface Chemistry of Soils. Oxford University Press, New York,
New York.
Strenge, D. L., and S. R. Peterson. 1989. Chemical Databases for the Multimedia
Environmental Pollutant Assessment System. PNL-7145, Pacific Northwest National
Laboratory, Richland, Washington.
Vine, E. N., R. D. Aguilar, B. P. Bayhurst, W. R. Daniels, S. J. DeVilliers, B. R. Erdal, F. O.
Lawrence, S. Maestas, P. Q. Oliver, J. L. Thompson, and K. Wolfsberg. 1980. Sorption-
Desorption Studies on Tuff. II. A Continuation of Studies with Samples form Jackass Flats,
Nevada and Initial Studies with Samples form Yucca Mountain, Nevada. LA-8110-MS, Los
Alamos Scientific Laboratory, Los Alamos, New Mexico.
H.26
-------
APPENDIX I
Partition Coefficients For Thorium
-------
Appendix I
Partition Coefficients For Thorium
1.1.0 BACKGROUND
Two generalized, simplifying assumptions were established for the selection of thorium Kd
values for the look-up table. These assumptions were based on the findings of the literature
review conducted on the geochemical processes affecting thorium sorption. The assumptions are
as follows:
• Thorium adsorption occurs at concentrations less than 10"9M. The extent of thorium
adsorption can be estimated by soil pH.
• Thorium precipitates at concentrations greater than 10"9 M. This concentration is based
on the solubility of Th(OH)4 at pH 5.5. Although (co)precipitation is usually quantified
with the solubility construct, a very large Kd value will be used in the look-up table to
approximate thorium behavior in systems with high thorium concentrations.
These assumptions appear to be reasonable for a wide range of environmental conditions.
However, these simplifying assumptions are clearly compromised in systems containing high
alkalinity (LaFlamme and Murray, 1987), carbonate (LaFlamme and Murray, 1987), or sulfate
(Hunter et al., 1988) concentrations, and low or high pH values (pH values less than 3 or greater
than 8) (Hunter et al., 1988; LaFlamme and Murray, 1987; Landa et al., 1995). These
assumptions will be discussed in more detail in the following sections.
Thorium Kd values and some important ancillary parameters that influence sorption were
collected from the literature and tabulated. Data included in this table were from studies that
reported Kd values (not percent adsorbed or Freundlich or Langmuir constants) and were
conducted in systems consisting of:
• Low ionic strength (< 0.1 M)
• pH values between 4 and 10.5
• Dissolved thorium concentrations less than 10"9 M
• Low humic material concentrations (<5 mg/1)
• No organic chelates (such as EDTA)
These aqueous chemistry constraints were selected to limit the thorium Kd values evaluated to
those that would be expected to exist in a far-field. The ancillary parameters included in these
tables were clay content, calcite concentration, pH, and CEC. Attempts were also made to
include the concentrations of organic matter and aluminum/iron oxides in the solid phase in the
data set. However, these latter ancillary parameters were rarely included in the reports
evaluated during the compilation of the data set. The data set included 17 thorium Kd values.
1.2
-------
The descriptive statistics of the thorium Kd data set are presented in Table 1.1. The lowest
thorium Kd value was 100 ml/g for a measurement made on a pH 10 soil (Rancon, 1973). The
largest thorium Kd value was 500,000 ml/g for a measurement made on a silt/quartz soil of schist
origin (Rancon, 1973). The average thorium Kd value for the 17 observations was 54,000 ±
29,944 ml/g.
Table LI. Descriptive statistics of thorium Kd value data set presented in Section 1.3.
Thorium K,,
(ml/g)
Clay
Content
(wt.%)
pH
CEC
(meq/100 g)
Calcite
(wt.%)
Al/Fe-
Oxides
(wt.%)
Organic
Matter
(wt.%)
Mean
54,000
26.8
6.1
13.7
29
—
—
Standard Error
29,944
6.3
0.4
11.2
13.4
--
--
Median
5,000
30
6
2.9
25
--
--
Mode
100,000
40
6
2.9
0
--
--
Standard Deviation
123,465
14.1
1.5
29.8
30.1
--
--
Sample Variance
1.5xl010
199.2
2.1
886.2
905
--
--
Minimum
100
12
4
1.7
0
--
--
Maximum
500,000
40
10
81.2
60
--
--
No. Observations
17
5
17
7
5
0
0
1.2.0 Approach and Regression Models
1.2.1 Correlations with Thorium Kd Values
A matrix of the correlation coefficients for thorium Kd values with soil parameters is
presented in Table 1.2. The correlation coefficients that are significant at or less than the
1 percent or 5 percent level of probability are identified. The parameter with the largest
correlation coefficient with thorium Kd was pH (r = 0.58, n = 16, P < 0.01, where r, n, and P
represent correlation coefficient, number of observations, and level of probability, respectively).
The pH range for this data set is 4 to 7.6. When Kd data for pH 10 is included in the regression
analysis, the correlation coefficient decreases to 0.14 (n = 17, P < 0.22). The nonsignificant
correlations with clay content, CEC, and calcite may in part be attributed to the small number of
values in the data sets.
1.3
-------
Table 1.2. Correlation coefficients (r) of the thorium Kd value data set presented in
Section 1.3.
Thorium Kd
Clay Content
pH
CEC
Thorium Kd
1
Clay Content
-0.79
1
pH
0.58 2
(0.14)3
-0.84 1
1
CEC
-0.15
--
-0.21
1
Calcite
0.76
-0.998 2
o
00
—
12 Correlation coefficient is significant at the 5 percent (P < 0.05) (indicated by footnote a) or 1 percent (P < 0.01)
(indicated by footnote b) level of significance, respectively. Significance level is in part dependent on the number
of observations, n, (more specifically, the degrees of freedom) and variance of each correlation comparison
(Table 1.1). Thus, it is possible for thorium Kd/clay correlation coefficient of -0.79 to be not significant and the
thorium Kd /pH correlation coefficient of 0.58 to be significant because the former has 4 degrees of freedom and
the latter has 15 degrees of freedom.
3 Excluding the Kd values at the highest pH value (pH 10), the correlation is 0.58 (n = 16). Including this Kd
value, the correlation coefficient decreases to 0.14.
1.2.2 Thorium Kd Values as a Function of pH
Thorium Kd values were significantly correlated to pH between the pH range of 4 to 8, but were
not correlated to pH between the range 4 to 10 (Figure 1.1 and Table 1.2). The pH dependence of
thorium sorption to solid phases has been previously demonstrated with pure mineral phases
(Hunter et al., 1987; LaFlamme and Murray, 1987). The pH dependence can be explained in
part by taking into consideration the aqueous speciation of thorium in groundwater. Thorium
aqueous speciation changes greatly as a function of groundwater pH (Table 1.3). As the pH
increases, the thorium complexes become more anionic or neutral, thereby becoming less prone
to be electrostatically attracted to a negatively charged solid phase. This decrease in electrostatic
attraction would likely result in a decrease in Kd values. Figure 1.1 shows an increase in thorium
Kd values between pH 4 and 8. This may be the result of the pH increasing the number of
exchange sites in the soil. At pH 10, the large number of neutral or anionic thorium complexes
may have reduced the propensity of thorium to sorb to the soil.
1.4
-------
7
6
5
A
H 4
"ao
° 3
2
1
3456789 10 11
pH
Figure LI. Linear regression between thorium Kd values
and pH for the pH range from 4 to 8. [The
single Kd value at pH 10 is identified by the
filled circle.]
Table 1.3. Calculated aqueous speciation of thorium as a function of pH. [The
composition of the water and details of the aqueous speciation calculations are
presented in Chapter 5. Total thorium concentration used in the aqueous
speciation calculations is 1 ng/ml.]
pH
Dominant
Percent (%) of
Aqueous Species
Total Dissolved Thorium
3
ThFf
54
ThFj
42
7
Th(HP04)r
98
9
Th(OH)° (aq)
99
1.5
-------
The regression equation between the pH range of 4 to 8 that is shown in Figure 1.1 is
log (ThKd) =-0.13 + 0.69(pH). (1.1)
The statistics for this equation are presented in Table 1.4. The fact that the P-value for the
intercept coefficient is >0.05 indicates that the intercept is not significantly (P > 0.05) different
than 0. The fact that the P-value for the slope coefficient is <0.05 indicates that the slope is
significantly (P > 0.05) different than 1. The lower and upper 95 percent coefficients presented
in Table 1.4 reflect the 95 percent confidence limits of the coefficients. They were used to
calculate the upper and lower limits of expected thorium Kd values at a given pH value.
1.2.3 Approach
Linear regression analyses were conducted with data collected from the literature. These
analyses were used as guidance for selecting appropriate Kd values for the look-up table. The Kd
values used in the look-up tables could not be based entirely on statistical consideration because
the statistical analysis results were occasionally nonsensible. For example, the data showed a
negative correlation between clay content and thorium Kd values. This trend contradicts well
established principles of surface chemistry. Instead, the statistical analysis was used to provide
guidance as to the approximate range of values to use and to identify meaningful trends between
the thorium Kd values and the solid phase parameters. Thus, the Kd values included in the
look-up table were in part selected based on professional judgment. Again, only low-ionic
strength solutions similar to that expected in far-field ground waters were considered in these
analyses.
Table 1.4. Regression coefficient and their statistics relating thorium Kd values and pH.
[log (Th Kd) = -0.13 + 0.69(pH), based on data presented in Figure I.I.]
Coefficients
Standard
Error
t- Statistic
P-value
Lower
95%
Upper
95%
Intercept Coefficient
2.22
1.06
0.47
0.64
-1.77
2.76
Slope Coefficient
0.57
0.18
3.24
0.006
0.19
0.95
1.6
-------
The look-up table (Table 1.5) for thorium Kd values was based on thorium concentrations and
pH. These 2 parameters have an interrelated effect on thorium Kd values. The maximum
concentration of dissolved thorium may be controlled by the solubility of hydrous thorium
oxides (Felmy et al., 1991; Rai et al., 1995; Ryan andRai, 1987). The dissolution of hydrous
thorium oxides may in turn vary with pH. Ryan and Rai (1987) reported that the solubility of
hydrous thorium oxide is ~10"8 5 to ~10"9 in the pH range of 5 to 10. The concentration of
dissolved thorium increases to ~10"26 M (600 mg/L) as pH decreases from 5 to 3.2. Thus,
2 categories, pH 3 - 5 and pH 5 - 10, based on thorium solubility were included in the look-up
table. Although precipitation is typically quantified by the solubility construct, a very large Kd
value was used in Table 1.5 to describe high thorium concentrations.
The following steps were taken to assign values to each category in the look-up table. For Kd
values in systems with pH values less than 8 and thorium concentrations less than the estimated
solubility limits, Equation 1.1 was used. This regression equation is for data collected between
the pH range of 4 to 8 as shown in Figure 1.1 [log (Th Kd) = -0.13 + 0.69(pH)]. pH values of 4
and 6.5 were used to estimate the "pH 3 to 5" and "pH 5 to 8" categories, respectively. The Kd
values in the "pH 8 to 10" category were based on the single laboratory experiment conducted at
pH 10 that had a Kd of 200 ml/g. Upper and lower estimates of thorium Kd values were
calculated by adding or subtracting 1 logarithmic unit to the "central estimates" calculated above
for each pH category (Figure 1.2). The 1 logarithm unit estimates for the upper and lower limits
are based on visual examination of the data in Figure 1.1. The use of the upper and lower
regression coefficient values at the 95 percent confidence limits (Table 1.5) resulted in calculated
ranges that were unrealistically large. At pH 4, for the "pH 3 to 5" category, the lower and upper
log (Th Kd) values were calculated to be 1 and 6.6, respectively; at pH 6.5, this range of Kd was -
0.5 to 9.0). All thorium Kd values for systems containing concentrations of dissolved thorium
greater than their estimated solubility limit (10"9 M for pH 5 to 10 and 10"2 6 M for pH < 5) were
assigned a Kd of 300,000 ml/g.
Table 1.5. Look-up table for thorium Kd values (ml/g) based on pH and dissolved
thorium concentrations. [Tabulated values pertain to systems consisting of
low ionic strength (<0.1 M), low humic material concentrations (<5 mg/1), no
organic chelates (such as EDTA), and oxidizing conditions.]
(ml/g)
pH
3-5
5-8
8-10
Dissolved Th (M)
Dissolved Th (M)
Dissolved Th (M)
<102 6
>1026
<109
>109
<109
>109
Minimum
62
300,000
1,700
300,000
20
300,000
Maximum
6,200
300,000
170,000
300,000
2,000
300,000
1.7
-------
Figure 1.2. Linear regression between thorium Kd values
and pH for the pH Range 4 to 8. [Values ±1
logarithmic unit from the regression line are
also identified. The single Kd value at pH 10
is identified by the filled circle)].
1.3.0 Kd Data Set for Soils
The data set of thorium Kd values used to develop the look-up table are listed in Table 1.6.
1.8
-------
Table 1.6. Data set containing thorium Kd values.
Thorium
K,
(ml/g)
PH
Clay
(wt.%)
CEC1
(meq/
lOOg)
OM1
(wt. %)
Fe-
Oxides
(wt.%)
Th
CM)
Calcite
(wt. %)
Solution
Chemistry
Soil ID and
Characteristics
Ref2
10,0000
7.6
3
Synthetic GW1,
pH 6.6
Soil A
1
500,000
6
40
0
Syn. GW, 232Th
Competing Ion
Silt+Qtz Sed., Schist soil
2
1,000
4
40
0
Syn. GW, 232Th
Competing Ion
Silt+Qtz Sed., Schist soil
2
100,000
8
12
60
Syn. GW, 232Th
Competing Ion
Silt+Qtz+OM+calcite,
Schist Soil
2
150,000
7
30
25
Syn. GW, 232Th
Competing Ion
Cadarache Sed.
2
100
10
12
60
Syn. GW, 232Th
Competing Ion
Silt+Qtz+OM+calcite,
Schist Soil
2
24,000
6
Groundwater
Glacial till, Clay
3
5,800
6
Groundwater
Fine Coarse Sand
3
1,028.6
5.1
2.9
Gleyed Dystric Brunisol, Ae
Horizon 4-15 cm
4
1,271
5.2
2.1
Gleyed Dystric Brunisol, Bf
Horizon 1 5-45 cm
4
5,000
4.5
Jefferson City, Wyoming,
Fine Sandstone and Silty
Clay
5
10,000
5.8
Jefferson City, Wyoming,
Fine Sandstone and Silty
Clay
5
15,000
7
Jefferson City, Wyoming,
Fine Sandstone and Silty
Clay
5
1,578
5.2
81.2
Groundwater
Gleyed Dystric Brunisol, Ah
Horizon
6
1,862.5
5.1
2.9
Groundwater
Gleyed Dystric Brunisol, Ae
Horizon
6
1,153.7
5.2
2.1
Groundwater
Gleyed Dystric Brunisol, Bf
Horizon
6
206.9
6.2
1.7
Groundwater
Gleyed Dystric Brunisol, C
Horizon
6
1 CEC = cation exchange capacity, OC = organic matter, GW = groundwater.
2 References: 1 =Legoux etal., 1992; 2 =Rancon, 1973; 3 = Bell and Bates, 1988; 4= Sheppard etal., 1987; 5 = Haji-Djafari et al.,
1981; 6 = Thibault et al., 1990.
1.9
-------
1.5.0 References
Ames, L. L., and D. Rai. 1978. Radionuclide Interactions with Soil and Rock Media.
Volume 1: Processes Influencing Radionuclide Mobility and Retention, element Chemistry
and Geochemistry, and Conclusions and Evaluation. EPA 520/6-78-007 A, Prepared for the
U.S. Environmental Protection Agency by the Pacific Northwest National Laboratory,
Richland, Washington.
Bell, J., and T. H. Bates. 1988. "Distribution Coefficients of Radionuclides Between Soils and
Groundwaters and Their Dependence on Various Test Parameters." The Science of the Total
Environment, 69:297-317.
Felmy, A. R., D. Rai, and D. A. Moore. 1993. "The Solubility of Hydrous Thorium(IV) Oxide
in Chloride Media: Development of an Aqueous Ion-Interaction Model." Radiochimica
Acta, 55:177-185.
Haji-Djafari, S., P. E. Antommaria, and H. L. Crouse. 1981. Attenuation of Radionuclides and
Toxic Elements by In Situ Soils at a Uranium Tailings Pond in Central Wyoming. In
Permeability and Groundwater Contaminant Transport, (eds.) T. F. Zimmie and C. O.
Riggs, pp. 221-242. American Society for Testing and Materials, Philadelphia,
Pennsylvania.
Hem, J. D. 1985. Study and Interpretation of the Chemical Characteristics of Natural Water.
U.S. Geological Survey Water Supply Paper 2254, U.S. Geological Survey, Alexandria,
Virginia. 1985
Hunter, K. A., D. J. Hawke, and L. K. Choo. 1988. "Equilibrium Adsorption of Thorium by
Metal Oxides in Marine Electrolytes." Geochimica et Cosmochimica Acta, 52:627-636.
LaFlamme, B. D., and J. W. Murray. 1987. "Solid/Solution Interaction: The Effect of
Carbonate Alkalinity on Adsorbed Thorium." Geochimica et Cosmochimica Acta,
51:243-250.
Landa, E. R., A. H. Le, R. L. Luck, and P. J. Yeich. 1995. "Sorption and Coprecipitation of
Trace Concentrations of Thorium with Various Minerals Under Conditions Simulating an
Acid Uranium Mill Effluent Environment." Inorganica Chimica Acta, 229:247-252.
Legoux, Y., G. Blain, R. Guillaumont, G. Ouzounian, L. Brillard, and M. Hussonnois. 1992.
"Kd Measurements of Activation, Fission, and Heavy Elements in Water/Solid Phase
Systems." Radiochimica Acta, 58/59:211-218.
1.10
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Rai, D., A. R. Felmy, D. A. Moore, and M. J. Mason. 1995. "The Solubility of Th(IV) and
U(IV) Hydrous Oxides in Concentrated NaHC03 and Na2C03 Solutions." In Scientific Basis
for Nuclear Waste Management XVIII, Part 2, T. Murakami andR. C. Ewing (eds.),
pp. 1143-1150, Materials Research Society Symposium Proceedings, Volume 353, Materials
Research Society, Pittsburgh, Pennsylvania.
Rancon, D. 1973. "The Behavior in Underground Environments of Uranium and Thorium
Discharged by the Nuclear Industry." In Environmental Behavior of Radionuclides Released
in the Nuclear Industry, pp. 333-346. IAEA-SM-172/55, International Atomic Energy
Agency Proceedings, Vienna, Austria.
Ryan, J. L., and D. Rai. 1987. "Thorium(IV) Hydrous Oxide Solubility." Inorganic Chemistry,
26:4140-4142.
Sheppard, M. I., D. H. Thibault, and J. H. Mitchell. 1987. "Element Leaching and Capillary
Rise in Sandy Soil Cores: Experimental Results." Journal of Environmental Quality,
16:273-284.
Thibault, D. H., M. I. Sheppard, and P. A. Smith. 1990. A Critical Compilation and Review of
Default Soil Solid/Liquid Partition Coefficients, Kd, for Use in Environmental Assessments.
AECL-10125, Whiteshell Nuclear Research Establishment, Atomic Energy of Canada
Limited, Pinawa, Canada.
1.11
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APPENDIX J
Partition Coefficients For Uranium
-------
Appendix J
Partition Coefficients For Uranium
J. 1.0 Background
The review of uranium Kd values obtained for a number of soils, crushed rock material, and
single-mineral phases (Table J.5) indicated that pH and dissolved carbonate concentrations are
the 2 most important factors influencing the adsorption behavior of U(VI). These factors and
their effects on uranium adsorption on soils are discussed below. The solution pH was also used
as the basis for generating a look-up table of the range of estimated minimum and maximum Kd
values for uranium.
Several of the studies identified in this review demonstrate the importance dissolved carbonate
through the formation of strong anionic carbonato complexes on the adsorption and solubility of
dissolved U(VI). This complexation especially affects the adsorption behavior of U(VI) at
alkaline pH conditions. Given the complexity of these reaction processes, it is recommended
that the reader consider the application of geochemical reaction codes, and surface complexation
models in particular, as the best approach to predicting the role of dissolved carbonate in the
adsorption behavior of uranium and derivation of Kd values when site-specific Kd values are not
available for U(VI).
J.2.0 Availability of Kd Values for Uranium
More than 20 references were identified that reported the results of Kd measurements for the
sorption of uranium onto soils, crushed rock material, and single mineral phases. These studies
were typically conducted to support uranium migration investigations and safety assessments
associated with the genesis of uranium ore deposits, remediation of uranium mill tailings,
agriculture practices, and the near-surface and deep geologic disposal of low-level and high-level
radioactive wastes (including spent nuclear fuel).
A large number of laboratory uranium adsorption/desorption and computer modeling studies
have been conducted in the application of surface complexation models (see Chapter 5 and
Volume I) to the adsorption of uranium to important mineral adsorbates in soils. These studies
are also noted below.
Several published compilations of Kd values for uranium and other radionuclides and inorganic
elements were also identified during the course of this review. These compilations are also
briefly described below for the sake of completeness because the reported values may have
applicability to sites of interest to the reader. Some of the Kd values in these compilations are
tabulated below, when it was not practical to obtain the original sources references.
J.2
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J. 2.1 Sources of Error and Variability
The Kd values compiled from these sources show a scatter of 3 to 4 orders of magnitude at any
pH value from pH 4 to 9. As will be explained below, a significant amount of this variation
represents real variability possible for the steady-state adsorption of uranium onto soils resulting
from adsorption to important soil mineral phases (e.g., clays, iron oxides, clays, and quartz) as a
function of important geochemical parameters (e.g., pH and dissolved carbonate concentrations).
However, as with most compilations of Kd values, those in this report and published elsewhere,
reported Kd values, and sorption information in general, incorporate diverse sources of errors
resulting from different laboratory methods (batch versus column versus in situ measurements),
soil and mineral types, length of equilibration (experiments conducted from periods of hours to
weeks), and the fact that the Kd parameter is a ratio of 2 concentrations. These sources of error
are discussed in detail in Volume I of this report.
Taking the ratio of 2 concentrations is particularly important to uranium, which, under certain
geochemical conditions, will absorb to soil at less than 5 percent (very small Kd) or up to more
than 95 percent (very large Kd) of its original dissolved concentration. The former circumstance
(<5 percent adsorption) requires the investigator to distinguish very small differences in the
analyzed initial and final concentrations of dissolved uranium. On the other hand, the latter
circumstance (>95 percent adsorption) requires analysis of dissolved uranium concentrations that
are near the analytical minimum detection limit. When comparing very small or very large Kd
values published in different sources, the reader must remember this source of uncertainty can be
the major cause for the variability.
In the following summaries, readers should note that the valence state of uranium is given as that
listed in the authors' publications. Typically, the authors describe their procedures and results in
terms of "uranium," and do not distinguish between the different valence states of uranium
[U(VI) and U(IV)] present. In most studies, it is fair for the reader to assume that the authors are
referring to U(VI) because no special precautions are described for conducting the adsorption
studies using a dissolved reductant and/or controlled environmental chamber under ultralow
oxygen concentrations. However, some measurements of uranium sorption onto crushed rock
materials may have been compromised unbeknownst to the investigators by reduction of U(VI)
initially present to U(IV) by reaction with ferrous iron [Fe(II)] exposed on fresh mineral
surfaces. Because a major decrease of dissolved uranium typically results from this reduction
due to precipitation of U(IV) hydrous-oxide solids (i.e., lower solubility), the measured Kd
values can be too large as a measure of U(VI) sorption. This scenario is possible when one
considers the geochemical processes associated with some in situ remediation technologies
currently under development. For example, Fruchter el al. (1996) [also see related paper by
Amonette et al. (1994)] describe development of a permeable redox barrier remediation
technology that introduces a reductant (sodium dithionite buffered at high pH) into contaminated
sediment to reduce Fe(III) present in the sediment minerals to Fe(II). Laboratory experiments
have shown that dissolved U(VI) will accumulate, via reduction of U(VI) to U(IV) and
subsequent precipitation as a U(IV) solid, when it contacts such treated sediments.
J.3
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J. 2.2 Uranium Kd Studies on Soils and Rock Materials
The following sources of Kd values considered in developing the uranium Kd look-up table are
listed in alphabetical order. Due to their extensive length, summary tables that list the uranium
Kd values presented or calculated from data given in these sources are located at the end of this
appendix.
Ames etal. (1982) studied the adsorption of uranium on 3 characterized basalts and associated
secondary smectite clay. The experiments were conducted at 23 and 60°C under oxidizing
conditions using 2 synthetic groundwater solutions. The compositions of the solutions were
based on those of groundwater samples taken at depth from the Columbia River basalt
formations. The basalts were crushed, and the 0.85-0.33 mm size fraction used for the
adsorption studies. The groundwater solutions were mixed with the basaltic material and
smectite in a ratio of 10 ml/1 g, and equilibrated for 60 days prior to analysis. Four initial
concentrations of uranium (l.OxlO"4, l.OxlO"5, l.OxlO"6, and l.OxlO"7 M uranium) were used for
the measurements. The pH values in the final solutions ranged from 7.65 to 8.48. Uranium Kd
values listed as "D" values in Ames et al. (1982, Table III) for the 23°C sorption measurements
are listed in Table J.5.
Bell and Bates (1988) completed laboratory uranium (and other radionuclides) Kd measurements
designed to evaluate the importance of test parameters such as pH, temperature, groundwater
composition, and contact time at site-relevant conditions. Materials used for the Kd
measurements included a sample of borehole groundwater that was mixed in a solution-to-solid
ratio of 10 ml/1 g with the <5-mm size fraction of each of 5 soil materials. For the experiments
conducted as a function of pH, the initial pH of the groundwater samples was adjusted by the
addition of HC1, NaOH, or NH4OH. The soils included a glacial till clay, sand, and 3 coarse
granular deposits (listed as CI :2, C.3, and C.6 by Bell and Bates). The Kd values were measured
using a batch method where the test vessel was agitated continuously at a fixed temperature for a
pre-determined length of time. The uranium Kd values measured for the 5 soils at pH 5.7 and
15°C sampled at 14 days are listed in Table J.5. Bell and Bates noted that steady-state
conditions were seldom achieved for 14 days contact at pH 5.7 and 15°C. For the clay and CI :2
soils, which exhibited the low-sorptive properties, the uranium Kd values doubled for each
temperature increase of 5°. No significant temperature dependence was observed in the uranium
Kd values measured using the other 3 soil materials. The uranium Kd values measured as a
function of pH showed a maximum in sorption near pH 6 and 10, for the sand and clay soils.
However, these 7-day experiments were affected by kinetic factors.
Erickson (1980) measured the Kd values for several radionuclides, including uranium, on abyssal
red clay. The dominant mineral in the clay was iron-rich smectite, with lesser amounts of
phillipsite, hydrous iron and manganese oxides. The Kd values were measured using a batch
equilibration technique with equilibration times of 2-4 days and an initial concentration of
dissolved uranium of approximately 3. lxlO"8 mg/ml. The uranium Kd values measured at pH
values of 2.8 and 7.1 by Erickson (1980) are listed in Table J.5.
J.4
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Erikson et al. (1993) determined the Kd values for the adsorption of uranium on soil samples
from the U.S. Department of Army munition performance testing sites at Aberdeen Proving
Ground, Maryland, and Yuma Proving Ground, Arizona. The soil samples included 2 silt loams
(Spesutie and Transonic) from the Aberdeen Proving Ground, and sandy loam (Yuma) from the
Yuma Proving Ground. The names of the soil samples were based on the sampling locations at
the study sites. The Kd measurements for the Spesutie and Transonic soil samples were
conducted with site-specific surface water samples. Because no representative surface water
existed at the Yuma site, the soil was equilibrated with tap water. The soil samples were
equilibrated in a ratio of 30 ml/1 g with water samples spiked with 200 |ig/l uranium. The
water/soil mixtures were sampled at 7 and 30 days. The Kd results are given in Table J.5. The
Kd values reported for the 30-day samples are 4360 (pH 6.8), 328 (pH 5.6), and 54 ml/g (pH 8.0),
respectively, for the Spesutie, Transonic, and Yuma soils. The lower Kd values measured for the
Yuma Soil samples were attributed to carbonate complexation of the dissolved uranium.
Giblin (1980) determined the Kd values for uranium sorption on kaolinite as a function of pH in a
synthetic groundwater. The measurements were conducted at 25°C using a synthetic
groundwater (Ca-Na-Mg-Cl-S04) containing 100 |ig/l uranium. Ten milliliters of solution was
mixed with 0.01 g of kaolinite for a solution-to-solid ratio of 1,000 ml/1 g. The pH of the
suspension was adjusted to cover a range from 3.8 to 10. Uranium Kd values from Giblin (1980,
Figure 1) are given in Table J.5.1 Giblin's results indicate that adsorption of uranium on
kaolinite in this water composition was negligible below pH 5. From pH 5 to 7, the uranium Kd
values increase to a maximum of approximately 37,000 ml/g. At pH values from 7 to 10, the
uranium adsorption decreased.
Kaplan et al. (1998) investigated the effects of U(VI) concentration, pH, and ionic strength on
the adsorption of U(VI) to a natural sediment containing carbonate minerals. The sediments
used for the adsorption measurements were samples of a silty loam and a very coarse sand taken,
respectively, from Trenches AE-3 and 94 at DOE's Hanford Site in Richland, Washington.
Groundwater collected from an uncontaminated part of the Hanford Site was equilibrated with
each sediment in a ratio of 2 ml/1 g for 14 or 30 days. The Kd values listed in Kaplan et al.
(1998) are given in Table J.5. The adsorption of U(VI) was determined to be constant for
concentrations between 3.3 and 100 |ig/l UO, at pH 8.3 and an ionic strength of 0.02 M. This
result indicates that a linear Kd model could be used to describe the adsorption of U(VI) at these
conditions. In those experiments where the pH was greater than 10, precipitation of
U(VI)-containing solids occurred, which resulted in apparent Kd values greater than 400 ml/g.
Kaplan et al. (1996) measured the Kd values for U(VI) and several other radionuclides at
geochemical conditions being considered in a performance assessment for the long-term disposal
of radioactive low-level waste in the unsaturated zone at DOE's Hanford Site in Richland,
1 The uranium Kd values listed in Table J.5 for Giblin (1980) were provided by E. A. Jenne
(PNNL, retired) based on work completed for another research project. The Kd values were
generated from digitization of the Kd values plotted in Giblin (1980, Figure 1).
J.5
-------
Washington. The studies included an evaluation of the effects of pH, ionic strength, moisture
content, and radionuclide concentration on radionuclide adsorption behavior. Methods used for
the adsorption measurements included saturated batch adsorption experiments, unsaturated batch
adsorption experiments, and unsaturated column adsorption experiments based on the
Unsaturated Flow Apparatus (UFA). The measurements were conducted using uncontaminated
pH 8.46 groundwater and the <2-mm size fraction of sediment samples collected from the
Hanford Site. The sediment samples included TBS-1 Touchet Bed sand, Trench AE-3 silty
loam, Trench-8 medium coarse sand, and Trench-94 very coarse sand. Dominant minerals
identified in the clay-size fraction of these sediment samples included smectite, illite,
vermiculite, and plagioclase. The reader should refer to Table 2.3 in Kaplan et al. (1996) for a
listing of the physical and mineralogical properties of these sediment samples. Uranium Kd
values estimated from results plotted in Kaplan et al. [1996, Figure 3.1 (400-day contact),
Figure 3.2 (all values as function of dissolved uranium concentrations), and Figure 3.5
(100 percent saturation values) are listed in Table J.5. Their results show that U(VI) Kd values
increased with increasing contact time with the sediments. For the concentration range from 3.3
to 100 |_ig/l dissolved uranium, the U(VI) Kd values were constant. The U(VI) Kd values
increased from 1.1 to 2.2 ml/g for pH values of 8 and 10, respectively, for these site-specific
sediments and geochemical conditions. Kaplan et al. noted that, at pH values above
approximately 10, the measured Kd values were affected by precipitation of uranium solids.
Their measurements also indicated that U(VI) Kd values varied as a function of moisture content,
although the trend differed based on sediment type. For a coarse-grained sediment, Kaplan et al.
noted the Kd values increased with increasing moisture saturation. However, the opposite trend
was observed for the U(VI) Kd values for fine-grained sediments. Kaplan et al. proposed that
this behavior was related to changes in tortuosity and effective porosity within the fine pore
spaces.
Kaplan and Serne (1995, Table 6.1) report Kd values for the adsorption of uranium on loamy
sand sediment taken from Trench 8 at DOE's Hanford Site in Richland, Washington. The
measurements were made using a column technique at unsaturated conditions (7 to 40 percent
saturated), neutral-to-high pH, low organic material concentrations, and low ionic strength
(I<0.1). The aqueous solutions consisted of a sample of uncontaminated groundwater from the
Hanford Site. The Kd values listed in Kaplan and Serne (1995) are given in Table J.5. The Kd
values ranged from 0.08 to 2.81 ml/g, and typically increase with increasing degree of column
saturation. Kaplan and Serne noted that Kd values measured using a batch technique are usually
greater than those obtained using the column technique due to the greater residence time and
greater mixing of the sediment and aqueous phase associated with the batch method.
Lindenmeier et al. (1995) conducted a series of flow-through column tests to evaluate
contaminant transport of several radionuclides through sediments under unsaturated (vadose
zone) conditions. The sediments were from the Trench 8 (W-5 Burial Ground) from DOE's
Hanford Site in Richland, Washington. The <2-mm size fraction of the sediment was used for
the measurements. The <2-mm size fraction had a total cation exchange capacity (CEC) of
5.2 meq/100 g, and consisted of 87 percent sand, 7 percent silt, and 6 percent clay-size materials.
Mineralogical analysis of <2-mm size fraction indicated that it consisted of 43.0 wt.% quartz,
J.6
-------
26.1 wt.% plagioclase feldspar, and minor amounts of other silicate, clay, hydrous oxide, and
carbonate minerals. The column tests were run using a site-specific groundwater and standard
saturated column systems, commercial and modified Wierenga unsaturated column systems, and
the Unsaturated Flow Apparatus (UFA). The results of the column tests indicated that the Kd
values for uranium on this sediment material decrease as the sediment becomes less saturated. A
Kd value of 2 ml/g was determined from a saturated column test conducted at a pore water
velocity of 1.0 cm/h and residence time of 1.24 h. However, at 29 percent water saturation, the
measured Kd value decreases by 70 percent to 0.6 ml/g (pore water velocity of 0.3 cm/h and
residence time of 20.6 h). The Kd values listed in Lindenmeier et al. (1995, Table 4.1) are given
in Table J. 5.
Salter et al. (1981) investigated the effects of temperature, pressure, groundwater composition,
and redox conditions on the sorption behavior of several radionuclides, including uranium, on
Columbia River basalts. Uranium Kd values were determined at 23 and 60°C under oxidizing
and reducing conditions using a batch technique. The measurements were conducted with
2 synthetic groundwater solutions (GR-1 and GR-2) that have compositions representative of the
groundwater present in basalt formations at DOE's Hanford Site, Richland, Washington. The
GR-1 and GR-2 solutions represent a pH 8 sodium bicarbonate-buffered groundwater and a
pH 10 silicic acid-buffered groundwater. The synthetic groundwater solutions were mixed with
the crushed basalt material (0.03-0.85 mm size fraction) in a ratio of 10 ml/1 g. The contact time
for the measurements was approximately 60 days. The Kd values were determined for initial
concentrations of 1.0x10"4, 1.0x10"5, l.OxlO"6, l.OxlO"7, and 2.15xl0"8 M uranium. The Kd
values listed in Table J.5 from Salter et al. (1981) include only those for 23°C under oxidizing
conditions. The reader is referred to Salter et al. (1981) for a description of the measurement
procedure and results for reducing conditions.
Serkiz and Johnson (1994) (and related report by Johnson et al., 1994) investigated the
partitioning of uranium on soil in contaminated groundwater downgradient of the F and H Area
Seepage Basins at DOE's Savannah River Site in South Carolina. Their study included
determination of an extensive set of field-derived Kd values for 238U and 235U for 48
soil/porewater samples. The Kd values were determined from analyses of 238U and 235U in soil
samples and associated porewaters taken from contaminated zones downgradient of the seepage
basins. It should be noted that the mass concentration of 235U is significantly less than (e.g., <1
percent) the concentration of 238U in the same soil sample and associated porewater. Serkiz and
Johnson used the geochemical code MINTEQA2 to model the aqueous complexation and
adsorption of uranium in their analysis of migration and partitioning in the contaminated soils.
Soil/porewater samples were collected over a range of geochemical conditions (e.g., pH,
conductivity, and contaminant concentration). The field-derived uranium Kd listed for 238U and
235U by Serkiz and Johnson are given in Table J.5. The uranium Kd values varied from 1.2 to
34,000 ml/g over a pH range from approximately 3 to 6.7 (Figure J.l). The reader should note
that the field-derived Kd values in Figures J. 1, J.2, and J.3 are plotted on a logarithmic scale. At
these site-specific conditions, the Kd values indicate that uranium adsorption increases with
increasing pH over the pH range from 3 to 5.2. The adsorption of uranium is at a maximum at
approximately pH 5.2, and then decreases with increasing pH over the pH range from 5.2 to 6.7.
J.7
-------
Serkiz and Johnson found that the field-derived Kd values for 238U and 235U were not well
correlated with the weight percent of clay-size particles (Figure J.2) or CEC (Figure J.3) of the
soil samples. Based on the field-derived Kd values and geochemical modeling results, Serkiz
and Johnson proposed that the uranium was not binding to the clays by a cation exchange
reaction, but rather to a mineral surface coating with the variable surface charge varying due to
the porewater pH.
100,000
10,000
1,000
100
10
n D ~
~ 9 B
~ ~ n
~
~
J3,
e 8i
pICcF3
n
~
o
~
5
PH
Figure J.l. Field-derived Kd values for 238U and 235U from Serkiz and
Johnson (1994) plotted as a function of porewater pH for
contaminated soil/porewater samples. [Square and circle
symbols represent field-derived Kd values for 238U and 235U,
respectively. Solid symbols represent minimum Kd values for
238U and 235U that were based on minimum detection limit
values for the concentrations for the respective uranium
isotopes in porewaters associated with the soil sample.]
J.8
-------
-2?
J
5
100,000
10,000
1,000
100
10
- ~ ~ ¦
~
• n ^Tl
~
s\
¦
~
¦
~
:
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• ~
•
•
^ en
; q5|§
T&S, ,
1 1 I I 1 1 1 I
1 1 1 1 I 1
1 1 1 I 1 I
10 20 30
Clay-Size Particle Content (wt%)
40
50
Figure 3.2. Field-derived Kd values for 238U and 235U from Serkiz and Johnson (1994)
plotted as a function of the weight percent of clay-size particles in the
contaminated soil/porewater samples. [Square and circle symbols represent
field-derived Kd values for 238U and 235U, respectively. Solid symbols
represent minimum Kd values for 238U and 235U that were based on minimum
detection limit values for the concentrations for the respective uranium
isotopes in porewaters associated with the soil sample.]
J.9
-------
ox
1
3
100,000
10,000
1,000
100
10
~
~
<3i
"8D*
y O
~1°
m° @p
n a 0
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o
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~
¦ I I I I I I I I I I I I I I I I I I
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10 15
CEC (meq/kg)
20
25
Figure J.3. Field-derived Kd values for 238U and 235U plotted from Serkiz and Johnson
(1994) as a function of CEC (meq/kg) of the contaminated soil/porewater
samples. [Square and circle symbols represent field-derived Kd values for
238U and 235U, respectively. Solid symbols represent minimum Kd values for
238U and 235U that were based on minimum detection limit values for the
concentrations for the respective uranium isotopes in porewaters associated
with the soil sample.]
Serne et al. (1993) determined Kd values for uranium and several other radionuclides at
geochemical conditions associated with sediments at DOE's Hanford Site in Richland,
Washington. The Kd values were measured using the batch technique with a well-characterized
pH 8.3 groundwater and the <2-mm size fraction of 3 well-characterized sediment samples from
the Hanford Site. The sediment samples included TBS-1 Touchet Bed sand, CSG-1 coarse
sand/gravel, and Trench-8 medium coarse sand. The <2-mm size fraction of 3 samples consisted
of approximately 70 to 90 wt.% plagioclase feldspar and quartz, and minor amounts of other
silicate, clay, hydrous oxide, and carbonate minerals. The solution-to-solid ratio was fixed at
30 ml/1 g. The contact time for adsorption measurements with TBS-1, CSG-1, and Trench-8
were, 35, 35, and 44 days, respectively. The average Kd values tabulated for uranium in Serne el
al. (1993) are given in Table J.5.
J.10
-------
Sheppard and Thibault (1988) investigated the migration of several radionuclides, including
uranium, through 3 peat1 types associated with mires2 typical of the Precambrian Shield in
Canada. Cores of peat were taken from a floating sphagnum mire (samples designated PCE,
peat-core experiment) and a reed-sedge mire overlying a clay deposit (samples designated SCE,
sedge-core experiment). Uranium Kd values were determined by in situ and batch laboratory
methods. The in situ Kd values were calculated from the ratio of uranium in the dried peat and
associated porewater solutions. The batch laboratory measurements were conducted over an
equilibration period of 21 days. The in-situ and batch-measured uranium Kd values tabulated in
Sheppard and Thibault (1988) are listed in Table J.5. Because the uranium Kd values reported
by Sheppard and Thibault (1988) represent uranium partitioning under reducing conditions,
which are beyond the scope of our review, these Kd values were not included in Figure J.4.
Sheppard and Thibault (1988) noted that the uranium Kd for these 3 peat types varied from 2,00
to 19,000 ml/g, and did not vary as a function of porewater concentration. The laboratory
measured Kd values were similar to those determined in situ for the SCE peat sample.
Thibault et al. (1990) present a compilation of soil Kd values prepared as support to 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 uranium, that would
be present in the nuclear fuel waste inventory. Some of the uranium Kd values listed by Thibault
et al. were collected from references that were not available during the course of our review.
These sources included studies described in reports by M. I. Sheppard, a coauthor of Thibault et
al. (1990), and papers by Dahlman et al. (1976), Haji-Djafari etal. (1981), Neiheisel (1983),
Ran<;on (1973) and Seeley and Kelmers (1984). The uranium Kd values, as listed in Thibault et
al. (1990), taken for these sources are included in Table J.5.
Warnecke and coworkers (Warnecke etal., 1984, 1986, 1988, 1994; Warnecke andHild, 1988;
and others) published several papers that summarize the results of radionuclide migration
experiments and adsorption/desorption measurements (Kd values) that were conducted in support
of Germany's investigation of the Gorleben salt dome, Asse II salt mine, and former Konrad iron
ore mine as disposal sites for radioactive waste. Experimental techniques included batch and
recirculation methods as well as flow-through and diffusion experiments. The experiments were
designed to assess the effects of parameters, such as temperature, pH, Eh, radionuclide
concentration, complexing agents, humic substances, and liquid volume-to-soil mass ratio, on
radionuclide migration and adsorption/desorption. These papers are overviews of the work
completed in their program to date, and provide very few details on the experimental designs and
individual results. There are no pH values assigned to the Kd values listed in these overview
1 Peat is defined as "an unconsolidated deposit of semicarbonized plant remains in a water
saturated environment" (Bates and Jackson, 1980).
2 A mire is defined as "a small piece of marshy, swampy, or boggy ground" (Bates and
Jackson, 1980).
J. 11
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papers. Warnecke et al. (1984) indicated that the measured pH values for the locations of soil
and groundwater samples at Gorleben site studies range from 6 to 9.
Warnecke et al. (1994) summarize experiments conducted during the previous 10 years to
characterize the potential for radionuclide migration at site-specific conditions at the Gorleben
site. Characteristic, minimum, and maximum Kd values tabulated by Warnecke et al. (1994,
Table 1) for uranium adsorbed to sandy and clayish sediments in contact with fresh or saline
waters are listed below in Table J. 1. No pH values were assigned to the listed Kd values.
Warnecke et al. noted that the following progression in uranium Kd values as function of
sediment type was indicated:
Kd (Clay) > Kd (Marl1) > Kd (Sandy) .
Warnecke and Hild (1988) present an overview of the radionuclide migration experiments and
adsorption/desorption measurements that were conducted for the site investigations of the
Gorleben salt dome, Asse II salt mine, and Konrad iron ore mine. The uranium Kd values listed
in Warnecke and Hild are identical to those presented in Warnecke et al. (1994). The uranium
Kd values (ml/g) listed by Warnecke and Hild (1988, Table II) for sediments and different water
types for the Konrad site are: 4 (Quaternary fresh water), 6 (Turonian fresh water), 6
(Cenomanian saline water), 20 [Albian (Hauterivain) saline water], 1.4 [Albian (Hils) saline
water], 2.6 (Kimmeridgian saline water), 3 (Oxfordian saline water), and 3 [Bajocian (Dogger)
saline water], Warnecke and Hild (1988, Table III) list minimum and maximum uranium Kd
values (0.54-15.2 ml/g) for 26 rock samples from the Asse II site. No pH values were assigned
to any of the tabulated Kd values, and no descriptions were given regarding the mineralogy of the
site sediment samples. Warnecke and Hild noted that sorption measurements for the Konrad
sediments, especially for the consolidated material, show the same trend as those for the
Gorleben sediments.
Table J.l. Uranium Kd values (ml/g) listed by Warnecke et al. (1994, Table 1).
Sediment
Type
Fresh Water
Saline Water
Typical
Kd Value
Minimum
Kd Value
Maximum
Kd Value
Typical
Kd Value
Minimum
Kd Value
Maximum
Kd Value
Sandy
27
0.8
332
1
0.3
1.6
Clayish
17
8.6
100
14- 1,400
14.1
1,400
1 Marl is defined as "an earthy substance containing 35-65 percent clay and 65-35 percent
carbonate formed under marine or freshwater conditions" (Bates and Jackson, 1980).
J. 12
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Warnecke et al. (1986) present an overview of the radionuclide migration experiments and
adsorption/desorption measurements that were conducted for the Gorleben salt dome, and
Konrad iron ore mine. The tabulated Kd values for the Gorleben and Konrad site sediments and
waters duplicate those presented Warnecke et al. (1994) and Warnecke and Hild (1988).
Warnecke et al. (1984) present a short summary of radionuclide sorption measurements that
were conducted by several laboratories in support of the Gorleben site investigation. Sediment
(especially sand and silt) and water samples were taken from 20 locations that were considered
representative of the potential migration path for radionuclides that might be released from a
disposal facility sited at Gorleben. The minimum and maximum Kd values listed by Warnecke et
al. (1984, Table III) are 0.5 and 3,000 ml/g, respectively (note that these values are not listed as a
function of pH).
Zachara et al. (1992) studied the adsorption of U(VI) on clay-mineral separates from subsurface
soils from 3 DOE sites. The materials included the clay separates (<2 |im fraction) from the
Kenoma Formation (Feed Materials Production Center, Fernald, Ohio), Ringold Formation
(Hanford Site, Richland, Washington), and Cape Fear Formation (Savannah River Site, Aiken,
South Carolina). Prior to the measurements the clay separates were treated with dithionite-
citrate buffer and hydrogen peroxide to remove amorphous ferric hydroxides and organic
materials. The measurements used clay suspensions (~ 1 meq of charge/1) spiked with 2 mg/1
(8.6 |imol/l) uranium and Ca(C104)2 or NaC104 as the electrolyte. The pH values of the
suspensions were adjusted over the pH range from 4.5 to 9.0 using sodium hydroxide. The
measurements were completed in a glovebox under an inert atmosphere to eliminate effects from
aqueous complexation of U(VI) by dissolved carbonate. Uranium Kd values calculated from
values of percent uranium adsorbed versus pH (Zachara et al., 1992, Figures 6 and 7) for the
Kenoma and Ringold clays are listed in Table J.5.1 The adsorption results for the Cape Fear clay
isolate were essentially the same as those for the Kenoma clay (Zachara et al., 1992, Figures 8).
The results for the Kenoma clay isolate show a strong dependence of uranium adsorption as a
function of ionic strength that is opposite to that expected for competitive sorption between
uranium and the electrolyte cation. Zachara et al. (1992) suggest that this increase in uranium
adsorption with increasing ionic strength may be due to the ionic strength dependence of the
hydrolysis of the uranyl ion.
J. 2.3 Uranium Kd Studies on Single Mineral Phases
1 The uranium Kd values listed in Table J.5 for Zachara et al. (1992) were provided by E. A.
Jenne (PNNL, retired) based on work completed for another research project. The Kd values
were derived from percent uranium adsorbed values generated from digitization of data plotted
in Zachara et al. (1992, Figures 6 and 7) for the Kenoma and Ringold clay isolates. Due the
inherent uncertainty and resulting exceptionally large Kd values, Jenne did not calculate Kd
values from any percent uranium adsorbed values that were greater 99 percent.
J.13
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Anderson et al. (1982) summarize an extensive study of radionuclides on igneous rocks and
related single mineral phases. They report Kd values for U(VI) sorption on apatite, attapulgite
(also known as palygorskite), biotite, montmorillonite, and quartz. The Kd values were
determined using a batch technique using 10"7-10"9 mol/1 uranium concentrations, synthetic
groundwater, and crushed (0.045-0.063 mm size fraction) mineral and rock material. The
solution-to-solid ratio used in the experiments was 50 ml/1 g. The synthetic groundwater had a
composition typical for a Swedish deep plutonic groundwater. Uranium Kd values from
Anderson et al. (1982, Figure 6a) are given in Table J.5.1
Ames et al. (1983a,b) investigated the effects of uranium concentrations, temperature, and
solution compositions on the sorption of uranium on several well-characterized secondary and
sheet silicate minerals. The secondary phases studied by Ames et al. (1983a, oxide analyses
listed in their Table 3) included clinoptilotite, glauconite, illite, kaolinite, montmorillonite,
nontronite, opal, and silica gel. The sheet silicate minerals used by Ames et al. (1983b, oxide
analyses listed in their Table 1) consisted of biotite, muscovite, and phlogopite. The sorption of
uranium on each mineral phase was measured with 2 solutions (0.01 M NaCl and 0.01 M
NaHC03) using 4 initial uranium concentrations. The initial uranium concentrations used for the
25°C experiments included l.OxlO"4, l.OxlO"5, 1.4xl0"6, and 4.4xl0"7 mol/1 uranium. The batch
experiments were conducted under oxidizing conditions at 5, 25, and 65°C in an environmental
chamber. Solutions were equilibrated with the mineral solids in a ratio of 10 ml/1 g. A
minimum of 30 days was required for the mineral/solution mixtures to reach steady state
conditions. Uranium Kd values calculated from the 25°C sorption results given in Ames et al.
(1983a, Table 6) are listed in Table J. 5.
Ames et al. (1983c) studied the effects of uranium concentrations, temperature, and solution
compositions on the sorption of uranium on amorphous ferric oxyhydroxide. The sorption of
uranium on amorphous ferric oxyhydroxide was measured with 2 solutions (0.01 M NaCl and
0.01 M NaHC03) using 4 initial uranium concentrations. The initial uranium concentrations
used for the 25°C experiments included l.OlxlO"4, 1.05xl0"5, 1.05xl0"6, and 4.89xl0"7 mol/1
uranium for the 0.01 MNaCl solution, and l.OlxlO"4, 1.05xl0"5, 1.53xl0"6, and 5.46xl0"7 mol/1
uranium for the 0.01 M NaHC03 solution. The batch experiments were conducted under
oxidizing conditions at 25 and 60°C. The solutions were equilibrated for 7 days with the
amorphous ferric oxyhydroxide in a ratio 3.58 1/g of iron in the solid. Uranium Kd values
calculated from the 25°C sorption results given in Ames et al. (1983c, Table II) are listed in
Table J. 5. Reflecting the high adsorptive capacity of ferric oxyhydroxide, the Kd values for the
25°C measurements range from approximately 2xl06 ml/g for the 0.01 M NaCl solution to
approximately 3xl04 ml/g for the 0.01 M NaHC03 solution.
1 The uranium Kd values listed in Table J.5 for Anderson et al. (1982) were provided by E. A.
Jenne (PNNL, retired) based on work completed for another research project. The Kd values
were generated from digitization of the Kd values plotted in Anderson et al. (1982, Figure 6a).
J. 14
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Borovec (1981) investigated the adsorption of U(VI) and its hydrolytic complexes at 20°C and
pH 6.0 on fine-grained kaolinite, illite, and montmorillonite. The results indicate that the Kd
values increase with decreasing concentrations of dissolved uranium. At uranium concentrations
less than 10"4 mol/1, the uranium Kd values for the individual minerals were constant. The Kd
values determined at 20°C and pH 6.0 ranged from 50 to 1,000. The values increased in the
sequence Kd (kaolinite) < Kd (illite) < Kd (montmorillonite). Borovec presents the following
linear equations for the maximum sorption capacity of uranium (am, in meq/100 g) on clays at
20°C and pH 6.0 with respect to CEC (in meq/100 g),
am = 0.90 CEC+1.56 (r = 0.99522),
and specific surface (A, in m2/g) of clays,
am = 0.11 A+ 2.05 (r = 0.97232).
2.4 Published Compilations Containing Kd Values for Uranium
Baes and Sharp (1983) present a model developed for annual-average, order-of-magnitude
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, in their
leaching model. A literature review was completed to evaluate appropriate distributions for Kd
values for various solutes, including uranium. Because Baes and Sharp (1983) are cited
frequently as a source of Kd values in other published Kd reviews (e.g, Looney el al., 1987;
Sheppard and Thibault, 1990), the uranium Kd values listed by Baes and Sharp are reported here
for the sake of completeness. Based of 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 45 ml/g
for uranium based on 24 uranium Kd values from 10.5 to 4,400 ml/g for agricultural soils and
clays in the pH range from 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) describe the estimation of 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 uranium, Looney et al. list a "recommended" Kd of
39.8 (10L6) ml/g, and a range for its Kd values of 0.1 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 using them in assessments at other sites. Nonetheless,
such data are often used as "default values" in radionuclide migration assessment calculations,
and are therefore listed here for the sake of completeness. It should be noted that the work of
Looney et al. (1987) predates the uranium-migration and field-derived uranium Kd study
reported for contaminated soils at the Savannah River Site by Serkiz and Johnston (1994)
(described above).
J. 15
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McKinley and Scholtis (1993) compare radionuclide Kd sorption databases used by different
international organizations for performance assessments of repositories for radioactive wastes.
The uranium Kd values listed in McKinley and Scholtis (1993, Tables 1, 2, and 4) are listed in
Table J.2. 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
J.2. The organizations listed in the tables in McKinley and Scholtis (1993) 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 PAGRIS SAFIR (Safety Assessment and Feasiblity 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).
J.16
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Table J.2. Uranium Kd values 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
Argillaceous (Clay)
Crystalline Rock
Soil/Soil
Sorbing
Material
Kd
(ml/g)
Sorbing
Material
Kd
(ml/g)
Sorbing
Material
K„
(ml/g)
AECL
Bentonite-Sand
100
Granite
5
Soil/Sediment
20
GSF
Sediment
2
IAEA
Pelagic Clay
500
KBS-3
Bentonite
120
Granite
5,000
NAGRA
Bentonite
1,000
Granite
1,000
Soil/Sediment
20
Clay
5,000
Soil/Sediment
100
NIREX
Clay Mudstone
10
NRC
Clay, Soil Shale
20
Granite
5
Basalt
4
Tuff
300
NRPB
Clay
300
Soil/Sediment
300
PAGIS
Bentonite
90
Soil/Sediment
1,700
Subseabed
100
PAGIS SAFIR
Clay
600
PSE
Sediment
0.02
RIVM
Sandy Clay
10
SKI
Bentonite
200
Granite
5,000
TVO
Bentonite
90
Crystalline
Rock, Reducing
200
Soil/Sediment
500
Baltic Sea
Sediment
500
Crystalline
Rock, Real.
5
Ocean Sediment
500
Lake Sediment
500
UKDoE
Clay
200
Soil/Sediment
50
Coastal Marine
Water
1000
n a similar comparison of sorption databases for use in performance assessments of radioactive
waste repositories, Stenhouse and Pottinger (1994) list "realistic" Kd values (ml/g) for uranium
J.17
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in crystalline rock/water systems of 1,000 (NAGRA), 5,000 [Svensk Karnbranslehantering AB
(Nuclear Fuel and Waste Management Company), Sweden; SKB], 1000 (TVO), and 6 (Canadian
Nuclear Fuel Waste Management Programme, CNFWM). For bentonite/groundwater systems,
they list 5,000 (NAGRA), 3,000 (SKB), and 500 (TVO). The reader should refer to sources
cited in Stenhouse and Pottinger for details regarding the source, derivation, and measurement of
these values.
Thibault et al. (1990) [also summarized in Sheppard and Thibault (1990)] updated a compilation
of soil Kd values that they published earlier (Sheppard et al., 1984). The compilations were
completed to support the assessments) of 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 uranium, that
would be present in the inventory associated with Canada's nuclear fuel wastes. Because
Thibault et al. (1990) and Sheppard and Thibault (1990) are frequently cited, their derived
uranium Kd values are reported here for the sake of completeness. The Kd values for each
element were categorized according to 4 soil texture types. These included sand {i.e., contains
>70 percent sand-size particles), clay {i.e., contains >35 percent clay-size particles), loam {i.e.,
contains an even distribution of sand-, clay-, and silt-size particles, or <80 percent silt-size
particles), and organic {i.e., contains >30 percent organic matter and are either classic peat or
muck sediments, or the litter horizon of a mineral sediment). Based on their previous
evaluations, Thibault et al. In-transformed and averaged the compiled Kd values to obtain a
single geometric mean Kd value for each element for each soil type. The Kd values for each soil
type and the associated range of Kd values listed for uranium by Thibault et al. (1990) are given
in Table J.3.
Table J.3. Geometric mean uranium Kd values derived by Thibault et al. (1990) for
sand, loam, clay, and organic soil types.
Soil Type
Geometric
Mean Kd
Values (ml/g)
Observed Range of
Kd Values (ml/g)
Number of
Kd Values
Sand
35
0.03 - 2,200
24
Loam
15
0.2-4,500
8
Clay
1,600
46 -395,100
7
Organic
410
33 - 7,350
6
J. 18
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J.3.0 Approach in Developing Kd Look-Up Table
The uranium Kd values listed in Table J.5 are plotted in Figure J.4 as a function of pH. The Kd
values exhibit large scatter. This scatter increases from approximately 3 orders of magnitude at
pH values below pH 5, to approximately 3 to 4 orders of magnitude from pH 5 to 7, and
approximately 4 to 5 orders of magnitude at pH values from pH 7 to 9. This comparison can be
somewhat misleading. At the lowest and highest pH regions, it should be noted that 1 to 2 orders
of the observed variability actually represent uranium Kd values that are less than 10 ml/g. At
pH values less than 3.5 and greater than 8, this variability includes extremely small Kd values of
less than 1 ml/g.
Figure J.4. Uranium Kd values used for development of Kd look-up table.
[Filled circles represent Kd values listed in Table J.5. Open
symbols (joined by dotted line) represent Kd maximum and
minimum values estimated from uranium adsorption
measurements plotted by Waite et al. (1992) for ferrihydrite
(open squares), kaolinite (open circles), and quartz (open
triangles). The limits for the estimated maximum and
minimum Kd values based on the values in Table J. 5 and
those estimated from Waite et al. (1992) are given by the "x"
symbols joined by a solid line.]
J.19
-------
J. 3.1 Kd Values as a Function ff pH
Although the uranium Kd values in Figure J.4 exhibit a great deal of scatter at any fixed pH
value, the Kd values show a trend as a function of pH. In general, the adsorption of uranium by
soils and single-mineral phases is low at pH values less than 3, increases rapidly with increasing
pH from pH 3 to 5, reaches a maximum in adsorption in the pH range from pH 5 to 8, and then
decreases with increasing pH at pH values greater than 8. This trend is similar to the in situ Kd
values reported by Serkiz and Johnson (1994) (see Figure J. 1), and percent adsorption values
measured for uranium on single mineral phases as described above and those reported for iron
oxides (Duff and Amrheim, 1996; Hsi and Langmuir, 1985; Tripathi, 1984; Waite etal., 1992,
1994; and others), clays (McKinley et al., 1995; Turner et al., 1996; Waite et al., 1992; and
others), and quartz (Waite et al., 1992). The adsorption data are similar to those of other
hydrolyzable metal ions with a sharp pH edge separating low adsorption at low pH from high
adsorption at higher pH values. As discussed in the surface complexation laboratory and
modeling studies [e.g., Tripathi (1984), Hsi and Langmuir (1985), Waite etal. (1992, 1994), and
Duff and Amrheim (1996)], this pH-dependent behavior is related to the pH-dependent surface
charge properties of the soil minerals and complex aqueous speciation of dissolved U(VI),
especially near and above neutral pH conditions where dissolved U(VI) forms strong anionic
uranyl-carbonato complexes with dissolved carbonate.
J. 3.2 Kd Values as a Function of Mineralogy
In addition to the sources of error and variability discussed above, the scatter in Kd values in
Figure J.4 is also related to heterogeneity in the mineralogy of the soils. Soils containing larger
percentages of iron oxide minerals and mineral coatings and/or clay minerals will exhibit higher
sorption characteristics than soils dominated by quartz and feldspar minerals. This variability in
uranium adsorption with respect to mineralogy is readily apparent in uranium Kd values
calculated from adsorption measurements (reported as percent uranium adsorbed versus pH) for
ferrihydrite, kaolinite, and quartz by Waite etal. (1992).
Uranium Kd values were estimated1 from the plots of percent uranium adsorption given for
ferrihydrite, kaolinite, and quartz by Waite et al. (1992). To estimate the maximum variability
that should be expected for the adsorption of uranium by different mineral substrates, Kd values
were calculated from plots of uranium adsorption data for ferrihydrite and kaolinite (minerals
with high adsorptive properties) that exhibited the maximum adsorption at any pH from 3 to 10,
and for quartz (a mineral with low adsorptive properties) that exhibited the minimum adsorption
1 The reader is cautioned that significant uncertainty may be associated with Kd values
estimated in this fashion because of the extreme solution-to-solid ratios used in some of these
studies, especially for highly adsorptive iron-oxide phases, and errors related to estimating the
concentrations of sorbed and dissolved uranium based on values for the percent of absorbed
uranium near 0 or 100 percent, respectively.
J.20
-------
at any pH. These estimated Kd values are shown, respectively, as open squares, circles, and
triangles (and joined by dotted lines) in Figure J.4. The difference in the maximum and
minimum Kd values is nearly 3 orders of magnitude at any fixed pH value in the pH range from 3
to 9.5. At pH values less than 7, the uranium Kd values for ferrihydrite and quartz calculated
from data in Waite et al. (1992) bound more than 95 percent of the uranium Kd values gleaned
from the literature. Above pH 7, the calculated uranium Kd values for ferrihydrite and kaolinite
effectively bound the maximum uranium Kd values reported in the literature.. In terms of
bounding the minimum Kd values, the values calculated for quartz are greater than several data
sets measured by Kaplan et al. (1996, 1998), Lindenmeirer et al. (1995), and Serne et al. (1993)
for sediments from the Hanford Site in Richland, Washington which typically contain a
significant quality of quartz and feldspar minerals. It should also be noted that some of the
values listed from these studies represent measurements of uranium adsorption on Hanford
sediments under partially saturated conditions.
J. 3.3 Kd Values As A Function Of Dissolved Carbonate Concentrations
As noted in several studies summarized above and in surface complexation studies of uranium
adsorption by Tripathi (1984), Hsi and Langmuir (1985), Waite et al. (1992, 1994), McKinley et
al. (1995), Duff and Amrheim (1996), Turner et al. (1996), and others, dissolved carbonate has a
significant effect on the aqueous chemistry and solubility of dissolved U(VI) through the
formation of strong anionic carbonato complexes. In turn, this complexation affects the
adsorption behavior of U(VI) at alkaline pH conditions. Even differences in partial pressures of
C02 have a major affect on uranium adsorption at neutral pH conditions. Waite et al. (1992,
Figure 5.7), for example, show that the percent of U(VI) adsorbed onto ferrihydrite decreases
from approximately 97 to 38 percent when C02 is increased from ambient (0.03 percent) to
elevated (1 percent) partial pressures. In those adsorption studies that were conducted in the
absence of dissolved carbonate (see surface complexation modeling studies listed above),
uranium maintains a maximum adsorption with increasing pH as opposed to decreasing with
increasing pH at pH values near and above neutral pH. Although carbonate-free systems are not
relevant to natural soil/groundwater systems, they are important to understanding the reaction
mechanisms affecting the aqueous and adsorption geochemistry of uranium.
It should be noted that it is fairly common to see figures in the literature or at conferences where
uranium adsorption plotted from pH 2 to 8 shows maximum adsorption behavior even at the
highest pH values. Such plots may mislead the reader into thinking that uranium adsorption
continues this trend {i.e., maximum) to even higher pH conditions that are associated with some
groundwater systems and even porewaters derived from leaching of cementitious systems.
Based on the uranium adsorption studies discussed above, the adsorption of uranium decreases
rapidly, possibly to very low values, at pH values greater than 8 for waters in contact with C02
or carbonate minerals .
No attempt was made to statistically fit the Kd values summarized in Table J.5 as a function of
dissolved carbonate concentrations. Typically carbonate concentrations were not reported and/or
discussed, and one would have to make assumptions about possible equilibrium between the
J.21
-------
solutions and atmospheric or soil-related partial pressures of C02 or carbonate phases present in
the soil samples. As will be discussed in a later section, the best approach to predicting the role
of dissolved carbonate in the adsorption behavior of uranium and derivation of Kd values is
through the use of surface complexation modeling techniques.
J. 3.4 Kd Values as a Function of Clay Content and CEC
No attempt was made to statistically fit the Kd values summarized in Table J.5 as a function of
CEC or concentrations of clay-size particles. The extent of clay concentration and CEC data, as
noted from information included in Table J.5, is limited to a few studies that cover somewhat
limited geochemical conditions. As discussed above, Serkiz and Johnson (1994) found no
correlation between their uranium in situ Kd values and the clay content (Figure J.2) or CEC
(Figure J.3) of their soils. Their systems covered the pH conditions from 3 to 7.
As noted in the studies summarized above, clays have an important role in the adsorption of
uranium in soils. Attempts have been made (e.g., Borovec, 1981) to represent this functionality
with a mathematical expression, but such studies are typically for limited geochemical
conditions. Based on the studies by Chisholm-Brause (1994), Morris etal. (1994), McKinley et
al. (1995), Turner et al. (1996), and others, uranium adsorption onto clay minerals is
complicated and involves multiple binding sites, including exchange and edge-coordination sites.
The reader is referred to these references for a detailed treatment of the uranium adsorption on
smectite clays and application of surface complexation modeling techniques for such minerals.
J. 3.5 Uranium Kd Look- Up Table
Given the orders of magnitude variability observed for reported uranium Kd values, a subjective
approach was used to estimate the minimum and maximum Kd values for uranium as a function
of pH. These values are listed in Table J.4. For Kd values at non-integer pH values, especially
given the rapid changes in uranium adsorption observed at pH values less than 5 and greater than
8, the reader should assume a linear relationship between each adjacent pair of pH-Kd values
listed in Table J.4.
Table J.4. Look-up table for estimated range of Kd values for uranium based on pH.
Kd
(ml/g)
pH
3
4
5
6
7
8
9
10
Minimum
<1
0.4
25
100
63
0.4
<1
<1
Maximum
32
5,000
160,000
1,000,000
630,000
250,000
7,900
5
J.22
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The minimum and maximum Kd values listed in Table J.4 were taken from the solid lines plotted
in Figure F.4. The area between the 2 solid lines contains more than 95 percent of uranium Kd
values collected in this review. The curve representing the minimum limit for uranium Kd values
is based on Kd values calculated (described above) for quartz from data given in Waite el al.
(1992) and the Kd values reported by Kaplan et al. (1996, 1998), Lindenmeirer et al. (1995), and
Serne et al. (1993). It is unlikely that actual Kd values for U(VI) can be much lower than those
represented by this lower curve. At the pH extremes along this curve, the uranium Kd values are
already very small. Moreover, if one considers potential sources of error resulting from
experimental methods, it is difficult to rationalize uranium Kd values much lower than this lower
boundary.
The curve representing the maximum limit for uranium Kd values is based on Kd values
calculated (described above) for ferrihydrite and kaolinite from data given in Waite et al. (1992).
It is estimated that the maximum boundary of uranium Kd values plotted in Figure J.4 is
conservatively high, possibly by an order of magnitude or more especially at pH values greater
than 5. This estimate is partially based on the distribution of measured Kd values plotted in
Figure J.4, and the assumption that some of the very large Kd measurements may have included
precipitation of uranium-containing solids due to starting uranium solutions being oversaturated.
Moreover, as noted previously, measurements of uranium adsorption onto crushed rock samples
may include U(VI)/U(IV) redox/precipitation reactions resulting from contact of dissolved U(VI)
with Fe(II) exposed on the fresh mineral surfaces.
J.4.0 Use of Surface Complexation Models to Predict Uranium Kd Values
As discussed in Chapter 4 and in greater detail in Volume I of this report, electrostatic surface
complexation models (SCMs) incorporated into chemical reaction codes, such as EPA's
MINTEQA2, may be used to predict the adsorption behavior of some radionuclides and other
metals and to derive Kd values as a function of key geochemical parameters, such as pH and
carbonate concentrations. Typically, the application of surface complexation models is limited
by the availability of surface complexation constants for the constituents of interest and
competing ions that influence their adsorption behavior.
The current state of knowledge regarding surface complexation constants for uranium adsorption
onto important soil minerals, such as iron oxides, and development of a mechanistic
understanding of these reactions is probably as advanced as those for any other trace metal. In
the absence of site-specific Kd values for the geochemical conditions of interest, the reader is
encouraged to apply this technology to predict bounding uranium Kd values and their
functionality with respect to important geochemical parameters.
Numerous laboratory surface complexation studies for uranium have been reported in the
literature. These include studies of uranium adsorption onto iron oxides (Duff and Amrheim,
1996; Hsi and Langmuir, 1985; Tripathi, 1984; Waite et al., 1992, 1994; and others), clays
(McKinley et al., 1995; Turner et al., 1996; Waite et al., 1992; and others), and quartz (Waite et
J.23
-------
al., 1992; and others). These references include derivation of the surface complexation constants
for surface coordination sites determined to be important.
In addition to these laboratory studies, there are numerous examples in the literature of the
application of surface complexation models and published binding constants to predict and
evaluate the migration of uranium in soil/groundwater systems. For example, KoB (1988)
describes the use of a surface complexation adsorption model to calculate the sorption of
uranium for soil-groundwater systems associated with the proposed site for a German geologic
radioactive waste repository at Gorleben. An apparent constant {i.e., apparent surface complex
formation constant based on bulk solution concentrations, Kapp) was derived for uranium
sorption using the MINEQL geochemical code and site-specific geochemical data for soil CEC
values, groundwater compositions, and measured uranium Kd values. Quartz (Si02) was the
main constituent in the soils considered in this study. Because the model incorporates the
aqueous speciation of uranium, it may be used tor compare Kd values for different soil systems
having equal sorption sites. The modeling results indicated that CEC, pH, ionic strength, and
dissolved carbonate concentrations were the main geochemical parameters affecting the sorption
of uranium in groundwater systems.
Puigdomenech and Bergstrom (1994) evaluated the use of surface complexation models for
calculating radionuclide sorption and Kd values in support of performance assessments studies of
geologic repositories for radioactive wastes. They used a triple layer surface complexation
model to predict the amount of uranium sorbed to a soil as a function of various environmental
parameters. They then derived Kd values based on the concentrations of adsorbed and dissolved
uranium predicted by the model. For the surface complexation modeling, they assumed (1) a
total uranium concentration of 10"5 mol/1, and (2) the adsorption of uranium on soil was
controlled by the soil concentration of iron oxyhydroxide solid, which was assumed to be 5
percent goethite [a-FeO(OH)]. Their modeling results indicated that pH, inorganic carbon {i.e.,
dissolved carbonate), and Eh (redox conditions) are major parameters that affect uranium Kd
values. Under oxidizing conditions at pH values greater than 6, their derived Kd values were
approximately 100 ml/g. At high concentrations of dissolved carbonate, and pH values greater
than 6, the Kd values for uranium decrease considerably. Their results indicate that the triple
layer surface complexation model using constants obtained under well controlled laboratory
conditions on well characterized minerals can easily be applied to estimate the dependence of
uranium adsorption and uranium Kd values as a function of a variety of important site
environmental conditions.
Efforts have also been made to compile site binding constants for radionuclides and other metals
to create "sorption databases" for use with geochemical codes such as MINTEQA2. For
example, Turner et al. (1993) and Turner (1993, 1995) describe the application of the surface-
complexation models (SCMs) [i.e., the diffuse layer model (DLM), constant capacitance model
(CCM), and triple layer model (TLM)] in the geochemical reaction code MINTEQA2 to
simulate potentiometric titration and adsorption data published for U(VI) and other radionuclides
on several single mineral phases. Their studies were conducted in support of developing a
uniform approach to using surface complexation models to predict radionuclide migration
J.24
-------
behavior associated with disposal of high-level radioactive waste in a geologic repository. The
parameter optimization code FITEQL was used for fitting and optimization of the adsorption
binding constants that were used in conjunction with MINTEQA2 and its thermodynamic
database. For those radionuclides having sufficient data, the surface-complexation models were
used to examine the effects of changing geochemical conditions (e.g., pH) on radionuclide
adsorption. Turner etal. (1993) and Turner (1993, 1995) include a detailed listing and
documentation of the adsorption reactions and associated binding constants used for the
MINTEQA2 DLM, CCM, and TLM calculations. Although all 3 models proved capable of
simulating the available adsorption data, the DLM was able to do so using the fewest parameters
(Turner, 1995). Compared to empirical approaches (e.g., Kd) for predicting contaminant
adsorption, Turner notes that surface complexation models based on geochemical principles have
the advantage of being used to extrapolate contaminant adsorption to environmental conditions
beyond the range measured experimentally.
J.5.0 Other Studies of Uranium
The following studies and adsorption reviews were identified during the course of this study.
Although they typically do not contain uranium Kd data, they discuss aspects of uranium
adsorption behavior in soils that might be useful to some readers searching for similar site
conditions. These studies and reviews are briefly discussed below.
Ames and Rai (1978) reviewed and evaluated the processes influencing the mobility and
retention of radionuclides. Their review for uranium discussed the following published
adsorption studies. The following descriptions are paraphrased from in their report.1
Dementyev and Syromyatnikov (1968) determined that the maximum adsorption
observed for uranium in the pH 6 region is due to the boundary between the dominant
uranium aqueous species being cationic and anionic at lower and higher pH values,
respectively.
Goldsztaub and Wey (1955) determined that 7.5 and 2.0 g uranium could be adsorbed per
100 g of calcined montmorillonite and kaolinite, respectively.
Horrath (1960) measured an average enrichment factor of 200 to 350 for the adsorption
of uranium on peat.
Kovalevskii (1967) determined that the uranium content of western Siberian
noncultivated soils increased as a function of their clay content and that clay soils
contained at least 3 times more uranium than sands.
1 The full citations listed for these references at the end of this appendix are provided exactly
as given by Ames and Rai (1978).
J.25
-------
Manskaya etal. (1956) studied adsorption of uranium on fulvic acids as a function of pH.
Results indicate a maximum removal of uranium of approximately 90 percent at pH 6,
and 30 percent removal at pH values of 4 and 7.
Masuda and Yamamoto (1971) showed that uranium from 1 to 100 mg/1 uranium
solutions was approximately completely adsorbed by volcanic ash, alluvial, and sandy
soils.
Rancon (1973) investigated the adsorption of uranium on several soils and single
minerals. The Kd values reported by Rancon (1973) are (in ml/g): 39 for river sediment
(quartz, clay, calcite, and organic matter); 33 for river peat; 16 for soil (quartz, clay,
calcite, and no organic matter); 270 for quartz-clay soil developed from an altered schist;
0 for quartz; 7 for calcite; and 139 for illite.
Ritchie et al. (1972) determined that the uranium content of a river sediment increased
with decreasing particle size.
Rozhkova et al. (1959) showed a maximum adsorption of uranium on lignite and humic
acids between pH 5 and 6.
Rubtsov (1972) found that approximately 58 percent of the total uranium was associated
with the
-------
accompanied by unreacted zeolite. The products of the reactions involving Na- and K-A
zeolites contained a phase similar to compreignacite (K2O6U03-11H20). Those experiments
conducted with Ca-A zeolite contained a phase similar to becquerelite ( CaO6U03-l 1H20).
Ho and coworkers studied the adsorption of U(VI) on a well-characterized, synthetic hematite
(a-Fe203) sol.1 Characterization data listed for the hematite sol by Ho and Doern (1985) and
cited in other studies by Ho and coworkers included a particle size of 0.12 |im, surface area of 34
m2/g, isoelectric point2 of pH 7.6, and composition of >98 percent a-Fe203 and <2 percent
P-FeO(OH). Ho and Doern (1985) studied the adsorption of U(VI) on the hematite sol as a
function of dissolved U(VI) concentration. Their procedure consisted of mixing 10 ml of the
hematite sol {i.e., constant particle concentration of 0.2 g/1) with 10 ml of uranyl nitrate solution.
The uranyl solutions and hematite sol were previously prepared at the required concentration,
pH, and ionic strength. The mixtures were equilibrated for 16 hr at 25°C. Over the pH range
from 3 to 6.2, Ho and Doern determined that adsorption of U(VI) on the hematite sol increased
with increasing concentrations of dissolved U(VI). Even though the particles of hematite sol had
a net positive charge in the pH range from 3 to 6.2, significant adsorption of U(VI) was
measured. The adsorption of U(VI) was greatest at pH of approximately 6.2, and decreased
significantly at lower pH values. Ho and Miller (1986) investigated the adsorption of U(VI)
from bicarbonate solutions as a function of initial U(VI) concentration over the pH range from
6.5 to 9.1 using the hematite sol described previously. Their experimental procedure was similar
to that described by Ho and Doern, except that the measurements were completed using a lxlO"3
mol/1 NaHC03 solution in which its pH was adjusted by the addition of dilute HC1. Over the pH
range from 6.5 to 9.1, Ho and Miller determined that the adsorption of uranium decreased
abruptly with increasing pH. In experiments conducted with an initial U(VI) concentration of
5xl0"6 mol/1, the reported percentages of U(VI) adsorbed on the hematite sol were approximately
98, 47, and 26 percent, respectively, at pH values of 7.1, 8.4, and 9.1. Ho and Miller (1985)
evaluated the effect of dissolved humic acid on the adsorption of U(VI) by the hematite sol
described in Ho and Hoern (1985) over the pH range from approximately 4.3 to 6.4. As used by
Ho and Miller, the term "humic acid" referred to the "fraction of humic substances soluble in
water at pH>4.30." The results of Ho and Miller (1985) indicate that the adsorption of U(VI) by
hematite is affected by the addition of humic acid and that the magnitude of this effect varies
with pH and concentration of humic acid. At low humic acid concentration of 3 mg/1, the
surface coverage of the hematite by the humic acid is low and the U(VI) adsorption by the
hematite sol is similar to that observed for bare hematite particles. However, as the
concentration of humic acid increases, the adsorption behavior of U(VI) changes. In the extreme
case of a high humic acid concentration of 24 mg/1, the U(VI) adsorption is opposite that
observed for bare hematite sol. At intermediate concentrations of humic acid, there is a change
1 A sol is defined as "a homogeneous suspension or dispersion of colloidal matter in a fluid"
(Bates and Jackson, 1980).
2 The isoelectric point (iep) is defined as "the pH where the particle is electrokinetically
uncharged" (Stumm and Morgan, 1981).
J.27
-------
from enhanced U(VI) adsorption at low pH to reduced adsorption at high pH for the pH range
from 4.3 to 6.4.
Tsunashima et al. (1981) investigated the sorption of U(VI) by Wyoming montmorillonite. The
experiments consisted of reacting, at room temperature, the <2-|im size fraction of
montmorillonite saturated with Na+, K+, Mg2+, Ca2+, and Ba2+ with U(VI) nitrate solutions
containing 1 to 300 ppm U(VI). The tests included systems with fixed volumes and variable
uranyl concentrations [50 mg of clay dispersed in 200 ml of U(VI) nitrate solutions with 1-40
ppm U(VI)] and systems with variable volumes and fixed amounts of U(VI) [100 mg clay
dispersed in 100 ml of solution]. The duration of the contact period for the clay-solution
suspensions was 5 days. Based on the conditions of the constant volume/constant ionic strength
experiments, the results indicated that adsorption of uranyl ions (U02 ) was strongly preferred
over Na+ and K+ by the clay, and less strongly preferred versus Mg2+, Ca2+, and Ba2+.
Vochten et al. (1990) investigated the adsorption of U(VI) hydrolytic complexes on well-
characterized samples of natural zeolites in relation to the double-layer potential of the minerals.
The zeolite samples included chabazite (CaAl2Si4012-6H20), heulandite
[(Ca,Na2)Al2Si7018-6H20], scolecite (CaAl2Si3O10-3H2O), and stilbite
[(Ca,Na2,K2)Al2Si7018-7H20], The adsorption measurements were conducted at 25°C over a pH
range from 4 to 7.5 using 0.1 g of powdered (35-75 |im) zeolite added to a 50 ml solution of
2xl0"5 mol/1 U(VI). The suspension was shaken for 1 week in a nitrogen atmosphere to avoid
the formation of U(VI) carbonate complexes. Given the relatively small dimension of the
channels in the zeolite crystal structure and ionic diameter of the non-hydrated U02+ ion (3.84
A), Vochten concluded that the adsorption of U(VI) was on the external surfaces of the zeolites.
The results indicate low adsorption of U(VI) to the 4 zeolites from pH 4 to 5. The amount of
U(VI) adsorption increases rapidly from pH 5 to 7 with the maximum rate of increase being
between pH 6 to 7.1 The adsorption results indicate that chabazite and scolecite had higher
sorptive capacities for U(VI) than heulandite and stilbite.
1 Based on experimental solubility [e.g., as Krupka et al. (1985) and others] and geochemical
modeling studies, the authors of this document suspect that Vochten et al. (1990) may have
exceeded the solubility of U(VI) above pH 5 and precipitated a U(VI) solid, such schoepite
(U03-2H20), during the course of their adsorption measurements conducted in the absence of (or
minimal) dissolved carbonate.
J.28
-------
Table J.5. Uranium Kd values selected from literature for development of look-up table.
PH
U Kd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
8.3
1.98
Hanford Groundwater
Trench 8 Loamy Sand
Kaplan and Serne (1995,
Part. Sat. Column, 40%)
8.3
0.49
Hanford Groundwater
Trench 8 Loamy Sand
Kaplan and Serne (1995,
Part. Sat. Column, 40%)
8.3
2.81
Hanford Groundwater
Trench 8 Loamy Sand
Kaplan and Serne (1995,
Part. Sat. Column, 38%)
8.3
0.62
Hanford Groundwater
Trench 8 Loamy Sand
Kaplan and Serne (1995,
Part. Sat. Column, 22%)
8.3
0.45
Hanford Groundwater
Trench 8 Loamy Sand
Kaplan and Serne (1995,
Part. Sat. Column, 30%)
8.3
0.54
Hanford Groundwater
Trench 8 Loamy Sand
Kaplan and Serne (1995,
Part. Sat. Column, 23%)
8.3
0.62
Hanford Groundwater
Trench 8 Loamy Sand
Kaplan and Serne (1995,
Part. Sat. Column, 25%)
8.3
0.40
Hanford Groundwater
Trench 8 Loamy Sand
Kaplan and Serne (1995,
Part. Sat. Column, 17%)
8.3
0.10
Hanford Groundwater
Trench 8 Loamy Sand
Kaplan and Serne (1995,
Part. Sat. Column, 7%)
8.3
0.08
Hanford Groundwater
Trench 8 Loamy Sand
Kaplan and Serne (1995,
Part. Sat. Column, 7%)
8.3
2.0
5.2
Hanford Groundwater
Trench 8 Loamy Sand
Lindenmeir et al. (1995,
Saturated Column 1)
8.3
0.5
5.2
Hanford Groundwater
Trench 8 Loamy Sand
Lindenmeir et al. (1995,
Saturated Column 1)
8.3
2.7
5.2
Hanford Groundwater
Trench 8 Loamy Sand
Lindenmeir et al. (1995,
Saturated Column 1)
8.3
1.0
5.2
Hanford Groundwater
Trench 8 Loamy Sand
Lindenmeir et al. (1995,
Unsat. Column 1, 65%)
8.3
0.5
5.2
Hanford Groundwater
Trench 8 Loamy Sand
Lindenmeir et al. (1995,
Unsat. UFA 1, 70%)
8.3
0.2
5.2
Hanford Groundwater
Trench 8 Loamy Sand
Lindenmeir et al. (1995,
Unsat. UFA 2, 24%)
8.3
1.1
5.2
Hanford Groundwater
Trench 8 Loamy Sand
Lindenmeir et al. (1995,
U nsat. Column 1, 63%)
8.3
1.1
5.2
Hanford Groundwater
Trench 8 Loamy Sand
Lindenmeir et al. (1995,
Unsat. Column 2, 43%)
8.3
0.6
5.2
Hanford Groundwater
Trench 8 Loamy Sand
Lindenmeir et al. (1995,
Unsat. UFA 1A, 29%)
8.3
0.6
5.2
Hanford Groundwater
Trench 8 Loamy Sand
Lindenmeir et al. (1995,
Unsat. UFA 1C, 29%)
J.29
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
8.4
0.20
5.3
6.3
Hanford Groundwater
Trench 94
Kaplan et al. (1998, Batch)
8.4
0.15
5.3
6.3
Hanford Groundwater
Trench 94
Kaplan et al. (1998, Batch)
8.4
0.09
5.3
6.3
Hanford Groundwater
Trench 94
Kaplan et al. (1998, Batch)
8.4
0.15
5.3
6.3
Hanford Groundwater
Trench 94
Kaplan et al. (1998, Batch)
8.4
0.14
5.3
6.3
Hanford Groundwater
Trench 94
Kaplan et al. (1998, Batch)
7.92
1.99
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1998, Batch)
8.05
1.92
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1998, Batch)
7.99
1.91
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1998, Batch)
7.99
2.10
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1998, Batch)
7.98
2.25
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1998, Batch)
7.97
2.44
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1998, Batch)
8.48
1.07
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1998, Batch)
8.26
1.46
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1998, Batch)
8.44
1.37
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1998, Batch)
9.12
2.12
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1998, Batch)
8.46
0.90
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1996, 100%
Unsaturated Batch)
8.46
1.70
5.3
6.3
Hanford Groundwater
Trench 94
Kaplan et al. (1996, 100%
Unsaturated Batch)
8.46
1.00
6.0
6.3
Hanford Groundwater
TSB-1
Kaplan et al. (1996, 100%
Unsaturated Batch)
8.46
1.10
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1996, Batch)
8.46
3.50
5.3
6.3
Hanford Groundwater
Trench 94
Kaplan et al. (1996, Batch)
8.46
2.10
6.0
6.3
Hanford Groundwater
TSB-1
Kaplan et al. (1996, Batch)
8.46
0.24
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1996)
8.46
0.64
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1996)
8.46
0.51
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1996)
8.46
0.46
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1996)
8.46
0.35
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1996)
8.46
0.53
6.4
14.8
Hanford Groundwater
Trench AE-3
Kaplan et al. (1996)
8.46
0.23
5.3
6.3
Hanford Groundwater
Trench 94
Kaplan et al. (1996)
8.46
0.15
5.3
6.3
Hanford Groundwater
Trench 94
Kaplan et al. (1996)
8.46
0.1
5.3
6.3
Hanford Groundwater
Trench 94
Kaplan et al. (1996)
8.46
0.16
5.3
6.3
Hanford Groundwater
Trench 94
Kaplan et al. (1996)
J.30
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
8.46
0.12
5.3
6.3
Hanford Groundwater
Trench 94
Kaplan et al. (1996)
2
8
Sand
Neiheisel [1983, as listed
in Thibault et al. (1990)]
1
7
Sand
Neiheisel [1983, as listed
in Thibault et al. (1990)]
3
15
Sand
Neiheisel [1983, as listed
in Thibault et al. (1990)]
750
36
Clayey Sand
Neiheisel [1983, as listed
in Thibault et al. (1990)]
770
21
Clayey Sand
Neiheisel [1983, as listed
in Thibault et al. (1990)]
550
19
Clayey Sand
Neiheisel [1983, as listed
in Thibault et al. (1990)]
2.00
100
Fine Sandstone and
Silty Sand
Haji-Djafari et al. [1981, as
listed in Thibault et al.
(1990)]
4.50
200
Fine Sandstone and
Silty Sand
Haji-Djafari et al. [1981, as
listed in Thibault et al.
(1990)]
5.75
1,000
Fine Sandstone and
Silty Sand
Haji-Djafari et al. [1981, as
listed in Thibault et al.
(1990)]
7.00
2,000
Fine Sandstone and
Silty Sand
Haji-Djafari et al. [1981, as
listed in Thibault et al.
(1990)]
5.6
25,000
Red-Brown Clayey
Seeley and Kelmers [1984, as
listed in Thibault et al.
(1990)]
5.6
250
Red-Brown Clayey
Seeley and Kelmers [1984, as
listed in Thibault et al.
(1990)]
5.20
58.4
Thibault et al. (1990, values
determined by coworkers)
5.10
294.9
Thibault et al. (1990, values
determined by coworkers)
5.20
160
Thibault et al. (1990, values
determined by coworkers)
6.20
45.4
Thibault et al. (1990, values
determined by coworkers)
J.31
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
7.00
450
36
28.0
Silty Loam Clay
Thibault et al. (1990, values
determined by coworkers)
7.30
1.2
15
17.0
Loam
Thibault et al. (1990, values
determined by coworkers)
4.90
0.03
2
5.8
Medium Sand
Thibault et al. (1990, values
determined by coworkers)
5.50
2900
1
120.0
Organic
Thibault et al. (1990, values
determined by coworkers)
7.40
1.9
10
9.1
Fine Sandy Loam
Thibault et al. (1990, values
determined by coworkers)
7.40
2.4
11
8.7
Fine Sandy Loam
Thibault et al. (1990, values
determined by coworkers)
6.60
590
10
10.8
Fine Sandy Loam
Thibault et al. (1990, values
determined by coworkers)
6.50
4500
10
12.6
Fine Sandy Loam
Thibault et al. (1990, values
determined by coworkers)
7.10
15
12
13.4
Fine Sandy Loam
Thibault et al. (1990, values
determined by coworkers)
7.00
16
Sand
Rancon [1973, as listed in
Thibault et al. (1990)]
7.00
33
Organic Peat
Rancon [1973, as listed in
Thibault et al. (1990)]
6.50
4400
Clay Fraction
Dahlman et al. [1976, as
listed in Thibault et al.
(1990)]
2.80
200
Abyssal Red Clay
Erickson (1980)
7.10
790,000
Abyssal Red Clay
Erickson (1980)
8.3
1.70
2.6
Hanford Groundwater
CGS-1 sand (coarse
gravel sand)
Serne et al. (1993, Batch)
8.3
2.30
5.2
Hanford Groundwater
Trench 8 Loamy Sand
(medium/coarse sand)
Serne et al. (1993, Batch)
8.3
79.30
6.0
Hanford Groundwater
TBS-1 Loamy Sand
(Touchet Bed sand)
Serne et al. (1993, Batch)
8.00
56.0
Hanford Groundwater,
GR-1
Umtanum Basalt
Salter et al. (1981)
8.00
7.5
Hanford Groundwater,
GR-1
Umtanum Basalt
Salter et al. (1981)
J.32
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
8.00
13.2
Hanford Groundwater,
GR-1
Umtanum Basalt
Salter et al. (1981)
8.00
17.8
Hanford Groundwater,
GR-1
Umtanum Basalt
Salter et al. (1981)
8.00
20.2
Hanford Groundwater,
GR-1
Umtanum Basalt
Salter et al. (1981)
8.00
13.0
Hanford Groundwater,
GR-1
Flow E Basalt
Salter et al. (1981)
8.00
2.7
Hanford Groundwater,
GR-1
Flow E Basalt
Salter et al. (1981)
8.00
2.2
Hanford Groundwater,
GR-1
Flow E Basalt
Salter et al. (1981)
8.00
3.2
Hanford Groundwater,
GR-1
Flow E Basalt
Salter et al. (1981)
8.00
2.9
Hanford Groundwater,
GR-1
Flow E Basalt
Salter et al. (1981)
8.00
16.0
Hanford Groundwater,GR-1
Pomona Basalt
Salter et al. (1981)
8.00
2.2
Hanford Groundwater,GR-1
Pomona Basalt
Salter et al. (1981)
8.00
3.5
Hanford Groundwater,GR-1
Pomona Basalt
Salter et al. (1981)
8.00
5.2
Hanford Groundwater,GR-1
Pomona Basalt
Salter et al. (1981)
8.00
5.8
Hanford Groundwater,GR-1
Pomona Basalt
Salter et al. (1981)
10.00
2.8
Hanford Groundwater,GR-2
Umtanum Basalt
Salter et al. (1981)
10.00
2.3
Hanford Groundwater,GR-2
Umtanum Basalt
Salter et al. (1981)
10.00
2.8
Hanford Groundwater,GR-2
Umtanum Basalt
Salter et al. (1981)
10.00
2.8
Hanford Groundwater,GR-2
Umtanum Basalt
Salter et al. (1981)
10.00
2.5
Hanford Groundwater,GR-2
Umtanum Basalt
Salter et al. (1981)
10.00
1.0
Hanford Groundwater,GR-2
Flow E Basalt
Salter et al. (1981)
10.00
0.5
Hanford Groundwater,GR-2
Flow E Basalt
Salter et al. (1981)
10.00
0.4
Hanford Groundwater,GR-2
Flow E Basalt
Salter et al. (1981)
10.00
0.8
Hanford Groundwater,GR-2
Flow E Basalt
Salter et al. (1981)
10.00
0.2
Hanford Groundwater,GR-2
Flow E Basalt
Salter et al. (1981)
10.00
0.9
Hanford Groundwater,GR-2
Pomona Basalt
Salter et al. (1981)
10.00
0.6
Hanford Groundwater,GR-2
Pomona Basalt
Salter et al. (1981)
10.00
0.8
Hanford Groundwater,GR-2
Pomona Basalt
Salter et al. (1981)
10.00
0.5
Hanford Groundwater,GR-2
Pomona Basalt
Salter et al. (1981)
10.00
0.4
Hanford Groundwater,GR-2
Pomona Basalt
Salter et al. (1981)
7.66
7.5
1.83
17.7
Hanford Groundwater,GR-1
Umtanum Basalt
Ames et al. (1982)
J.33
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
7.66
13
1.83
17.7
Hanford Groundwater,GR-1
Umtanum Basalt
Ames et al. (1982)
7.66
18
1.83
17.7
Hanford Groundwater,GR-1
Umtanum Basalt
Ames et al. (1982)
7.66
20
1.83
17.7
Hanford Groundwater,GR-1
Umtanum Basalt
Ames et al. (1982)
8.38
2.4
1.83
17.7
Hanford Groundwater,GR-2
Umtanum Basalt
Ames et al. (1982)
8.38
2.9
1.83
17.7
Hanford Groundwater,GR-2
Umtanum Basalt
Ames et al. (1982)
8.38
2.9
1.83
17.7
Hanford Groundwater,GR-2
Umtanum Basalt
Ames et al. (1982)
8.38
2.5
1.83
17.7
Hanford Groundwater,GR-2
Umtanum Basalt
Ames et al. (1982)
7.65
2.7
1.5
10.3
Hanford Groundwater,GR-1
Flow E Basalt
Ames et al. (1982)
7.65
2.2
1.5
10.3
Hanford Groundwater,GR-1
Flow E Basalt
Ames et al. (1982)
7.65
3.2
1.5
10.3
Hanford Groundwater,GR-1
Flow E Basalt
Ames et al. (1982)
7.65
2.9
1.5
10.3
Hanford Groundwater,GR-1
Flow E Basalt
Ames et al. (1982)
8.38
0.55
1.5
10.3
Hanford Groundwater,GR-2
Flow E Basalt
Ames et al. (1982)
8.38
0.38
1.5
10.3
Hanford Groundwater,GR-2
Flow E Basalt
Ames et al. (1982)
8.38
0.78
1.5
10.3
Hanford Groundwater,GR-2
Flow E Basalt
Ames et al. (1982)
8.38
0.19
1.5
10.3
Hanford Groundwater,GR-2
Flow E Basalt
Ames et al. (1982)
7.90
2.2
4.84
31.2
Hanford Groundwater,GR-1
Pomona Basalt
Ames et al. (1982)
7.90
3.5
4.84
31.2
Hanford Groundwater,GR-1
Pomona Basalt
Ames et al. (1982)
7.90
5.2
4.84
31.2
Hanford Groundwater,GR-1
Pomona Basalt
Ames et al. (1982)
7.90
5.8
4.84
31.2
Hanford Groundwater,GR-1
Pomona Basalt
Ames et al. (1982)
8.48
0.57
4.84
31.2
Hanford Groundwater,GR-2
Pomona Basalt
Ames et al. (1982)
8.48
0.83
4.84
31.2
Hanford Groundwater,GR-2
Pomona Basalt
Ames et al. (1982)
8.48
0.47
4.84
31.2
Hanford Groundwater,GR-2
Pomona Basalt
Ames et al. (1982)
8.48
0.42
4.84
31.2
Hanford Groundwater,GR-2
Pomona Basalt
Ames et al. (1982)
7.7
27
71.66
646
Hanford Groundwater,GR-1
Smectite, secondary
Ames et al. (1982)
7.7
39
4.84
31.2
Hanford Groundwater,GR-1
Smectite, secondary
Ames et al. (1982)
7.7
127
4.84
31.2
Hanford Groundwater,GR-1
Smectite, secondary
Ames et al. (1982)
7.7
76
4.84
31.2
Hanford Groundwater,GR-1
Smectite, secondary
Ames et al. (1982)
7.7
12
4.84
31.2
Hanford Groundwater,GR-2
Smectite, secondary
Ames et al. (1982)
7.7
42
4.84
31.2
Hanford Groundwater,GR-2
Smectite, secondary
Ames et al. (1982)
7.7
48
4.84
31.2
Hanford Groundwater,GR-2
Smectite, secondary
Ames et al. (1982)
7.7
22
4.84
31.2
Hanford Groundwater,GR-2
Smectite, secondary
Ames et al. (1982)
6.85
477,285
0.01 NaCl
Amor Fe(III)
Hydroxide
Ames et al. (1983c)
6.80
818,221
0.01 NaCl
Amor Fe(III)
Hydroxide
Ames et al. (1983c)
J.34
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
6.90
1,739,87
7
0.01 NaCl
Amor Fe(III)
Hydroxide
Ames et al. (1983c)
6.90
1,690,52
2
0.01 NaCl
Amor Fe(III)
Hydroxide
Ames et al. (1983c)
8.60
4,313
0.01 NaHCOj
Amor Fe(III)
Hydroxide
Ames et al. (1983c)
8.65
14,098
0.01 NaHCOj
Amor Fe(III)
Hydroxide
Ames et al. (1983c)
8.65
21,362
0.01 NaHCOj
Amor Fe(III)
Hydroxide
Ames et al. (1983c)
8.80
26,269
0.01 NaHCOj
Amor Fe(III)
Hydroxide
Ames et al. (1983c)
7.15
8.4
15.3
1.59
0.01 NaCl
Biotite
Ames et al. (1983b)
7.15
43.9
15.3
1.59
0.01 NaCl
Biotite
Ames et al. (1983b)
7.15
253.5
15.3
1.59
0.01 NaCl
Biotite
Ames et al. (1983b)
7.15
544.3
15.3
1.59
0.01 NaCl
Biotite
Ames et al. (1983b)
7.15
113.7
0.95
1.88
0.01 NaCl
Muscovite
Ames et al. (1983b)
7.15
251.0
0.95
1.88
0.01 NaCl
Muscovite
Ames et al. (1983b)
7.15
459.7
0.95
1.88
0.01 NaCl
Muscovite
Ames et al. (1983b)
7.15
68.2
0.95
1.88
0.01 NaCl
Muscovite
Ames et al. (1983b)
7.15
67.9
1.17
1.22
0.01 NaCl
Phlogopite
Ames et al. (1983b)
7.15
85.4
1.17
1.22
0.01 NaCl
Phlogopite
Ames et al. (1983b)
7.15
95.4
1.17
1.22
0.01 NaCl
Phlogopite
Ames et al. (1983b)
8.65
0.9
15.3
1.59
0.01 NaHCOj
Biotite
Ames et al. (1983b)
8.65
3.4
15.3
1.59
0.01 NaHCOj
Biotite
Ames et al. (1983b)
8.65
23.0
15.3
1.59
0.01 NaHCOj
Biotite
Ames et al. (1983b)
8.65
80.8
15.3
1.59
0.01 NaHCOj
Biotite
Ames et al. (1983b)
8.65
2.2
0.95
1.88
0.01 NaHCOj
Muscovite
Ames et al. (1983b)
8.65
26.9
0.95
1.88
0.01 NaHCOj
Muscovite
Ames et al. (1983b)
8.65
602.5
0.95
1.88
0.01 NaHCOj
Muscovite
Ames et al. (1983b)
8.65
3489.6
0.95
1.88
0.01 NaHCOj
Muscovite
Ames et al. (1983b)
8.65
0.6
1.17
1.22
0.01 NaHCOj
Phlogopite
Ames et al. (1983b)
8.65
1.1
1.17
1.22
0.01 NaHCOj
Phlogopite
Ames et al. (1983b)
8.65
0.6
1.17
1.22
0.01 NaHCOj
Phlogopite
Ames et al. (1983b)
7
544.5
25
116.1
0.01 NaCl
Illite, only lowest U
cone
Ames et al. (1983a)
J.35
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
8.5
90.5
25
116.1
0.01 NaHCOj
Illite, only lowest U
cone
Ames et al. (1983a)
7
657.8
12.2
68.3
0.01 NaCl
Kaolinite, only lowest
U cone
Ames et al. (1983a)
8.5
400.8
12.2
68.3
0.01 NaHCOj
Kaolinite, only lowest
U cone
Ames et al. (1983a)
7
542.0
120
747
0.01 NaCl
Montmorillonite, only
lowest U cone
Ames et al. (1983a)
8.5
1.8
120
747
0.01 NaHCOj
Montmorillonite, only
lowest U cone
Ames et al. (1983a)
7
299.9
95
861
0.01 NaCl
Nontronite, only
lowest U cone
Ames et al. (1983a)
8.5
4.1
95
861
0.01 NaHCOj
Nontronite, only
lowest U cone
Ames et al. (1983a)
7
138.0
16.03
137.3
0.01 NaCl
Glauconite, only
lowest U cone
Ames et al. (1983a)
8.5
114.2
16.03
137.3
0.01 NaHCOj
Glauconite, only
lowest U cone
Ames et al. (1983a)
7
66.5
140.2
20
0.01 NaCl
Clinoptilolite, only
lowest U cone
Ames et al. (1983a)
8.5
0.6
140.2
20
0.01 NaHCOj
Clinoptilolite, only
lowest U cone
Ames et al. (1983a)
7
225.7
3.18
46.8
0.01 NaCl
Opal, only lowest U
cone
Ames et al. (1983a)
8.5
1.7
3.18
46.8
0.01 NaHCOj
Opal, only lowest U
cone
Ames et al. (1983a)
7
300.5
2.79
626.3
0.01 NaCl
Silica Gel,, only
lowest U cone
Ames et al. (1983a)
8.5
639.9
2.79
626.3
0.01 NaHCOj
Silica Gel,, only
lowest U cone
Ames et al. (1983a)
7.3
4200.0
4.36
Spesutie (silt loam)
Erikson et al. (1993)
6.2
136.0
1.29
Transonic (silt loam)
Erikson et al. (1993)
8.0
44
9.30
Yuma (sandy loam)
Erikson et al. (1993)
6.8
4360
4.36
Spesutie (silt loam)
Erikson et al. (1993)
5.6
328
1.29
Transonic (silt loam)
Erikson et al. (1993)
8.0
54
9.30
Yuma (sandy loam)
Erikson et al. (1993)
39
River Sediment
(Quartz, clay, calcite,
organic matter)
Rancon (1973) as cited
by Ames and Rai (1978)
33
River Peat
Rancon (1973) as cited
by Ames and Rai (1978)
J.36
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
16
River Sediment
(Quartz, clay, calcite)
Rancon (1973) as cited
by Ames and Rai (1978)
270
Soil (Quartz and Clay,
from Altered Schist)
Rancon (1973) as cited
by Ames and Rai (1978)
0
Quartz
Rancon (1973) as cited
by Ames and Rai (1978)
7
Calcite
Rancon (1973) as cited
by Ames and Rai (1978)
139
Illite
Rancon (1973) as cited
by Ames and Rai (1978)
27
(0.8-
332)
Fresh Water
Gorleben Salt Dome,
Sandy Sediment
Warnecke etal. (1984, 1986,
1994), Warnecke and Hild
(1988)
1
(0.3-1.6)
Fresh Water
Gorleben Salt Dome,
Sandy Sediment
Warnecke etal. (1984, 1986,
1994), Warnecke and Hild
(1988)
17
(8.5-
100)
Saline Water
Gorleben Salt Dome,
Clayish Sediment
Warnecke etal. (1984, 1986,
1994), Warnecke and Hild
(1988)
14-1,400
Saline Water
Gorleben Salt Dome,
Clayish Sediment
Warnecke etal. (1984, 1986,
1994), Warnecke and Hild
(1988)
4
Quaternary fresh water
Former Konrad Iron
Ore Mine
Warnecke etal. (1986),
Warnecke and Hild (1988)
6
Turonian fresh water
Former Konrad Iron
Ore Mine
Warnecke etal. (1986),
Warnecke and Hild (1988)
6
Cenomanian saline water
Former Konrad Iron
Ore Mine
Warnecke etal. (1986),
Warnecke and Hild (1988)
20
Albian (Hauterivain) saline
water
Former Konrad Iron
Ore Mine
Warnecke etal. (1986),
Warnecke and Hild (1988)
1.4
Albian (Hils) saline water
Former Konrad Iron
Ore Mine
Warnecke etal. (1986),
Warnecke and Hild (1988)
2.6
Kimmeridgian saline water
Former Konrad Iron
Ore Mine
Warnecke etal. (1986),
Warnecke and Hild (1988)
3
Oxfordian saline water
Former Konrad Iron
Ore Mine
Warnecke etal. (1986),
Warnecke and Hild (1988)
3
Bajocian (Dogger) saline
water
Former Konrad Iron
Ore Mine
Warnecke etal. (1986),
Warnecke and Hild (1988)
3.83
310
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
3.90
235
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
J.37
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
3.94
741
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
3.96
211
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
4.03
694
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
4.13
720
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
4.28
898
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
4.33
630
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
4.36
247
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
4.53
264
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
4.58
903
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
4.61
324
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
4.71
522
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
4.81
1,216
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
4.95
1,185
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
4.84
3,381
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.00
2,561
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.10
2,635
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.11
3,807
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.19
4,293
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.52
4,483
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.15
4,574
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.24
5,745
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.16
7,423
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
J.38
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
5.28
3,214
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.52
5,564
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.44
6,687
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.54
6,185
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.58
6,615
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.85
7,124
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.45
8,146
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.56
8,506
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.74
9,332
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.50
10,462
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.69
10,681
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.54
11,770
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.66
13,616
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.81
14,675
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.86
14,417
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.75
20,628
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
6.01
24,082
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
6.20
22,471
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
5.95
26,354
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
6.35
26,078
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
6.40
25,601
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
6.35
27,671
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
J.39
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
6.46
30,529
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
6.13
31,477
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
6.26
33,305
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
6.80
37,129
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
6.86
37,657
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
6.81
32,312
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
7.10
29,390
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
7.85
33,583
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
7.67
26,518
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
8.40
30,523
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
8.51
19,632
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
9.45
23,177
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
9.80
17,763
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
9.90
14,499
Synthetic Groundwater,
function of pH
Kaolinite
Giblin (1980)
3.8
2
Synthetic Groundwater,
function of pH
Quartz
Andersson etal. (1982)
3.5
5
Synthetic Groundwater,
function of pH
Quartz
Andersson etal. (1982)
3.7
8
Synthetic Groundwater,
function of pH
Quartz
Andersson etal. (1982)
3.7
69
Synthetic Groundwater,
function of pH
Quartz
Andersson etal. (1982)
4.0
116
Synthetic Groundwater,
function of pH
Quartz
Andersson etal. (1982)
6.4
1,216
Synthetic Groundwater,
function of pH
Quartz
Andersson etal. (1982)
6.5
1,824
Synthetic Groundwater,
function of pH
Quartz
Andersson etal. (1982)
J.40
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
6.6
2,679
Synthetic Groundwater,
function of pH
Quartz
Andersson etal. (1982)
7.7
7,379
Synthetic Groundwater,
function of pH
Quartz
Andersson etal. (1982)
8.0
2,506
Synthetic Groundwater,
function of pH
Quartz
Andersson etal. (1982)
8.3
21,979
Synthetic Groundwater,
function of pH
Quartz
Andersson etal. (1982)
8.6
3,999
Synthetic Groundwater,
function of pH
Quartz
Andersson etal. (1982)
9.0
14,689
Synthetic Groundwater,
function of pH
Quartz
Andersson etal. (1982)
3.4
27
Synthetic Groundwater,
function of pH
Biotite
Andersson etal. (1982)
4.4
326
Synthetic Groundwater,
function of pH
Biotite
Andersson etal. (1982)
4.4
522
Synthetic Groundwater,
function of pH
Biotite
Andersson etal. (1982)
4.7
418
Synthetic Groundwater,
function of pH
Biotite
Andersson etal. (1982)
5.1
1,489
Synthetic Groundwater,
function of pH
Biotite
Andersson etal. (1982)
5.2
2,512
Synthetic Groundwater,
function of pH
Biotite
Andersson etal. (1982)
6.4
2,812
Synthetic Groundwater,
function of pH
Biotite
Andersson etal. (1982)
7.3
7,228
Synthetic Groundwater,
function of pH
Biotite
Andersson etal. (1982)
7.3
16,634
Synthetic Groundwater,
function of pH
Biotite
Andersson etal. (1982)
7.4
9,840
Synthetic Groundwater,
function of pH
Biotite
Andersson etal. (1982)
8.1
4,732
Synthetic Groundwater,
function of pH
Biotite
Andersson etal. (1982)
9.0
8,337
Synthetic Groundwater,
function of pH
Biotite
Andersson etal. (1982)
3.3
207
Synthetic Groundwater,
function of pH
Apatite
Andersson etal. (1982)
3.8
324
Synthetic Groundwater,
function of pH
Apatite
Andersson etal. (1982)
4.0
726
Synthetic Groundwater,
function of pH
Apatite
Andersson etal. (1982)
J.41
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
4.0
668
Synthetic Groundwater,
function of pH
Apatite
Andersson etal. (1982)
4.4
3,767
Synthetic Groundwater,
function of pH
Apatite
Andersson etal. (1982)
4.5
4,732
Synthetic Groundwater,
function of pH
Apatite
Andersson etal. (1982)
5.0
16,218
Synthetic Groundwater,
function of pH
Apatite
Andersson etal. (1982)
5.3
8,241
Synthetic Groundwater,
function of pH
Apatite
Andersson etal. (1982)
6.0
140,605
Synthetic Groundwater,
function of pH
Apatite
Andersson etal. (1982)
7.7
24,660
Synthetic Groundwater,
function of pH
Apatite
Andersson etal. (1982)
3.6
460
Synthetic Groundwater,
function of pH
Attapulgite
(Palygorskite)
Andersson etal. (1982)
4.1
1,514
Synthetic Groundwater,
function of pH
Attapulgite
(Palygorskite)
Andersson etal. (1982)
4.2
7,194
Synthetic Groundwater,
function of pH
Attapulgite
(Palygorskite)
Andersson etal. (1982)
4.5
6,471
Synthetic Groundwater,
function of pH
Attapulgite
(Palygorskite)
Andersson etal. (1982)
4.7
4,753
Synthetic Groundwater,
function of pH
Attapulgite
(Palygorskite)
Andersson etal. (1982)
5.1
23,335
Synthetic Groundwater,
function of pH
Attapulgite
(Palygorskite)
Andersson etal. (1982)
5.9
12,531
Synthetic Groundwater,
function of pH
Attapulgite
(Palygorskite)
Andersson etal. (1982)
6.4
266,686
Synthetic Groundwater,
function of pH
Attapulgite
(Palygorskite)
Andersson etal. (1982)
7.3
645,654
Synthetic Groundwater,
function of pH
Attapulgite
(Palygorskite)
Andersson etal. (1982)
7.8
82,224
Synthetic Groundwater,
function of pH
Attapulgite
(Palygorskite)
Andersson etal. (1982)
8.7
46,132
Synthetic Groundwater,
function of pH
Attapulgite
(Palygorskite)
Andersson etal. (1982)
3.2
1,175
Synthetic Groundwater,
function of pH
Montimorillonite
Andersson etal. (1982)
4.4
12,503
Synthetic Groundwater,
function of pH
Montimorillonite
Andersson etal. (1982)
6.6
3,917
Synthetic Groundwater,
function of pH
Montimorillonite
Andersson etal. (1982)
J.42
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
7.0
10,139
Synthetic Groundwater,
function of pH
Montimorillonite
Andersson etal. (1982)
7.0
28,054
Synthetic Groundwater,
function of pH
Montimorillonite
Andersson etal. (1982)
7.3
10,715
Synthetic Groundwater,
function of pH
Montimorillonite
Andersson etal. (1982)
8.2
21,528
Synthetic Groundwater,
function of pH
Montimorillonite
Andersson etal. (1982)
8.4
20,370
Synthetic Groundwater,
function of pH
Montimorillonite
Andersson etal. (1982)
9.0
18,621
Synthetic Groundwater,
function of pH
Montimorillonite
Andersson etal. (1982)
5.1
7,391
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
5.0
1,177
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
5.1
2,180
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
5.4
3,680
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
5.3
4,437
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
5.5
7,265
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
5.5
7,108
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
5.8
23,603
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
5.8
22,948
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
4.7
176
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
4.8
176
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
5.0
283
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
5.0
297
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
5.4
708
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
5.7
1,961
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
J.43
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
5.6
2,367
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
5.9
4,283
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
5.9
4,936
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
6.0
7,936
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
6.1
8,586
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
6.2
17,631
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
6.3
19,553
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
6.4
30,963
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
6.5
43,756
45
99
Ca Electrolyte, C02 Free
Kenoma Clay, <2um
fraction
Zachara et al. (1992, Fig 6)
5.1
508
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.2
554
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.2
676
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.4
874
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.4
1,136
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.6
1,136
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.7
2,143
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.8
2,363
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.9
9,829
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.9
11,966
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
6.0
33,266
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
6.1
37,596
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
J.44
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
4.8
377
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
4.8
399
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.1
620
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.0
637
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.5
1,476
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.5
1,603
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.8
3,091
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
6.1
6,047
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
6.1
5,823
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
6.3
13,713
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
6.4
13,341
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
4.9
918
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.1
1,168
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.1
1,251
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.6
2,719
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
5.7
2,928
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
6.7
14,848
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
6.8
13,036
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
7.0
13,827
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
7.0
18,042
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
7.0
19,150
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
7.1
21,771
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
J.45
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
7.1
18,097
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
7.4
26,008
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
7.4
19,488
59
112
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
7.7
31,032
Ca Electrolyte, C02 Free
Ringold Clay Isolate,
<2um Fraction
Zachara et al. (1992, Fig 7)
6.28
3,400
Reducing Conditions
PCE Surface Core, 0-8
cm
Sheppard and Thibault
(1988, In Situ)
6.28
2,800
Reducing Conditions
PCE Surface Core,
9-16 cm
Sheppard and Thibault
(1988, In Situ)
6.28
3,000
Reducing Conditions
PCE Surface Core,
17-24 cm
Sheppard and Thibault
(1988, In Situ)
6.28
11,600
Reducing Conditions
PCE Surface Core,
25-32 cm
Sheppard and Thibault
(1988, In Situ)
6.28
18,600
Reducing Conditions
PCE Surface Core,
33-40 cm
Sheppard and Thibault
(1988, In Situ)
6.09
3,200
Reducing Conditions
PCE Deep Core, 9-16
cm
Sheppard and Thibault
(1988, In Situ)
6.09
8,900
Reducing Conditions
PCE Deep Core, 17-24
cm
Sheppard and Thibault
(1988, In Situ)
6.09
9,400
Reducing Conditions
PCE Deep Core, 25-32
cm
Sheppard and Thibault
(1988, In Situ)
6.09
12,500
Reducing Conditions
PCE Deep Core, 33-40
cm
Sheppard and Thibault
(1988, In Situ)
5.94
3,000
Reducing Conditions
SCE Surface Core, 0-5
cm
Sheppard and Thibault
(1988, In Situ)
6.82
8,800
Reducing Conditions
SCE Surface Core,
6-20 cm
Sheppard and Thibault
(1988, In Situ)
7.28
2,600
Reducing Conditions
SCE Surface Core,
21-25 cm
Sheppard and Thibault
(1988, In Situ)
7.28
1,700
Reducing Conditions
SCE Surface Core,
26-30 cm
Sheppard and Thibault
(1988, In Situ)
7.28
700
Reducing Conditions
SCE Surface Core,
31-40 cm
Sheppard and Thibault
(1988, In Situ)
1,300
Reducing Conditions
PCE Surface Core,
0-40 cm
Sheppard and Thibault
(1988, Batch)
2,100
Reducing Conditions
PCE Deep Core, 40-80
cm
Sheppard and Thibault
(1988, Batch)
2,000
Reducing Conditions
SCE Surface Core,
1-10 cm
Sheppard and Thibault
(1988, Batch)
2,900
Reducing Conditions
SCE Surface Core,
10-30 cm
Sheppard and Thibault
(1988, Batch)
J.46
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
870
Reducing Conditions
SCE Surface Core,
30-40 cm
Sheppard and Thibault
(1988, Batch)
5.7
46
2.3
Site Borehole Groundwater
Clay (Glacial Till,
Less Than 5 mm)
Bell and Bates (1988)
5.7
46
3.0
Site Borehole Groundwater
Cl:2 (Brown, Slightly
Silty, Less Than 5
mm)
Bell and Bates (1988)
5.7
900
2.7
Site Borehole Groundwater
C3 (Dark Brown
Coarse Granular
Deposit, Less Than 5
mm)
Bell and Bates (1988)
5.7
2,200
2.9
Site Borehole Groundwater
C6 (Brown Coarse
Granular Deposit,
Less Than 5 mm)
Bell and Bates (1988)
5.7
560
0.8
Site Borehole Groundwater
Sand (Light Brown
Coarse Granular
Deposit, Less Than 5
mm)
Bell and Bates (1988)
4.16
85.0
0.5
1.11
A12
Serkiz and Johnson (1994)
4.99
170.0
3.3
1.82
A13
Serkiz and Johnson (1994)
3.42
5.3
3
3.74
A13R
Serkiz and Johnson (1994)
3.19
2.1
1.5
1.39
A22
Serkiz and Johnson (1994)
3.01
1.7
4.5
1.4
A23
Serkiz and Johnson (1994)
3.19
3.7
4.4
7.92
A31
Serkiz and Johnson (1994)
3.5
1.4
3.1
1
A3 2
Serkiz and Johnson (1994)
3.29
1.2
4.7
2.1
A42
Serkiz and Johnson (1994)
5.42
2,200.0
2.5
0.68
A52
Serkiz and Johnson (1994)
3.72
2.3
2
0.42
A53
Serkiz and Johnson (1994)
3.24
2.7
2.8
4.71
B13
Serkiz and Johnson (1994)
3.93
8.5
3.9
3.06
B14
Serkiz and Johnson (1994)
3.86
10.1
4.9
B23
Serkiz and Johnson (1994)
4.02
5.2
2.5
3.8
B23R
Serkiz and Johnson (1994)
3.83
14.0
7.5
5.69
B24
Serkiz and Johnson (1994)
4.62
390.0
6.2
2.5
B32
Serkiz and Johnson (1994)
4.64
180.0
5.5
8.42
B33
Serkiz and Johnson (1994)
4.67
190.0
12.6
21.4
B42
Serkiz and Johnson (1994)
3.66
6.4
1.2
3.02
B43
Serkiz and Johnson (1994)
4.09
39.0
8.2
15.1
B51
Serkiz and Johnson (1994)
J.47
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
3.61
5.3
B52
Serkiz and Johnson (1994)
4.69
530.0
3.3
2.39
B52R
Serkiz and Johnson (1994)
3.68
6.4
C13
Serkiz and Johnson (1994)
3.75
23.0
6.4
C14
Serkiz and Johnson (1994)
3.96
30.0
1.28
C22
Serkiz and Johnson (1994)
4.17
980.0
6.4
6.12
C23
Serkiz and Johnson (1994)
5.53
3,600.0
5.5
2.54
C32
Serkiz and Johnson (1994)
4.64
6,300.0
6.1
8.54
C33
Serkiz and Johnson (1994)
5.27
14,000.0
7.9
11.4
C42
Serkiz and Johnson (1994)
4.51
13,000.0
3
5.04
C43
Serkiz and Johnson (1994)
6.78
11,000.0
5.3
1.96
D13
Serkiz and Johnson (1994)
4.14
13.0
D13RA
Serkiz and Johnson (1994)
9.3
2
2.55
D13RB
Serkiz and Johnson (1994)
4
320.0
10.5
11.4
E13
Serkiz and Johnson (1994)
4.04
310.0
4.5
8.5
E14
Serkiz and Johnson (1994)
5.85
2,700.0
6.4
15.5
E23
Serkiz and Johnson (1994)
4.32
980.0
3.9
13.3
E23R
Serkiz and Johnson (1994)
3.87
290.0
7.3
13.8
E24
Serkiz and Johnson (1994)
4.27
1,500.0
6.5
11.5
E33
Serkiz and Johnson (1994)
4.05
380.0
3.7
10.5
E34
Serkiz and Johnson (1994)
5.27
16,000.0
31.8
20.6
E41
Serkiz and Johnson (1994)
4.87
18,000.0
14.5
20.6
E42
Serkiz and Johnson (1994)
4.3
7,500.0
15.5
16.1
F12
Serkiz and Johnson (1994)
4.9
830.0
8.51
F13
Serkiz and Johnson (1994)
4.69
160.0
8.1
7.48
F22
Serkiz and Johnson (1994)
6.48
16,000.0
13
11.6
F23
Serkiz and Johnson (1994)
4.85
8,700.0
14.2
15.1
F32
Serkiz and Johnson (1994)
4.77
2,900.0
18.3
13.6
F33
Serkiz and Johnson (1994)
5.2
34,000.0
17.2
11.8
F42
Serkiz and Johnson (1994)
4.12
330.0
14.2
F43
Serkiz and Johnson (1994)
5.91
5,500.0
42.2
19.9
F52
Serkiz and Johnson (1994)
5.63
27,000.0
16.3
13.3
F53
Serkiz and Johnson (1994)
4.16
139.0
0.5
1.11
A12
Serkiz and Johnson (1994)
4.99
361.0
3.3
1.82
A13
Serkiz and Johnson (1994)
3.42
9.46
3
3.74
A13R
Serkiz and Johnson (1994)
J.48
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
3.19
3.79
1.5
1.39
A22
Serkiz and Johnson (1994)
3.01
1.55
4.5
1.4
A23
Serkiz and Johnson (1994)
3.19
4.43
4.4
7.92
A31
Serkiz and Johnson (1994)
3.5
1.38
3.1
1
A3 2
Serkiz and Johnson (1994)
3.29
1.19
4.7
2.1
A42
Serkiz and Johnson (1994)
5.42
160.0
2.5
0.68
A52
Serkiz and Johnson (1994)
3.72
16.0
2
0.42
A53
Serkiz and Johnson (1994)
3.24
2.0
2.8
4.71
B13
Serkiz and Johnson (1994)
3.93
10.4
3.9
3.06
B14
Serkiz and Johnson (1994)
3.86
10.7
4.9
B23
Serkiz and Johnson (1994)
4.02
4.0
2.5
3.8
B23R
Serkiz and Johnson (1994)
3.83
11.3
7.5
5.69
B24
Serkiz and Johnson (1994)
4.62
332.0
6.2
2.5
B32
Serkiz and Johnson (1994)
4.64
212.0
5.5
8.42
B33
Serkiz and Johnson (1994)
4.67
180.0
12.6
21.4
B42
Serkiz and Johnson (1994)
3.66
7.1
1.2
3.02
B43
Serkiz and Johnson (1994)
4.09
20.8
8.2
15.1
B51
Serkiz and Johnson (1994)
3.61
2.6
B52
Serkiz and Johnson (1994)
4.69
180.0
3.3
2.39
B52R
Serkiz and Johnson (1994)
3.68
5.6
C13
Serkiz and Johnson (1994)
3.75
28.3
6.4
C14
Serkiz and Johnson (1994)
3.96
27.4
1.28
C22
Serkiz and Johnson (1994)
4.17
823.0
6.4
6.12
C23
Serkiz and Johnson (1994)
5.53
540.0
5.5
2.54
C32
Serkiz and Johnson (1994)
4.64
690.0
6.1
8.54
C33
Serkiz and Johnson (1994)
5.27
1,400.0
7.9
11.4
C42
Serkiz and Johnson (1994)
4.51
460.0
3
5.04
C43
Serkiz and Johnson (1994)
6.78
690.0
5.3
1.96
D13
Serkiz and Johnson (1994)
4.14
26.6
D13RA
Serkiz and Johnson (1994)
22.6
2
2.55
D13RB
Serkiz and Johnson (1994)
4
650.0
10.5
11.4
E13
Serkiz and Johnson (1994)
4.04
190.0
4.5
8.5
E14
Serkiz and Johnson (1994)
4.32
310.0
3.9
13.3
E23R
Serkiz and Johnson (1994)
3.87
360.0
7.3
13.8
E24
Serkiz and Johnson (1994)
4.27
470.0
6.5
11.5
E33
Serkiz and Johnson (1994)
J.49
-------
PH
UKd
(rnl/g)
Clay
Cont.
(wt.%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Soil Identification
Reference / Comments
4.05
270.0
3.7
10.5
E34
Serkiz and Johnson (1994)
5.27
870.0
31.8
20.6
E41
Serkiz and Johnson (1994)
4.87
630.0
14.5
20.6
E42
Serkiz and Johnson (1994)
4.3
690.0
15.5
16.1
F12
Serkiz and Johnson (1994)
4.9
2,200.0
8.51
F13
Serkiz and Johnson (1994)
4.69
1,200.0
8.1
7.48
F22
Serkiz and Johnson (1994)
6.48
950.0
13
11.6
F23
Serkiz and Johnson (1994)
4.85
660.0
14.2
15.1
F32
Serkiz and Johnson (1994)
4.77
220.0
18.3
13.6
F33
Serkiz and Johnson (1994)
5.2
910.0
17.2
11.8
F42
Serkiz and Johnson (1994)
4.12
700.0
14.2
F43
Serkiz and Johnson (1994)
5.91
600.0
42.2
19.9
F52
Serkiz and Johnson (1994)
5.63
960.0
16.3
13.3
F53
Serkiz and Johnson (1994)
J.50
-------
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