United States
Environmental Protection
Agency
Office of
Air and Radiation
66O2J
EPA402-R-99-004B
August 1999
Understanding Variation
In Partition Coefficient,
KJ, Values
Volume II: Review Of
eochemistry And Available
Values For Cadmium,
Cesium, Chromium, Lead,
Plutonium, Radon, Strontium,
Thorium, Tritium (3H),
And Uranium
Case I: Kd = 1 mi/g
Continuous Source of Contamination
f y y y
Steady State
Flow
Case II: Kd =iOmi/g
Continuous Source of Contamination
C/q, =0.1
Steady State
Flow
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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.
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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 n 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.
m
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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. P4ge/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
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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
401 M Street, SW (6602J)
Washington, DC 20460
Values
or
webmaster.oria@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 defensibiliry 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 (^H), 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
NOTICE ii
FOREWORD iii
ACKNOWLEDGMENTS v
FUTURE UPDATES vi
ABSTRACT vii
LIST OF FIGURES xiii
LIST OF TABLES xv
1.0 Introduction 1.1
2.0 TheKdModel 2.1
3.0 Methods, Issues, and Criteria for Measuring Kd Values 3.1
3.1 Laboratory Batch Methods 3.1
3.2 Laboratpry 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
vm
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5.2.6.2.5 Limits of Kd Values with Respect to Dissolved Calcium,
Magnesium, and Sulfide Concentrations, andRedox 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
IX
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Concentrations Extractable Iron Content 5.34
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
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5.9.1 Overview: Important Aqueous- and Solid-Phase Parameters
Controlling Retardation 5.53
5.9.2 General Geochemistry 5.54
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
XI
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Appendix A - Acronyms and Abbreviations A. 1
AppendixB - Definitions B.I
Appendix C - Partition Coefficients for Cadmium C.I
Appendix D - Partition Coefficients for Cesium D. 1
Appendix E - Partition Coefficients for Chromium E.I
Appendix F - Partition Coefficients for Lead F.I
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
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LIST OF FIGURES
'age
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 ug/1 total dissolved U(VT) 5.70
Figure 5.6b. Calculated distribution of U(VI) hydrolytic species as a function of pH
at 1,000 ug/1 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.I. Relation between cadmium Kd values and pH in soils C.5
Figure D.I. 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.I6
xni
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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.I. 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 (K^) 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.l. 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 ^U 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 ^U and ^U 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 2XU and ^U 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
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.I. 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.I. 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. 11
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
xvi
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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 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 aiqueous 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.I. Summary of Kd values for Cr(VT) adsorption on soils. E.5
Table E.2. Data from Rai et al. (1988) for the adsorption of Cr(VT) 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. 11
Table E.5. Data from Rai et al. (1988) on effects of competing anions on Cr(VT)
adsorption on Kenoma soil E.12
Table F.I. 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.I. Plutonium adsorption data for soil samples G.10
Table G.2. Regression models for plutonium adsorption G.I3
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
xvn
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Table H.3. Simple and multiple regression analysis results involving
strontium Kd values, CEC (meq/100 g), pH, and clay content (percent) H.8
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 I.I. 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.I. Uranium Kd values (ml/g) listed by Warnecke et al. (1994, Table 1) 112
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.I 7
xvm
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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,1>2 values and their use in formulating the retardation factor, Rf. The contaminant retardation
factor (Rf) is the parameter commonly used in transport models to describe the chemical
interaction between the contaminant and geological materials (i.e., soil, sediments, rocks, and
geological formations, henceforth simply referred to as soils3). It includes processes such as
surface adsorption, absorption into the soil structure, precipitation, and physical filtration of
colloids. Specifically, it describes the rate of contaminant transport relative to that of
groundwater. This report is provided for technical staff from EPA and other organizations who
are responsible for prioritizing site remediation and waste management decisions. The
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 Rf, (2) the difference between the original thermodynamic Kd
parameter derived from the ion-exchange literature and its "empiricized" use in contaminant
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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transport codes, and (3) the explicit and implicit assumptions underlying the use of the Kd
parameter in contaminant transport codes.
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
Krf 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 K^ 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 H 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
defensibiliry issues concerning available Kd data and the measurement of Kd values is also
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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 MTNTEQA2 code is also 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(VT) (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)-aniomc 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
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exist, or where steep pH or oxidation-reduction (redox) gradients exist. Adsorption/desorption
will likely be the key process controlling contaminant migration in areas where chemical steady
state exist, such as in areas far from the point source. Diffusion flux spreads solute via a
concentration gradient (i.e., Pick's law). Diffusion is a dominant transport mechanism when
advection is insignificant, and is usually a negligible transport mechanism when water is being
advected in response to various forces.
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2.0 The 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 adsprbate adsorbed per
mass of solid to the amount of the adsorbate remaining in solution at equilibrium. For the
reaction
C = A,
(2.1)
the mass action expression for Kd is
„ _ Mass of Adsorbate Sorbed _ ^
d Mass of Adsorbate in Solution Cj
where A = free or unoccupied surface adsorption sites
Cj = total dissolved adsorbate remaining in solution at equilibrium
Ai = amount of adsorbate on the solid at equilibrium.
The Kd is typically given in units of ml/g. Describing the Kd in terms of this simple reaction
assumes that A is in great excess with respect to Q and that the activity of AJ is equal to 1.
Chemical retardation, Rf, is defined as,
(2.2)
(2.3)
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
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retardation factor 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
p - 1 4- b "K"
j\_ —• 1 *r —— JS.J
where pb = porous media bulk density (mass/length3)
n,. = effective porosity of the media at saturation.
(2.4)
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
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track many more parameters and some numerical solving techniques (such as closed-form
analytical solutions) can no longer be used to perform the integration necessary to solve for the Kd
value and/or retardation factor, Rf. Generally, computer codes that can accommodate the
parametric-Kd model use a chemical subroutine to update the Kd value used to determine the RF,
when called 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 con-
centration, 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 K^ 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 K00 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 experiments can be completed quickly for a wide variety of elements and
chemical environments. The primary disadvantage of the batch technique for measuring Kd is that
it does not necessarily reproduce the chemical reaction conditions that take place in the real
environment. For instance, in a soil column, water passes through at a finite rate and both
reaction time and degree of mixing between water and soil can be much less than those occurring
in a laboratory batch test. Consequently, Kd values from batch experiments can be high relative to
the extent of sorption occurring in a real system, and thus result in an estimate of contaminant
retardation that is too large. Another disadvantage of batch experiments is that they do not
accurately simulate desorption of the radionuclides or contaminants from a contaminated soil or
solid waste source. The Kd values are frequently used with the assumption that adsorption and
desorption reactions are reversible. This assumption is contrary to most experimental
observations that show that the desorption process is appreciably slower than the adsorption
process, a phenomenon referred to as hysteresis. The rate of desorption may even go to zero, yet
a significant mass of the contaminant remains sorbed on the soil. Thus, use of Kd values
determined from batch adsorption tests in contaminant transport models is generally considered to
provide estimates of contaminant remobilization (release) from soil that are too large (i.e.,
estimates of contaminant retention that are too low).
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
3.1
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of interest. The retardation factor (a ratio of the velocity of the contaminant to that of water) is
measured directly from the experimental data. A Kd value can be calculated from the retardation
factor. It is frequently useful to compare the back-calculated Kd value from these experiments
with those derived directly from the batch experiments to evaluate the influence of limited
interaction between solid and solution imposed by the flow-through system.
One potential advantage of the flow-through column studies is that the retardation factor can be
inserted directly into the transport code. However, if the study site contains different hydrological
conditions (e.g., porosity and bulk density) than the column experiment, 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 man 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
K^ 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
3.2
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containing solid, gaseous, and liquid phases. Each phase contains several different constituents.
The use of simplified systems containing single mineral phases and aqueous phases with 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 chemi-
cally 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 (i.e., 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 K.J 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
3.3
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(e.g., soil). 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 ground-
water 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 Kj 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
O.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
TWf
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 MINEQL code with
the thermodynamic database and geochemical attributes of the WATEQ3 code. The MINTEQA2
code includes submodels to calculate aqueous speciation/complexation, oxidation-reduction, gas-
phase equilibria, solubility and saturation state (i.e., saturation index), precipitation/dissolution of
solid phases, and adsorption. The most current version of MINTEQA2 available from EPA is
compiled to execute on a personal computer (PC) using the MS-DOS computer operating system.
The MINTEQA2 software package includes PRODEFA2, a computer code used to create and
modify input files for MINTEQA2.
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(H), Cr(in), and Cr(VI); cadmium; lead; strontium; and uranium,
including the valence states U(Tfl), U(IV), U(V), and U(VT). 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 K"*, 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-
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complexation constants for the contaminants of interest and the assumed sorptive phase. The
EPA (EPA 1992, 1996) has used the MINTEQA2 model and this approach to estimate Kd values
for several metals under a variety of geochemical conditions and metal concentrations to support
several waste disposal issues. The EPA in its "Soil Screening Guidance" determined
MINTEQA2-estimated Kd values for barium, beryllium, cadmium, Cr(ffl), 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 IQ Values
The important geochemical factors affecting the sorptiori1 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 dissolved gases and complexing ligands,
such as carbonate, may be less in a groundwater as a result of infiltration of surface water through
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
-------
the soil column. Additionally, the redox potential of groundwaters, especially deep ground-waters,
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 (Le., 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
Silica, as H4SiO4
Ca
Mg
Na
K
Inorganic Carbon, as CO3
S04
ci
F
NO3
PO.
Total Concentration
mg/1
20.8
15
4.1
6.3
2.3
57
11
7.8
1
1
0.0767
mol/I
2.16 xW4
3.7 x 1C"4
1.7 xlO"4
2.7 x 10-4
5.9 xlO'5
9.5 x 10-4
1.1 x 10-4
2.2 x 10"4
5 x 10'5
2 x 10'5
8.08 x lO'7
1 Most values from this table were taken, from Hem (1 985: 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 (1 985, p. 128) lists 25 (jg/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 (PO4) listed above.
5.3
-------
Table 5.2. Concentrations of contaminants used in the aqueous
species distribution calculations.
Element
Cd
Cs
Cr
Pb
Pu
Rn
Sr
Th
3H
U
Total Cone.
(ug/0
1.0
-
1.4
1.0
3.2 x 10-'
—
110
1.0
—
0.1 and
1,000
Reference for Concentration of Contaminant
Used in Aqueous Spetiation Calculations
Hem (1985, p. 142) lists this value as a median concentration of dissolved
cadmium based on the reconnaissance study of Duram et al. (1971) of metal
concentrations in surface waters in the United States.
Distribution of aqueous species was not modeled, because mobility of dissolved
cesium is not significantly affected by complexation (see Section 5.3).
Hem (1985, p. 138) lists this value as an average concentration estimated by
Kharkar et al. (1968) for chromium in river waters.
Hem (1985, p. 144) lists this value as an average concentration estimated by
Duram et al. (1971) for lead in surface-water samples from north- and southeastern
sections of the United States.
This concentration is based on the maximum activity of 239-240Pu measured by
Simpson et al. (1984) in 33 water samples taken from the highly alkaline Mono
Lake in California.
Aqueous speciation was not calculated, because radon migrates as a dissolved gas
and is not affected by complexation (see Section 5.7).
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).
Hem (1985, p. 150) gives 0.01 to 1 ug/1 as Hie range expected for thorium
concentrations hi fresh waters.
Aqueous speciation was not calculated, because tritium (3H) migrates as tritiated
water.
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 ug/1) of total dissolved
uranium were used to model the species distributions. Hem (1985, p. 148) gives
0. 1 to 10 ug/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 ug/1.
5.4
-------
5.2 Cadmium Geochemistry andKd 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-S M SO2.") is the uncomplexed Cd2+ species. The
dominant cadmium solution species in groundwater at pH values greater than 8.2 are CdCOg (aq)
and to a smaller extent CdCl+. Both precipitation/coprecipitation/dissolution and
adsorption/desorption reactions control cadmium concentrations. Several researchers report that
otavite (CdCO3) limits cadmium solution concentrations in alkaline soils. The solid Cd3(PO4)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
-------
5.2.3 Aqueous Speciation
Cadmium forms soluble complexes with inorganic and organic Hgands 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 ug/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 et al. (1971) of
metal concentrations in surface waters in the United States. These MENTEQA2 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 [CdHCOj 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 CdCOa (aq). The species
CdSOS (aq), CdHCOj, CdCf, and CdOET 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 ug/1.
Table 5.3. Cadmium aqueous species included
in the speciation calculations.
Aqueous Species
Cd2
CdOH+, Cd(OH)° (aq), Cd(OH);,
Cd2OH3+
CdHCO*, CdCO| (aq), Cd(CO3#
(aq), Cd(S04
CdCl+, CdCl°2 (aq), CdCl;, CdOHCl0 (aq)
CdF+, CdFj (aq)
5.6
-------
o
I
€
I
100
90
SO
70
60
50
40
30
20
10
0
CdCI
CdC03 (aq)
CdSO4°(aq)
CdHCO3
67
pH
10
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 ug/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 et 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.
5.7
-------
5.2.4 Dissolution/Predpitation/Copredpitafion
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 (10'7M), indicating that cadmium at these concentrations is not
likely to precipitate. At pH levels greater than 7.5, the solubilities of Cd3(PO4)2 or CdCO3 may
control the concentrations of cadmium in soils. Cavallaro and McBride (1978) and McBride
(1980) demonstrated that otavite, CdCO3, 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 et 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 CdCO3. 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)CO3. 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 et al., 1977). Seme (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 O2 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 et 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.
5.8
-------
A third mechanism involves pyrite that may be present in soils_pr sediments and gets oxidized
when exposed to air.1 The pyrite oxidizes to form FeSO4., 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 etal, 1976; McBride etal, 1981;
Navrot et al., 1978; Petruzelli et at., 1978). During cation exchange, cadmium generally
exchanges with adsorbed calcium and magnesium (McBride et 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 CdOtT 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 et 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, 1980). In contrast, zinc almost completely displaces
1 D. M. Kargbo (1998, personal communication).
5.9
-------
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 et
a/., 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
et a/., 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 ofKd 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 et 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 um.
5.10
-------
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 oxidi2ing conditions.]
KH (ml/g)
Minimum
Maximum
pH
3-5
1
130
5-8
8
4,000
8-10
50
12,600
5.11
-------
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.
5.12
-------
5.3 Cesium Geochemistry andKd 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.
5.5.2 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 (tVl) = 2.05 y], 135Cs (tVi = 3 x 106 y), and 137Cs (t,/2 = 30.23 y) are at significant
concentrations 10 y after separation from nuclear fuels (Schneider and Platt, 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 et al, 1970)>
Ce > Fe > Mn > Co ^ Ru ;> Sr > Cs.
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.
5.13
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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)25H2O]} 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 et 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.
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
5.14
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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 et 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.
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
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 et al. (1999), and Oscarson and Hume
(1998) were identified and may be of interest to the reader.
5.15
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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 KJ 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.
5.16
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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)
Minimum
Maximum
CEC (meq/100 g) / day Content (wt.%)
<3/<4
10
3,500
3-10/4-20
30
9,000
10 - 50 / 20 -
60
80
26,700
Table 5.6. Estimated range of Kd values (ml/g) for cesium based
on CEC or clay content for sysitems 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.]
KH (ml/g)
Minimum
Maximum
CEC (meq/100 g) / Clay Content (wt.%)
<3/<4
30
9,000
3-10/4-20
70
22,000
10-50/20-
60
210
66,700
5.17
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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 andKd 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(in) because the former is less likely to adsorb or precipitate to the solid phase.
Chromium(VT) is also appreciably more toxic than Cr(ni). It exhibits significant subsurface
mobility in neutral and basic pH environments. In acid environments, Cr(VT) may be moderately
adsorbed by pH-dependent charge minerals, such as iron- and aluminum-oxide minerals. The
reduction of Cr(VT) to Cr(m) by ferrous iron, organic matter, and microbes is generally quite
rapid whereas the oxidation of Cr(in) to Cr(VT) 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(VT) 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(in)
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 ore processing, plating operations, and
5.18
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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 Puls (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(HI) 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(in) exists predominantly as Cr3* below pH 3.5 in a Cr(m)-H2O system.
With increasing pH, hydrolysis of Cr* yields CrOH2*, Cr(OH)^, Cr(OH)3(aq), and Cr(OH);,
Cr2(OH)f", and Cr3(OH)4+. At higher chromium concentrations, polynuclear species, such as
Cr2(OH)2+ and Cr3(OH)f, can form slowly at 25°C (Baes and Mesmer, 1976). Chromium(VI)
hydrolyses extensively, forming primarily anionic species. These species are HCrO4 (bichromate),
CrO4" (chromate), and Cr-jOy" (dichromate) (Baes and Mesmer, 1976; Palmer and Wittbrodt,
1991; Richard and Bourg, 1991). Palmer and Puls (1994) presented some Cr(VI) speciation
diagrams representative of groundwater conditions. They showed that above pH 6.5, CrO4"
generally dominates. Below pH 6.5, HCrO; dominates when the total concentration of dissolved
Cr(VI) is low (<30 mM). When Cr(VT) concentrations are greater than 30 mM, Cr2O2' is the
dominant aqueous species relative to HCrO4 at acidic conditions (Palmer and Puls, 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(ni) in soils. Rai and Zachara (1984) concluded that most Cr(HT) solubility-
controlling solids in nature are either Cr(OH)3 or Cr(IH) coprecipitated with iron oxides. Then-
conclusion was supported by 3 observations: (1) the thermodynamic treatment of the data where
the solubility of chromite (FeCr2O4) 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(]H) ionic radii, and
(3) the observations that aqueous Cr(in) 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 Cr2O3. Because Cr(in) minerals are sparingly soluble, the
aqueous concentration of Cr(in) 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 Puls,
1994).
Several Cr(VI)-containing mineral phases may be present at chromium-contaminated sites.
Palmer and Wittbrodt (1990) identified PbCrO4 (crocoite), PbCrO4-H2O (iranite), andK2CrO4
5.19
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(tarapacaite) in chromium sludge from a plating facility. They also reported that BaCrO4 formed
a complete solid solution with BaSO4. 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(H), ferrous iron minerals, reduced sulfur, microbes, and organic matter
(Bartlett and Kimble, 1976; Nakayama et al, 1981). Studies indicate that Cr(VT) can be reduced
to Cr(ni) by ferrous iron derived from magnetite (Fe3O4) and ilmenite (FeTiO3) (White and
Hochella, 1989), hematite (FeA) (Eary and Rai, 1989),1 and pyrite (FeS2) (Blowes and Ptacek,
1992).
The reduction of Cr(VT) by Fe(H) is very rapid. The reaction can go to completion in a matter of
minutes (Eary and Rai, 1989). The rate of reduction of Cr(VT) increases with decreasing pH and
increasing initial Cr(VI) and reductant concentrations (Palmer and Puls, 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(ni) can precipitate with Fe(in) 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(IH) 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 Puls, 1994). Only 2 constituents in the environment
are known to oxidize Cr(m): dissolved oxygen and manganese-dioxide minerals [e.g., pyrolusite
(P-MnOz)]. Eary and Rai (1987) reported that the rate of Cr(IH) oxidation was much greater in
the presence of manganese-dioxide minerals than dissolved oxygen.
5.4.5 Sorption/Desorption
The extent to which Cr(TII) sorbs to soils is appreciably greater than that of Cr(VI) because the
former exists in groundwater as a cation, primarily as Cr3* (and its complexed species), whereas
the latter exists as an anion, primarily as CrO2; or HCrO;. Most information on Cr(VT) adsorption
comes from studies with pure mineral phases (Davis and Leckie, 1980; Griffin et al., 1977; Leckie
et al, 1980). These studies suggest that Cr(VI) adsorbs strongly to gibbsite (cc-Al2O3) and
amorphous iron oxide [Fe2O3-H2O(am)] at low to medium pH values (pH 2 to 7) and adsorbs
weakly to silica (SiO^) at all but very low pH values (Davis and Leckie, 1980; Griffin et al., 1977;
Leckie et al, 1980). These results can be explained by considering the isoelectric points (EEP)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 below the isoelectric point, the mineral has a net positive
1 Eary and Rai (1989) attributed the reduction of Cr(VI) to Cr(IH) by hematite (Fe2O3) as
containing having trace quantities of Fe(IT).
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.
5.20
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charge. Hence, anion adsorption generally increases as the pH becomes progressively lower than
the isoelectric point. The isoelectric point of gibbsite (a-Al2O3) is 9.1, amorphous iron oxide
[Fe2CVH2O (am)] is 8. 1, and silica is 2.0 (Stumm and Morgan, 198 1).
The presence of competing and, less commonly, complexing ions may significantly alter chromate
adsorption. Although sulfate is adsorbed less strongly on Fe2O3-H2O(am) than chromate, sulfate
may compete for adsorption sites when present in higher concentration (Leckie et 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
(am) (Leckie etal, 1980).
Adsorption of Cr(ni) to soils has received only a nominal amount of research attention. The
reason for this may be that sorption of Cr(HI) by soil is commonly ascribed to solid phase
formation. Chromium(in) 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(HI), like other +3 cationic metals, is
strongly and specifically absorbed by soil iron and manganese oxides (Korte et aL, 1976).
However, when Cr(ni) is present in solution at high concentrations, it may undergo exchange
reactions with aluminosilicates (Griffin et al, 1977). Chromium(HI) 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 et 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(HI) in soil solutions are typically controlled by
dissolution/precipitation reactions.
Increasing pH decreases adsorption (decrease in K^ 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(VT) which results in precipitation (higher
5.21
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Soils containing Mn oxides oxidize Cr(m) into Cr(VT) form thus resulting in lower
K,j 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 et al.
(1980), Davis and Leckie (1980), Griffin et al. (1977), and Rai et 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 et al., 1976). However, these investigators did not publish
either Kj 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 et
al. (1988) on 4 different soils confirmed that Kd values decrease with increasing equilibrium
Cr(VI) concentration. 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 the Cr(HI) oxidation state. The reduced Cr
has been shown to coprecipitate with ferric hydroxide. Therefore, observed removal of Cr(VT)
from solution when contacted with chromium-reductive soils may stem from both adsorption and
precipitation reactions. Similarly, Rai et al. (1988) also showed that certain soils containing
manganese oxides may oxidize Cr(HT) to Cr(VI). Depending on solution concentrations, the
oxidized form (+6) of chromium may also precipitate in the form of Ba(S,Cr)O4 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 SOj", COf/HCOg, HPOf, H2PO;
NOa and Cl", significantly reduce Cr(VT) adsorption on oxide minerals and soils (Leckie et al.,
1980; MacNaughton, 1977; Rai etaL, 1986; Rai etal, 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(VT), the results indicated that organic matter did not influence Cr(VT) adsorption
properties (see Appendix E).
5.4.6.2 KdLook-Up Tables
Among all available data for Cr(VI) adsorption on soils, the most extensive data set was
developed by Rai et 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
5.22
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competing anions such as SO^", and COfYHCOj. 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(VT) 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 et al. (1988) indicate the soils with DCB
extractable iron contents above -0.3 mmol/g can reduce Cr(VI) to Cr(Tfl). Therefore the
measured K^ 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(VT). 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 (iO.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 et 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 2x 10"3M(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).
5.23
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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"s 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 [PbCO3(aq), and Pb(CO3)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 PbCO3, Pb3(OH)2(CO3)25 PbSO4,
Pbs(PO4)3(Cl), and Pb4SO4(CO3)2(OH)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(i_x)PbxSO4), apatite [Ca(1_x)Pbx(PO4)3OH], calcite (Ca^^CCy, 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 ug/kg (Ter Haar et al, 1967). In contaminated soils,
lead concentrations may be as high as 18 percent by weight (Mattigod and Page, 1983; Ruby et
al, 1994). Lead in nature occurs in 4 stable isotopic forms (?*Pb, 206Pb, 207Pb, and 208Pb). The
isotopes, 206Pb, 207Pb, and 20SPb are the stable end products of the **U, S9U, and 232Th thorium
decay series, respectively (Robbins, 1980). Additionally., heavier isotopes of lead (*10Pb, 211Pb,
212Pb, and 214Pb) are known to occur in nature as intermediate products of uranium and thorium
decay (Robbins, 1978). The
5.25
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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).
5.5.5 Aqueous Speciation
Lead exhibits typical amphoteric1 metal ion behavior by forming hydrolytic species (Baes and
Mesmer, 1976). Formation of monomeric hydrolytic species, such as PbOH+, Pb(OH)2(aq) and
Pb(OH);, is well established. Although several polymeric hydrolytic species such as Pb2OH3+,
Pb3(OH)l+, Pb4(OH)4+, and Pb6(OH)f 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"9 M)
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 ug/1 based on the data for natural waters tabulated by Duram
et 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 PbHCOs (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(CO3)i'. 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 pg/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 on dissolution/precipitation of lead-solid phases and the adsorption/desorption
1 Amphoteric behavior is the ability of an aqueous complex or solid material to have a negative,
neutral, or positive charge.
5.26
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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 et 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 (PbCO3), hydrocerussite [Pb3(OH)2(CO3)2], anglesite (PbSO4), or
chloropyromorphite [Pb5(PO4)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 mat the mineral
leadhillite [Pb4SO4(CO3)2(OH)2] may be the solubility controlling solid for lead in a mine-waste
contaminated soil.
5.27
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Table 5.8. Lead aqueous species included in the
speciation calculations.
Aqueous Species
2+
Pb
PbOFT, Pb(OH)^(aq), Pb(OH);, Pb(OH#
Pb2(OH)3+,Pb3(OH)
PbCO°(aq),
PbSO^aq), Pb(S04)|-
PbNO3
PbCl+, PbClS(aq), PbCl^, PbCl|-
PbF+, PbF^(aq), PbFg, PbF"'
5.28
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I
•**
jfi
.3
100
80
60
40
20
10
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 ug/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)PbxSO4], apatite [Ca^jPbxOPO^OH], calcite [Ca(1_x)PbxCO3],
and iron sulfides, [Fe(1_x)PbxSj (Driesens, 1986; Goldschmidt, 1954; Nriagu and Moore, 1984;
Rickard and Nriagu, 1978). Consequently, the equilibrium solubility of lead controlled by these
phases will be less than the concentrations controlled by corresponding pure phases, namely
PbSO4, Pbs(PO4)3OH, PbCO3, and PbS, respectively.
5.29
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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 et 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.
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
5.30
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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.
• 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.
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.
5.31
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Adsorption of lead increases with increasing organic matter content of soils.
Increasing equilibrium solution concentrations correlates with decreasing lead adsorption
(decrease in
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 1 1), lead
adsorption increases with increasing pH (Braids et al, 1972; Bittel and Miller, 1974; Griffin and
Shimp, 1976; Haji-Djafari etal, 1981; Hildebrand and Blum, 1974; Overstreet and Krishamurthy,
1950; Scradato 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.
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
5.32
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effect on lead adsorption have been established by Genitse et al. (1982) and Soldatini et al.
(1976).
Lead adsorption by a subsurface soil sample from Hanford, Washington was investigated by
Rhoads et 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).
5.5.5.2 KdLook-Up Tables
Among all available data, Gerritse et al (1982) obtained aidsorption 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 et 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 et al. (1982) and Rhoads et 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 (ug/I)
0.1-0.9
1.0-9.9
10 - 99.9
100 - 200
Kd (rnl/g)
Minimum
Maximum
Minimum
Maximum
Minimum
Maximum
Minimum
Maximum
Soil pH
4.0 - 6.3
940
8,650
420
4,000
190
1,850
150
860
6.4 - 8.7
4,360
23,270
1,950
10,760
900
4,970
710
2,300
8.8 - 11.0
11,520
44,580
5,160
20,620
2,380
9,530
1,880
4,410
5.33
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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/l) were selected based
on the experimental data generated by Gerritse et al. (1982) and Rhoads et 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 andKd 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(VT) are common, whereas, under reducing conditions, Pu(ni) and Pu(IV) 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 PuO2 xH2O) 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(n) 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
Plutonium is produced by fissioning uranium fuel and is used hi 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, ^Pu [t,/2 (half life) =
5.34
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86 y], y39Pu (tt/z = 24,400 y), 240Pu (^ = 6,580 y), 241Pu (t,A = 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(in)
species would be dominant up to pH values approaching about 8.5, beyond which the Pu(IV)
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,
/., 1980b).
Of the contaminated sites considered in EPA/DOE/NRC (1993), radioactive contamination by
238pUj 23<>pUj ancj/or 240^ 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
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(IV) > Pu(IH) > Pu(VI) > Pu(V)
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++ 2H2O = 2Pu3++ PuOf+4H+
= Pu3+ + 2PuOf +2H2O.
5.35
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(Choppin, 1983). Plutonium hydrolytic species may have up to 4 coordinated hydroxyls.
The tendency ofplutonium 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, PuQV) exhibits the highest ionic potential and therefore forms
the strongest complexes with various ligands. Based on the equilibrium constants (K° 29g) for the
plutonium complexation reactions, ligands, such as chloride and nitrate, form weak complexes
(log K°>29g of 1 to 2) with plutonium, whereas fluoride, sulfate, phosphate, citrate, and oxalate
form stronger complexes (log K°29g of 6 to 30). Among the strongest complexes ofplutonium
are the hydroxy-carbonate mixed ligand complexes [e.g., Pu(OH)2(CO3)2"] (Tait et al, 1995;
Yamaguchi et al, 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 ofplutonium may
be the dominant soluble species in natural environments at lower pH (below 5 to 6) values (Allard
andRydberg, 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
(1.36xlO'ls M) total dissolved plutonium. This concentration is based on the maximum activity of
239*240Pu 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 pHTEh 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
The pe is related to Eh by the equation
Eh = 1303*1 pe
(5.2)
where R = universal gas constant (1.9872 cal/niol-K)
1 The electron activity is defined as unity for the standard hydrogen electrode.
5.36
-------
T = temperature in degrees kelvin
F = Faraday constant (96,487 coulombs/equivalent).
At25.0°C(298K),
Eh (mV) = 59.2 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 MINTEQA2 calculations of plutonium aqueous speciation. The plutonium aqueous species
that were included in the computation scheme are tabulated in Table 5.10. Thefmodynamic 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, PuF2+ and PuO2 are the dominant species of plutonium. The free
ionic species, PuO2 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 PuO2, and Pu(OH)2(CO3)2~,
with minor species being the neutral hydrolytic species Pu(OH)°4(aq) and the phosphate complex
Pu(HPO4)4". At pH values exceeding 6.5, the bulk of the dissolved plutonium (-90 percent)
would be comprised of the Pu(OH)2(CO3)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(OH)2(CO3)|~.
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 PuO2 xH2O. Many observations show that
plutonium associated with soils and paniculate organic matter is present in tetravalent oxidation
state (Nelson andLovett, 1980; Nelson etal, 1987; Silver, 1983). Calculations by Allard and
Rydberg (1983) based on available thermodynamic data show that, under reducing conditions, the
solubility of dissolved plutonium would be limited by the solid phase PuO2 at pH values greater
than 8, and by the solid phase Pu2(CO3)3 of trivalent plutonium at lower pH values.
5.37
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Table 5.10. Plutonium aqueous species included in the speciation calculations.
Redox
State
Aqueous Species
Pu(in)
Pu3+, PuOH2+,
PuC03+, Pu(C03)2,
PuSOj, Pu(SO4>2
PuH2POf, PuCP
Pu(IV)
3, Pu(OH):(aq)
Pu4+, PuOH3+,
Pu(OH)4(C03)t, Pu(OH)2(C03)l-
PuSOf , Pu(S04)°(aq), PuHPOf, Pu(HPO4)^(aq),
Pu(HP04)f , Pu(HP04)t
PuCl3+, PuF3+, PuFf,
Pu(V)
^, Pu02OH°(aq), (PuO2)2OH+
Pu(VI)
PuOf'PuO2OBT, PuO2(OH)^(aq),
Pu02(OH)3, (Pu
Pu02CO|(aq), PuO2(CO3)i-, PuO2(CO3)t
Pu02Cl+, Pu02F+, Pu02F°(aq), PuO2F3, PuO2Ff
Pu02SO^(aq), PuO2H2PO;
5.38
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100
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 et al.
(1994) indicated that a freshly precipitated amorphous PuO2 xH2O phase controls the equilibrium
solubility of plutonium. Solubility on aged precipitates by Rai et al. (1980a) and Delegard (1987)
also showed that equilibrium plutonium concentrations would be controlled by a partially
crystallized PuO2 xH2O phase at concentrations about 2 orders of magnitude less than that of
amorphous PuO2 xH2O. Therefore, under oxidizing conditions, amorphous PuO2 xH2O, 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, PuO2 xH2O may control plutonium
concentrations at about 1CT10 M (Rai et al, 1980a).
5.39
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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;
Coughtreye/a/., 1985; Thibaulte/a/., 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 et 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 et 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.
5.40
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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 et 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. Efowever, 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 (HI, 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 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
1 Since the completion of our review and analysis of Kd data for the selected contaminants and
radionuclides, the studies by Duffer al. (1999) and Fisher et al. (1999) were identified and may be
of interest to the reader.
5.41
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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 etal, 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 etal., 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 et 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(IV) adsorption on goethite have been quantified by Sanchez et al. (1985). They found
that carbonate concentrations exceeding 100 mmol/1 significantly reduced adsorption of Pu(IV)
on goethite. In contrast, under soil saturation extract conditions in which carbonate
concentrations typically range from 0.1 to 6 mmol/1 HCO3', Pu(TV) adsorption appears to increase
with increasing carbonate concentration (Glover et 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 6.8xlO"7 to
IxlO"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 etal, 1975), attapulgite (Billon, 1982), and
oxides, hydroxides, and oxyhydroxides (Evans, 1956; CharyuluetaL, 1991; Sanchez et al, 1985;
Tamura, 1972; Ticknor, 1993; Van Dalen etal, 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 et 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(IV) 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 et al (1976) conducted a set of experiments using 10"s M initial
concentration to study the adsorption behavior of Pu(IV) on 14 different soil samples from
5.42
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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 KdLook-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 um.
Table 5.11. Estimated range of Kd values for plutonium as a function of the
soluble carbonate and soil clay content values.
K* (ml/g)
Minimum
Maximum
Clay Content (wt%)
0-30
Soluble Carbonate
(meq/1)
0.1-2
5
420
3-4
80
470
5-6
130
520
31 - 50
Soluble Carbonate
(meq/1)
0.1-2
380
1,560
3-4
1,440
2,130
5-6
2,010
2,700
51-70
Soluble Carbonate
(meq/1)
0.1-2
620
1,980
3-4
1,860
2,550
5-6
2,440
3,130
5.43
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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 HCO3"). 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 et 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 et al. (1985). The data by Glover et 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 andKd 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 hi
water, and the hydrostatic pressure on ground water below the water table is sufficient to keep
dissolved radon in solution.
j
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 hi a total environmental system, including air transport pathways.
5.44
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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 (tv) 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 Superfund 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 ^U decay series
• 220Rn (ty=54 sec) by alpha decay from ^Ra in the 232Th decay series
• 219Rn ft,/=3.9 sec) by alpha decay from 223Ra in the ^U 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., Melson et
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 (A/2=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 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.
5.45
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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 et al. (1990), do not list any Kd values for radon. A Kd value of
zero should be considered for radon.
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 (SrCO3) control strontium concentrations in solutions.
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 ^Rn from
5.46
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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 CaCO3 and/or precipitation as SrCO3 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"1" 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 et a/., 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 (f,/2) = 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 (t,/2 = 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
There is little tendency for strontium to form complexes with inorganic ligands (Faure and Powell,
1972). The solubility of the free Sr2'1' 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 avoids 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
5.47
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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 Sr*+ 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, SrCO^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 ug/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 (SrSO4) and strontianite (SrCO3), have calcium
counterparts, anhydrite (CaSO4), and calcite (CaCO3). 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.
Table 5.12. Strontium aqueous species included in the
speciation calculations.
Aqueous Species
Sr*, SrOfT
SrCO^aq), SrSOftaq), SrNOj
SrCf, SrF+
SrPO;, SrHPO^aq), SrH2PO;, SrP2Of
5.48
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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)SO4 in more-alkaline environments (Ainsworth
andRai, 1987; Felmy etal., 1993).
5.&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 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 etal, 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 adsorbed 90 Sr as readily as
similar soils containing recently adsorbed strontium, indicating that 90Sr does not become more
recalcitrant to leaching with time. Furthermore, these studies suggested that cation exchange, and
not (co)precipitation, was responsible for 90Sr 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-free 90Sr was not easily
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) +
= 90SrX(s) + Ca2+(aq)
where X designates an exchange surface site. Adsorption phenomena are discussed in more detail
in Volume I of this report.
5.49
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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 of ^Sr present in these sediments was extracted by warm ITVNaCl 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, 1978), pH (Juo and Barber, 1970; Prout, 1958; Rhodes, 1957), and solution ionic
strength (Rhodes, 1957; Routson et 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 NaNO3. 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 NaNO3
(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
5.50
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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 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-
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 et al. (1999), Oscarson and Hume (1998),
and Wang et al. (1998) were identified and may be of interest to the reader.
5.51
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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
KJ 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
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.
5.52
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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)
Minimum
Maximum
CEC (meq/100 g) / Clay Content (wt.%)
3 /<4
pH
<5
1
40
5-8
2
60
8-10
3
120
3-10/4-20
pH
<5
10
150
5-8
15
200
8-10
20
300
10 - 50 / 20 - 60
pH
<5
100
1,500
5-8
200
1,600
8-10
300
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 decrease 90Sr Kd values due to formation of the weaker
strontium-carbonate aqueous complex.
5.9 Thorium Geochemistry andKd 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 carbonate and hydroxyl-carbonate complexes, such as
Th(OH)3CO3, at pH values greater than 7.5. Species distributions calculated using the stability
constants for thorium citrate, oxalate, and ethylenediamine complexes indicate that thorium
5.53
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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:
. 238
!U decay series: 234Th [fa (half life) = 24.1 d) and ^Th (fa = 8.0 x 104 y)
5Th decay series: ^Th (fa = 1.41 x 1010 y) and ^Th (fa = 1.913 y)
MSU decay series: ^Th (fa = 25.5 h) and 227Th (fa = 18.5 d).
Natural thorium consists of essentially 1 isotope, ^Th, with trace quantities of the other isotopes.
Thorium is fertile nuclear material in that the principal isotope ^Th can be converted by capture
of a thermal neutron and 2 beta decays to fissionable ^U which does not exist in nature. The
application of thorium as a reactor fuel in the ThO2 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, 1976; Cotton and Wilkinson, 1980). The thorium content in natural water is
very low. The concentration range in natural fresh water rarely exceeds 1 ug/1 (0.1 pCi/1232Th),
5.54
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although mg/1 concentrations of ^Th 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,ete.)SiO4] and thorianite (crystalline ThO2). 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)PO4] and
zircon (ZrSiO4). 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 et a/., 1967; Frondel, 1958).
Concentrations of thorium can be several weight percent in these deposits.
Because of their long half lives, 228Th (t,/z = 1.913 y), ^Th (t,/z = 8.0 x 104 y), and ^Th (t% =
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 ^Th 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 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.
1 A detrital mineral is defined as "any mineral grain resulting from mechanical disintegration of
parent rock" (Bates and Jackson, 1980).
5.55
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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 (og/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 ng/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)f, 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)f, Th4(OH)g+, and Th6(OH)i5 based on
the review of Bases and Mesmer (1976).
Table 5.14. Thorium aqueous species included in the
speciation calculations.
Aqueous Species
Th4+, ThOH3+, Th(OH)l+, Th(OH£, Th(OH)$(aq),
Th4(OH)*+, Th6(OH)£
Th(OH)3CC>3 andTh(CO3)
ThF3+, ThFf, ThF^, ThF5(aq)
ThCl3+, ThClf, ThCl^, ThCl^aq)
ThSOf, Th(S04)l(aq), Th(SO4)r, Th(SO4)t
ThH3P04+, ThH2POf , Th(H2PO4)f ,
Th(HP04)l(aq), Th(HPO4):
,2-
5.56
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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 et al. (1994) have recently published
thermodynamic constants for the thorium carbonate complexes Th(OH)3CO3 and Th(CO3)5~ that
are based on their solubility studies of microcrystalline ThO2 at different partial pressures of CO2
in aqueous media.
s
•
100
so
60
40
20
678 9 10
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 ug/1 total
dissolved thorium in pure water (i.e., absence of complexing ligands other
than OH") and thermodynamic data from Langmuir and Herman (1980).]
5.57
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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 ug/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 Th(OK)3CO~3 and Th(CO3)f" from
Osthols et 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)3CO3. The species distribution illustrated in 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 ug/1. At
the higher concentration of dissolved thorium, the stability of Th(OH)3CO'3 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 (C6H5Q*'), oxalate (CaOf), and
ethylenediamine tetra-acetic acid (EDTA) (C10H12O8N^) to show the possible role of organic
complexes hi the mobility of thorium hi 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.)SiO4], thorianite (crystalline
ThOz), monazite [(Ce,La,Th)PO4) and zircon (ZrSiO4), 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.
5.58
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100
a
o
.0
•**
.2
Q
e
3
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 ug/1 total dissolved thorium and thermodynamic data
from Langmuir and Herman (1980) and Osthols et al. (1994, for
Th(OH)3CO^ and Th(CO3)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 MNaClO4 solutions, the measured solubility of hydrous thorium oxide ranges
5.59
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from about lO'8-5 mol/1 (0.0007 mg/1) 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'z6 mol/1 (600 mg/1) as pH decreases from 5.0 to 3.2.
Felmy et 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 NaClO4 solutions. Moreover, the pH at which hydrous thorium oxide
exhibits rapid increases in solubility with decreasing pH changes from pH 5 in 0.1 M NaClO4 to
approximately pH 7 in 3.0 M NaCl. In studies conducted at high hydroxide and carbonate
concentrations, Rai et 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 ug/1) by adsorption processes. Humic substances are considered particularly important in
the adsorption of thorium (Gascoyne, 1982). Thibault et 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 et 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 (SiO2). 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 et al. (1988) studied the adsorption of thorium on goethite (a-FeOOH) and nsutite
(y-MnO2) in marine electrolyte solutions. Their experiments indicate that adsorption of thorium
1 Surface complexation models are discussed in Volume I of this report.
5.60
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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-MnO2 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 adsoiption 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 then" experiments, LaFlamme and Murray attributed this reduction to the
competition for surface sites by CO|" 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"9M. 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 (LaFleimme and Murray, 1987), or sulfate
(Hunter et al, 1988) concentrations, and high or low pH values (pH: 3 < x > 8: Hunter et al,
1988; LaFlamme and Murray 1987; Landa et al, 1995). These assumptions will be discussed hi
more detail in the following sections.
5.61
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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 et al., 1991; Rai et a/., 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
5.62
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used in the look-up table to describe high thorium concentrations (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 T), 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 (CaCO3) 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), low humic material concentrations (<5 mg/1), no organic
chelates (e.g., EDTA), and oxidizing conditions.]
Kd (ml/g)
Minimum
Maximum
pH
3-5
Dissolved Th,M
<10-2.e
62
6,200
>10-2.6
300,000
300,000
5-8
Dissolved Th,M
<10>9
1,700
170,000
>109
300,000
300,000
8-10
Dissolved Th, A/
<109
20
2,000
>109
300,000
300,000
5.63
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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 (£/2) 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 fH) is a radioactive isotope of hydrogen. Three isotopes of hydrogen are known. These
include the 2 stable isotopes XH (protium or H) and 2H (deuterium or D), and the radioactive
isotope 3H (tritium or T). Tritium has a half life (£/2) of 12.3 y, and disintegrates into helium-3
(3He) by emission of a weak beta (P") particle (Rhodehamel et 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 ^ 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 et 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.
5.64
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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 et 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 andKd 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
5.65
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uranium. Uranium(VT) species dominate in oxidizing environments. Uranium(VT) retention by
soils and rocks in alkaline conditions is poor because of the predominance of neutral or negatively
charged species. An increase in CO2 pressure in soil solutions reduces U(VI) adsorption by
promoting the formation of poorly sorting carbonate complexes. Uranium(IV) 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(IV) 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(TV) 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
23*U, 0.711 percent ^U, and 0.0054 percent ^U by weight. The half-lives of these isotopes are
4.51 x 109 y, 7.1 x 10s y, and 2.47 x 10s 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 (UO2 through UO2.25) and coffinite [USiOJ (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 [(K2(UO2)2(VO4)2], schoepite
(UO3-2H2O), rutherfordine (UO2CO3), tyuyamunite [Ca(UO2)2(VO4)2], autunite
[Ca(UO2)2(PO4)2], potassium autunite tK2(UO2)2(PO4)2], and uranophane [Ca(UO2)2(SiO3OH)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 ^U, 235V, and/or **U 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(DI) easily oxidizes to U(TV) under most reducing conditions
found in nature. The U(V) aqueous species (UOa) readily disproportionates to U(IV) and UOT).1
Consequently, U(IV) and U(VT) 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, UO2,* and U4"1", hydrolyze readily. The U4+ ion is more readily hydrolyzed
than UO|+, as would be expected from its higher ionic charge. Langmuir (1978) calculated U(IV)
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(IV) at pH values greater than 3 is
dominated by hydrolytic species such as U(OH)3 and \J(OK)l(a.q). 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 ug/1, because of the low solubility of U(IV) solid
phases (Bruno et a/., 1988; Bruno et al., 1991). Precipitation is discussed further in the next
section.
Dissolved U(VT) 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(VT)
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 ug/1, were used in these calculations. Hem (1985, p. 148)
1 Disproportionation is defined in the glossary at the end of this letter report. This particular
disproportionation reaction can be described as:
2UO2 + 4H3O+ = UCf + U4+.
2 A polynuclear species contains more than 1 central cation moiety, e.g., (UO2)2CO3(OH)3 and
5.67
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gives 0.1 to 10 ug/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 ug/1.
In a U(VI)-water system, the dominant species were UO2+ at pH values less than 5,
UO2(OH)2 (aq) at pH values between 5 and 9, and UO2(OH)3 at pH values between 9 and 10.
This was true for both uranium concentrations, 0.1 ug/1 (Figure 5.6a) and 1,000 ug/1 dissolved
U(VI) (Figure 5.6b). At 1,000 ug/1 dissolved uranium, some polynuclear species, (UO2)3(OH)5
and (UO2)2(OH)!'1', were calculated to exist between pH 5 and 6. Morris et 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 ug/1 dissolved U(VI).
At pH values less than 5, the UO2F+ species dominates the system, whereas at pH values greater
than 5, carbonate complexes [UO2CO3(aq), UO2(CO3)^, UO2(CO3)£j dominate the system.
These calculations clearly show the importance of carbonate chemistry on U(VT) 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 ug/1. At the lower
concentration of dissolved uranium, the species (UO2)2CO3(OH)^ is no longer present as a
dominant aqueous species.
Sandino and Bruno (1992) showed that UOf-phosphate complexes \\JO2BPO°4(aq) and UO2PO;]
could be important in aqueous systems with a pH between 6 and 9 when the total concentration
ratio PO4(total)/CO3(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 UO2+. Kim (1986) concluded that, in general, +6 actinides, including
U(VT), 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,
5.68
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U(TV) and U(VI) can form stable organic complexes, thereby increasing their solubility and
mobility.
Table 5.16. Uranium(VT) aqueous species included in the
speciation calculations.
Aqueous Species
, U02OEr, U02(OH)^(aq)3 UO2(OH)i , UO2(OH)*
(U02)2OH3+, (U02)2(OH)r, (U02)3(OH)f , (UO2)3(OH&
(U02)3(OH);, (U02)4(OH);,
), UO2(CO3>r, UO2(CO3#, UO2(CO3)f ,
(U02)3(C03)66-, (U02)n(C03)6(OH&, (U02)2C03(OH)3
UO2PO;, UO2EDPO;(aq), UO2H2PO^ UO2H3POf ,
U02(H2P04)^(aq), UO2(H2PO4)(H3PO4)+,
U02SO:(aq),U02(S04)
U02C1+, U02Cl°(aq), UO2F+,
UO2SiO(OH)^
), UO2F3,
5.11.4 Dissolution/Precipitation/Coprecipitation
Dissolution, precipitation, and coprecipitation have a much greater effect on the concentrations of
U(TV) 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(IV) to precipitate. Important U(TV) minerals
include uraninite (compositions ranging from UO2 to UO2 25), coffinite (USiO4), and ningyoite
[CaU(PO4)2-2H2O] (Frondel, 1958; Langmuir, 1978). Important U(VT) minerals include caraotite
[(K2(UO2)2(VO4)2], schoepite (UO3-2H2O), rutherfordine (UO2CO3), tyuyamunite
[Ca(UO2)2(VO4)2], autunite [Ca(UO2)2^O4)2], potassium autunite [K2(UO2)2(PO4)23, and
uranophane [Ca(UO2)2(SiO3OH)2] (Frondel, 1958; Langmuir, 1978). Carnotite, aU(VI) mineral,
is found in the oxidized zones of uranium ore deposits and uraninite, a U(IV) mineral, is a primary
5.69
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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 constants.
100
80
.2
Q
20
Figure 5.6a. Calculated distribution of U(VT) hydrolytic species as a function of pH
at 0.1 ug/1 total dissolved U(VT). [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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100
I
S
Figure 5.6b. Calculated distribution of U(VI) hydrolytic species as a function of pH at
1,000 ug/1 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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o
•-C
8
1
100
so
60
40
20
U02(C03)3
- U02F
U02C03 (aq) UO2(CO3)2
(U02)2C03(OH)3~
2-
10
pH
Figure 5.7. Calculated distribution of U(VT) 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 ug/1 total dissolved U(VT) and thermodynamic data
from Wanner and Forest (1992).]
5.11.5 Sorption/Desorption
In low ionic strength solutions with low U(VT) concentrations, dissolved uranyl concentrations
will likely be controlled by cation exchange and adsorption processes. The uranyl ion and its
complexes adsorb onto clays (Ames et al, 1982; Chisholm-Brause et al., 1994), organics
(Borovec et al., 1979; Read et al., 1993; Shanbhag and Choppin, 1981), and oxides (Hsi and
Langmuir, 1985; Waite et 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, forc-
ing it into solution. For this reason, the uranyl ion is particularly mobile in high ionic-strength
5.72
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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 et 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(VT) 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 etal., 1985; Borovec etal, 1979; Idiz etal, 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 etal., 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 etal., 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 et 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 ofKd 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(VT).
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 Kj 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 etal, 1992, 1994), clays (McKinley et
al, 1995; Turner et a!., 1996; Waite et al, 1992), and quartz (Waite et 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(VT) 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 uranium Kd values, a subjective approach was used to estimate the
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 et al. (1998), Redden et al. (1998),
Rosentreter et al. (1998), and Thompson etal. (1998) were identified and may be of interest to
the reader.
5.74
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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.
Ki
(ml/g)
Minimum
Maximum
pH
3
<1
32
4
0.4
5,000
5
25
160,000
6
100
1,000,000
7
63
630,000
8
0.4
250,000
9
<1
7,900
10
<1
5
The boundary representing the minimum limit for uranium Kd values is based on values calculated
for quartz from data given in Waite et 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 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 Kld values is based on Kd values calculated
for ferrihydrite and kaolinite from data given in Waite et 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(rV) redox/precipitation reactions resulting from contact of dissolved U(VI) with Fe(H)
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 etal. (1992, 1994),
McKinley etal. (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 carbonate complexes. In turn, this complexation affects
the adsorption behavior of U(VI) at alkaline pH conditions.
5.75
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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 CO2 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(VT) 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
otfiers, 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 K^ values for the geochemical conditions of interest, the reader is encouraged to
5.76
-------
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 NO3" and Cl", commonly adsorb to geological materials to a limited
extent. This is also the case observed for the sorption behavior of anionic Cr(VT).
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 MTNTEQA2 (Version 3.11, Allison et 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"1", and HTO,
respectively (Ames and Rai, 1978; Rai and Zachara, 1984). Chromium (as chromate, CrO^"),
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(VT) also forms polynuclear complex
species [complexes containing more than 1 mole of uranyl [e.g., (UO2)2CO3OH"].
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 catiomc 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 hi 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.
Element
Cd
Cs
Cr
Pb
Pu
Rn
Sr
Th
3H
u
Radio-
nuclide1
X
X
X
X
X
X
X
Primary Species at pH 7
and Oxidizing Conditions
Cationic
X
X
X
X
Anionic
X
X
X
X
X
Neutral
X
X
X
X
Redox
Sensitive2
X
X
X
X
X
Transport Through
Soils at pH 7
Not
Retarded3
X
X
X
Retarded3
X
X
X
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
Element
Cd
Cs
Cr
Pb
Pu
Rn
Sr
Th
3H
U
0.1 ug/1
U
1,000 jjg/1
pH3
Species
Cd2+
Cs+
HCr04
Pb2*
PbSO^aq)
PuFf
PuOJ
Pu3+
Rn°
Sr2*
TOFf
ThFj
HTO
uo2r
uor
U02Ff(aq)
U02F*
uor
U02FS(aq)
%
97
100
99
96
4
69
24
5
100
99
54
42
100
62
31
4
61
33
4
pH7
Species
Cd2*
CdHCOJ
CdCO^(aq)
Cs+
CrOi'
HCrO;
PbCO^(aq)
Pb2+
PbHCO^
PbOH+
Pu(OH)2(C03)i-
Pu(OH)°4(aq)
Rn°
Sr2*
Th(HPO4)^
Th(OH)3CO3
HTO
U02(C03)1-
U02(OH)^(aq)
U02CO^(aq)
U02P04
U02(C03)i-
(U02)2C03(OH)S
U02(OH)^(aq)
UO2CO|(aq)
%
84
6
6
100
78
22
75
15
7
3
94
5
100
99
76
22
100
58
19
17
3
41
30
13
12
pHIO
Species
CdCOf(aq)
Cs+
CrO2,'
PbCOf(aq)
Pb(C03)i-
Pb(OH)^(aq)
Pb(OH)+
Pu(OH)2(C03)2-
Pu(OH)5(aq)
Rn°
Sr2*
SrCOf(aq)
ThCO^COs
HTO
U02(C03)f
U02(OH)3-
U02(C03)2"
uo^co,),4-
U02(OH)i
U02(C03)1-
%
96
100
99
50
38
9
3
90
10
100
86
12
99
100
63
31
4
62
32
4
1 Only species comprising 3 percent or more of the total contaminant distribution are
presented. Hence, the total of the. percent distributions presented in table will not
always equal 100 percent.
5.79
-------
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(VT) (SO2',
sulfate) will convert into S(E) (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
Cd
Cs
Cr
Pb
Pu
Rn
Sr
Th
3H
U
Important Aqueous- and Solid-Phase Parameters Influencing
Contaminant Sorption2
[Aluminum/Iron-Oxide Minerals], [Calcium], Cation Exchange
Capacity, [Clay Mineral], [Magnesium], [Organic Matter], pH, Redox,
[Sulfide]
[Aluminum/Iron-Oxide Minerals], [Ammonium], Cation Exchange
Capacity, [Clay Mineral], [Mica-Like Clays], pH, [Potassium]
[Aluminum/Iron-Oxide Minerals], [Organic Matter], pH, Redox
[Aluminum/Iron-Oxide Minerals], [Carbonate, Fluoride, Sulfate,
Phosphate], [Clay Mineral], [Organic Matter], pH, Redox
[Aluminum/Iron-Oxide Minerals], [Carbonate, Fluoride, Sulfate,
Phosphate], [Clay Mineral], [Organic Matter], pH, Redox
None
Cation Exchange Capacity, [Calcium], [Carbonate], pH, [Stable
Strontium]
[Aluminum/Iron-Oxide Minerals], [Carbonate], [Organic Matter], pH
None
[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
-------
-------
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
-------
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)
UK DoE United Kingdom Department of the Environment
UNSCEAR United Nations Scientific Committee on the Effects of Atomic Radiation
A.3
-------
A.2.0 List of Symbols for the Elements and Corresponding Names
Symbol
Ac
Ag
Al
Am
Ar
As
At
Au
B
Ba
Be
Bi
Bk
Br
C
Ca
Cb
Cd
Ce
Cf
Cl
Cm
Co
Cr
Cs
Cu
Dy
Er
Es
Eu
F
Fe
Fm
Fr
Ga
Element
Actinium
Silver
Aluminum
Americium
Argon
Arsenic
Astatine
Gold
Boron
Barium
Beryllium
Bismuth
Berkelium
Bromine
Carbon
Calcium
Columbium
Cadmium
Cerium
Californium
Chlorine
Curium
Cobalt
Chromium
Cesium
Copper
Dysprosium
Erbium
Einsteinium
Europium
Fluorine
Iron
Fermium
Francium
Gallium
Symbol
Gd
Ge
H
He
Hf
Hg
Ho
I
In
Ir
K
Kr
La
Li
Lu
Lw
Md
Mg
Mn
Mo
N
Na
Nb
Nd
Ne
Ni
No
Np
O
Os
P
Pa
Pb
Pd
Pm
Element
Gadolinium
Germanium
Hydrogen
Helium
Hafnium
Mercury
Holmium
Iodine
Indium
Indium
Potassium
Krypton
Lanthanum.
Lithium
Lutetium
Lawrencium
Mendelevium
Magnesium
Manganese
Molybdenum
Nitrogen
Sodium
Niobium
Neodymium
Neon
Nickel
Nobelium
Neptunium
Oxygen
Osmium
Phosphorus
Protactinium
Lead
Palladium
Promethium
Symbol
Po
Pr
Pt
Pu
Ra
Rb
Re
Rh
Rn
Ru
S
Sb
Sc
Se
Si
Sm
Sn
Sr
Ta
Tb
Tc
Te
Th
Ti
Tl
Tm
U
V
w
w
Xe
Y
Yb
Zn
Zr
Element
Polonium
Praseodymium
Platinum
Plutonium
Radium
Rubidium
Rhenium
Rhodium
Radon
Ruthenium
Sulfur
Antimony
Scandium
Selenium
Silicon
Samarium
Tin
Strontium
Tantalum
Terbium
Technetium
Tellurium
Thorium
Titanium
Thallium
Thulium
Uranium
Vanadium
Tungsten
Wolfram
Xenon
Yttrium
Ytterbium
Zinc
Zirconium
A.4
-------
A.3.0 List of Symbols and Notation
A
ads
A,
am
aq
CEC
Ci
d
dpm
e"
Eh
F
g
3H
h
I
IAP
EEP
Kd
Porous media bulk density (mass/length3)
1
M
m
mCi
meq
mi
ml
mol
mV
N
n
He
pCi
pE
PH
ppm
Angstrom, 10"10 meters
Adsorption or adsorbed
Concentration of adsorbate (or species) I on the solid.phase at equilibrium
Amorphous
Aqueous
Cation exchange capacity
Curie
Day
Disintegrations per minute
Free electron
Redox potential of an aqueous system relative to the standard hydrogen electrode
Faraday constant, 23,060.9 cal/V-mol
Gram
Tritium
Hour
Ionic strength
Ion activity product
Isoelectric point
Concentration-based partition (or distribution) coefficient
Equilibrium constant at 298 K
Equilibrium constant at temperature T
Liter
Molar
Meter
Millicurie, 10"3 Curies
Milliequivalent
Mile
Milliliter
Mole
Millivolt
Constant in the Freundlich isotherm model
Total porosity
Effective porosity
Picocurie, 10"12 Curies
Negative common logarithm of the free-electron activity
Negative logarithm of the hydrogen ion activity
pH for zero point of charge
Parts per million
A.5
-------
R
Rf
s
sec
SI
T
t
tvi
TDS
TU
VP
y
z
Ideal gas constant, 1.9872 cal/mol-K
Retardation factor
Solid phase species
Second
Saturation index, as defined by log (IAP/K,. T)
Absolute temperature, usually in Kelvin unless otherwise specified
Time
Half life
Total dissolved solids
Tritium unit which is equivalent to 1 atom of 3H (tritium) per 1018 atoms
of *£[ (protium)
Velocity of contaminant through a control volume
Velocity of the water through a control volume
Year
Valence state
Charge of ion
Activity
Concentration
A.6
-------
-------
APPENDIX B
Definitions
-------
-------
Appendix B
Definitions
Adsorption - partitioning of a dissolved species onto a solid surface.
Adsorption Edge - the pH range where solute adsorption sharply changes from -10% 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 "He atom (Bates and Jackson 1980).
Alpha Recoil - displacement of an atom from its structural position, as in a mineral, resulting
from radioactive decay of the release an alpha particle from its parent isotope (e.g., alpha
decay of 222Rn from 226Ra).
Amphoteric Behavior - the ability of the aqueous complex or solid material to have a negative,
neutral, or positive charge.
Basis Species - see component species.
Cation Exchange - reversible adsorption reaction in which an aqueous species exchanges with an
adsorbed species. Cation exchange reactions are approximately stoichiometric and can be
written, for example, as
CaX(s) + "Sr^Caq) = 90SrX(s) + Ca2+(aq)
where X designates an exchange surface site.
Cation Exchange Capacity (CEC) - the sum total of exchangeable cations per unit mass of
soil/sediment that a soil can adsorb.
B.2
-------
Clay Content - particle size fraction of soil that is less than 2 urn (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
MHSTTEQA2 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.
Disproportionate - is a chemical reaction in which a single compound serves as both oxidizing
and reducing agent and is thereby converted into more oxidized and a more reduced
derivatives (Sax and Lewis 1987). For the reaction to occur, conditions in the system must be
temporarily changed to favor this reaction (specifically, the primary energy barrier to the
reaction must be lowered). This is accomplished by a number of ways, such as adding heat or
microbes, or by radiolysis occurring. Examples of plutonium disproportionation reactions are:
3Pu4+ + 2H2O = 2Pu3+ + PuOf +4IT
SPuOj + 4H+ = Pu3+ + 2PuOf +2H2O.
Electron Activity - unity for the standard hydrogen electrode.
i
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).
B.3
-------
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
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.
B.4
-------
Polynuclear Species - an aqueous species containing more than 1 central cation moiety, e.g.,
and Pb4(OH)f.
Protium (H) - stable isotope JH of hydrogen.
Retrograde Solubility - solubility that decreases with increasing temperature, such as those of
calcite (CaCO3) and radon. The solubility of most compounds (e.g., salt, NaCl) increases with
increasing temperature.
Species - actual form in which a dissolved molecule or ion is present in solution.
Specific Adsorption - surface complexation via a strong bond to a mineral surface. For example,
several transition metals and actinides are specifically adsorbed to aluminum- and iron-oxide
minerals.
Sol - a homogeneous suspension or dispersion of colloidal matter in a fluid.
Solid Solution - a solid material in which a minor element is substituted for a major element in a
mineral structure.
Thoron - name occasionally used, especially in older documents, to refer to 220Rn which forms
from the decay of thorium.
j
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).
B.5
-------
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 Kj 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. 1.
C.2
-------
Table C.I. Descriptive statistics of the cadmium Kd data set for soils.
Mean
Standard
Error
Median
Mode
StdDev
Sample
Variance
Range
Minimum
Maximum
No.
Samples
Cadmium
Kd
(ml/g)
226.7
44.5
121.8
80.0
586.6
344086
4359
0.50
4360
174
Clay
Content
(wt%)
14.2
1.7
10.24
6
13.5
182
86.2
.9
87.1
62
pH
5.88
0.09
5.83
6.8
1.16
1.34
6.20
3
9.2
174
CEC
(meq/lOOg)
21
3
23
2
15
245
58
2
60
22
TOC
(mg/1)
5.5
0.85
2.0
0.4
6.8
45.9
32.4
0.2
32.6
63
Cd Cone.
(mg/l)
3.67
0.48
0.01
0.01
6.27
39.4
34.9
0.01
35
172
Fe Oxides
(wt.%)
1.32
0.53
0.38
0.19
2.12
4.51
8.28
0.01
8.29
16
C.2.0 Approach and Regression Models
C2.1 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,
logOQ
Clay Cone.
pH
CEC
TOC
Cd Cone.
Fe Oxide
Cone.
Cadmium
K,,
1
0.69
-0.04
0.50
0.40
0.20
-0.02
0.18
logCK,,)
1
0.03
0.75
0.41
0.06
-0.10
0.11
Clay
Content
1
0.06
0.62
0.13
-0.39
-0.06
PH
1
0.35
-0.39
0.22
0.16
CEC
1
0.27
-0.03
0.19
TOC
1
-0.09
0.18
Cd Cone.
1
0.01
C.2.2 Cadmium Kd Values as a Function ofpH
The cadmium Kd values plotted as a function of pH are presented in Figure C.I. 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.I is:
CdKd = -0.54 + 0.45(pH).
(C.I)
The minimum and maximum values were estimated based on the scatter of data points observed in
Figure C.I.
C.4
-------
4
3.5
3
/—N
M' 2'5
T3 •?
P. 2
1.5
1
0.5
0
00
O 00
= -0.55 + 0.45x; r = 0.75
.2 L
J U
J L
6 7 8 9 10
pH
Figure C.I. Relation between cadmium Kd values and pH in soils.
Table C.3. Look-up table for estimated range of Kj 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.]
IQ (ml/g)
Minimum
Maximum
pH
3-5
1
130
5-8
8
4,000
8-10
50
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)
52.5
288.4
13.9
186.6
52.7
91.2
28.8
97.9
5.5
755.1
Clay
Cont.
(wt%)
54.7
8.3
51.2
0.9
17.6
28.2
2.8
6.2
3.8
23.9
pH
4.8
5.7
5.4
5.9
3.9
6
6.9
6.6
4.3
7.6
CEC
(meq/
100 g)
30.2
2
2.4
22.54
26.9
11
4.1
8.6
2.7
48.1
TOC
(wt%)
1.54
0.61
0.26
6.62
11.6
1.67
0.21
0.83
1.98
4.39
[Cd]
(mg/1)
1
1
1
1
1
1
1
1
1
1
Fe
Oxides
(wt%)
0.33
0.1
0.08
1.68
1.19
0.19
0.06
0.3
0
0.19
Solution
0.005 M
CaNO3
0.005 M
CaNO3
0.005 M
CaN03
0.005 M
CaNOj
0.005 M
CaNO3
0.005 M
CaNO3
0.005 M
CaNO3
0.005 M
CaNO3
0.005 M
CaNO3
0.005 M
CaNO3
Soil
Identification
Alligator Ap
Cecil Ap
CecilB
KulaApl
LafitteAp
Molokai Ap
Norwood Ap
Olivier Ap
Spodisol
Webster Ap
Comments
Converted
Freund. to K^
Using Ippm
Converted
Freund. to Kj
Using Ippm
Converted
Freund. to Ka
Using Ippm
Converted
Freund. to Ki
Using Ippm
Converted
Freund. to IQ
Using Ippm
Converted
Freund. to Ka
Using Ippm
Converted
Freund. to Kj
Using Ippm
Converted
Freund. to K,!
Using Ippm
Converted
Freund. to K^
Using Ippm
Converted
Freund. to K,j
Using Ippm
Ref."
1
1
1
1
1
1
1
1
1
1
C.6
-------
(ml/g)
14.4
87.1
33.88
120.42
10.47
80
200
1 133.3
181.8
266.7
g
11 l^
32
64
92
11 110
| 250
Clay
Cont
(wt%)
2.8
8
8
8
8
8
8
8
pH
5.3
8.4
5.2
5.8
6
8.2
7.8
8.3
7.6
7.9
3.7
4.8
5.3
6
6.2
6.8
7.3
CEC
(meq/
100 g)
2
60
33.8
23.8
25
8.2
15.4
18.9
31.8
37
TOC
(wt%)
2.03
1.44
32.6
3
3.2
0.21
0.83
0.23
0.79
0.86
1.6
1.6
1.6
1.6
1.6
1.6
1.6
[Cd]
(mg/1)
1
1
1
1
1
35
'25
30
25
15
11.2
11.2
11.2
11.2
11.2
11.2
11.2
Fe
Oxides
(wt%)
0.42
1.07
8.29
1.07
Solution
0.005 M
CaNO3
Water
Water
Water
Water
0.01 M
NaCl
0.01 M
NaCl
0.01 M
NaCl
0.01 M
NaCl
0.01 M
NaCl
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaN03
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaN03
0.01 M
NaN03
Soil
Identification
Windsor Ap
Vertic
Torrifluvent
Organic
Boomer, Ultic
Haploxeralf
UlticPalexeral
f
Gevulot
BetYizhaq
Gilat
Maaban
Michael
Hahoterim
Downer
Loamy Sand
Downer
Loamy Sand
Downer
Loamy Sand
Downer
Loamy Sand
Downer
Loamy Sand
Downer
Loamy Sand
Downer
Loamy Sand
Comments
Converted
ireund. to K,j
Using Ippm
Converted
ireund. to Kj
Jsing Ippm
Converted
Freund. to K^
Using Ippm
Converted
Freund. to Kj
Using Ippm
Converted
Freund. to Ka
Using Ippm
Calc.Figl.
Calc.Fig 1.
Calc.Figl.
Calc.Fig 1.
Calc.Fig 1.
Ref.°
1
2
2
2
2
3
3
3
3
3
4
4
4
4
4
C.7
-------
CdK*
(ml/g)
580
0.5
3.3
7.5
10
34
45
80
150
420
900
2.1
10
30
57
Clay
Cont
(wt%)
8
6
6
6
6
6
6
6
6
6
6
13
13
13
13
PH
8.5
3.1
3.8
4.5
5.5
6.1
6.8
7.5
8
8.4
9.1
3
3.7
4.2
4.6
CEC
(meq/
100 g)
TOC
(wt%)
1.6
0.4
0.4
0.4
0.4
0.4
0.4
0.4
0.4
0.4
0.4
16.8
16.8
16.8
16.8
[Cd]
(mg/1)
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
Fe
Oxides
(wt%)
Solution
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaN03
0.01 M
NaN03
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaN03
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaNOj
SoU
Identification
Downer
Loamy Sand
Freehold
Sandy Loam A
Horizon
Freehold
Sandy Loam A
Horizon
Freehold
Sandy Loam A
Horizon
Freehold
Sandy Loam A
Horizon
Freehold
Sandy Loam A
Horizon
Freehold
Sandy Loam A
Horizon
Freehold
Sandy Loam A
Horizon
Freehold
Sandy Loam A
Horizon
Freehold
Sandy Loam A
Horizon
Freehold
Sandy Loam A
Horizon
Boonton Loam
Boonton Loam
Boonton Loam
Boonton Loam
Comments
Ref.»
4
4
4
4
4
4
4
4
4
4
4
4
4
4
4
C.8
-------
CdKa
(ml/g)
101
195
420
1,200
4,000
1.2
7.1
27
53
170
300
390
910
1,070
43
67
130
Clay
Cont
(wt%)
13
13
13
13
13
16
16
16
16
16
16
16
16
16
10
10
10
pH
5
5.2
5.8
6.2
6.8
3.3
4.1
4.8
5.1
5.6
6.1
6.2
6.5
6.8
4.8
5.7
6.3
CEC
(meq/
100 g)
TOC
(wt%)
16.8
16.8
16.8
16.8
16.8
9.8
9.8
9.8
9.8
9.8
9.8
9.8
9.8
9.8
2.4
2.4
2.4
[Cd]
(mg/1)
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2
11.2 '
11.2
Fe
Oxides
(wt%)
Solution
0.01 M
NaNO3
0.01 M
NaN03
0.01 M
NaNO3
0.01 M
NaN03
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaN03
0.01 M
NaN03
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaNO3
SoU
Identification
Boonton Loam
Boonton Loam
Boonton Loam
Boonton Loam
Boonton Loam
Rockaway
Stony Loam
Rockaway
Stony Loam
Rockaway
Stony Loam
Rockaway
Stony Loam
Rockaway
Stony Loam
Rockaway
Stony Loam
Rockaway
Stony Loam
Rockaway
Stony Loam
Rockaway
Stony Loam
Fill Material -
Delaware
River
Fill Material -
Delaware
River
Fill Material -
Delaware
River
Comments
Ref."
4
4
4
4
4
4
4
4
4
4
4
4
4
4
4
4
4
C.9
-------
CdlQ
(ml/g)
150
370
880
1,950
1,000
4,360
536.8
440
9
23.4
15.8
11.3
31.2
32.5
23
17.1
Clay
Cont.
(wt%)
10
10
10
10
12
12.4
25.2
25.2
PH
6.7
7.3
8
9.2
8
8
6.8
6.8
4.3
4.3
4.4
4.5
4.5
4.5
4.5
4.7
CEC
(meq/
100 g)
27.5
27.5
TOC
(wt%)
2.4
2.4
2.4
2.4
[Cd]
(mg/1)
11.2
11.2
11.2
1L2
1
1
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
Fe
Oxides
(wt%)
3.7
2.5
Solution
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaNO3
0.01 M
NaNO3
Carbonate
Grouiidwate
T
Carbonate
Groundwate
r
0.01 M
NsiCl
0.01 M
NaCl
0.001M
CaClj
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.00 1M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
Soil
Identification
Fill Material -
Delaware
River
Fill Material -
Delaware
River
Fill Material -
Delaware
River
Fill Material -
Delaware
River
Interbed
Alluvium
Soil A
Soil A
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Comments
pHof
Groundwater
pHof
Groundwater
Desorption
Desorption
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Re£°
4
4
4
4
5
5
6
6
7
7
7
7
7
7
7
7
C.10
-------
CdK,,
(ml/g)
13.1
24.9
26.8
36.2
32.9
37.2
29.2
28.3
22.6
37.4
40.9
63.5
25.2
29.9
33.7
44.3
42.8
53.5
56.2
Clay
Cont
(wt%)
pH
4.8
4.6
4.7
4.7
4.7
4.7
4.8
4.8
4.9
4.9
4.9
4.7
5.4
5.3
5.2
5.1
5.1
5
4.9
CEC
(meq/
100 g)
TOC
(wt%)
[Cd]
(mg/1)
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
Fe
Oxides
(wt.%)
Solution
0.001M
CaCl2
0.001M
CaClj
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
Soil
Identification
Agricultural
Danish Soil
.Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Comments
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Re£"
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
C.ll
-------
CdK,,
(ml/g)
68.7
82.3
75.7
95.2
103
160
43.3
55.2
52.2
40.3
56.1
67.5
102.9
164.4
163.8
202.1
172.4
149
72.8
Clay
Cont
(wt%)
pH
5
5.1
5
4.8
4.8
4.8
5.4
5.4
5.3
5.6
5.5
5.5
5.4
5.5
5.3
5.2
5.2
5.2
5.6
CEC
(meq/
100 g)
•
TOC
(wt%)
[Cd]
(n»g/l)
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
Fe
Oxides
(wt%)
Solution
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCI2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCk
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.00 1M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
Soil
Identification
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Comments
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Ref."
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
C.12
-------
CdK,,
(ml/g)
81.6
90
94.3
48.1
56.5
81
122.3
121.4
101.5
99.3
107.8
219.5
179.2
177
360.4
305.2
236.8
186.3
174.8
Clay
Cont
(wt%)
pH
5.7
5.7
5.6
6.2
6.4
6.5
6.4
6.2
6
6
6
6.2
6.2
6.1
6
6
5.9
5.9
5.8
CEC
(meq/
100 g)
TOC
(wt%)
[Cd]
(mg/1)
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
Fe
Oxides
(wt%)
Solution
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.00 1M
CaCl2
0.001M
CaC^
0.001M
CaCl2
Soil
Identification
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Comments
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Ref."
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
C.13
-------
CdK,,
(ml/g)
138.7
132.5
375.6
403.3
510.8
225.9
227.3
248
253.1
277.2
240.7
227.8
281.1
551.2
519.8
418.7
353.7
400.8
609.2
Clay
Cont.
(wt%)
PH
5.8
5.7
5.9
5.8
5.8
5.7
5.7
5.7
5.6
5.6
6.4
6.5
6.6
6.2
6.2
6.2
6.2
6.4
6.3
CEC
(meq/
100 g)
TOC
(wt%)
[CdJ
(mg/1)
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
Fe
Oxides
(wt%)
Solution
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
Ca.Cl2
0.001M
Ca.Cl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
Soil
Identification
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Comments
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Ref."
7
7
7
7
7
7
7-
7
7
7
7
7
7
7
7
7
7
7
7
C.14
-------
CdlQ
(ml/g)
545.7
515.9
545.7
760.9
665.7
503.2
515.2
488.9
481
461.6
1,151
868.7
637.2
970.9
950.5
886.2
1,106
970.9
2,248
Clay
Cont
(wt%)
pH
6.3
6.4
6.4
6.4
6.5
6.5
7
6.9
6.9
6.9
6.5
6.6
6.7
6.7
6.8
6.9
6.9
7
7.1
CEC
(meq/
100 g)
TOC
(wt%)
[Cd]
(mg/1)
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
0.01
Fe
Oxides
(wt.%)
Solution
0.001M
CaCl2
0.00 1M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.00 1M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
0.001M
CaCl2
Soil
Identification
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Comments
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co =0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Ref."
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
7
C.15
-------
CdKa
(ml/g)
1,909
1,411
1,383
2,337
Clay
Cont.
(wt%)
pH
7.2
7.3
7.4
7.5
CEC
(meq/
100 g)_
TOC
(wt%)
[Cd]
(mg/1)
0.01
0.01
0.01
0.01
Fe
Oxides
(wt.%)
Solution
0.001M
CaCl2
0.001M
CaClj
0.001M
CaCl2
0.001M
CaCl2
Soil
Identification
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Agricultural
Danish Soil
Comments
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Co = 0.7 to
12.6 ppb
Ref."
7
7
7
7
a 1 = Buchter etal, 1989; 2 = Garcia-Miragaya, 1980; 3 = Navrot et al, 1978; 4 = Allen et al, 1995; 5 = Del Debbio,
1 99 1 ; 6 = Madrid et al. , 1 992; 7 = Anderson and Christensen , 1 988
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."
Radiochimica Acta, 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 et a/.,
1986) (Table D.I). 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.I. Descriptive statistics of cesium Kd data set including soil and pure mineral
phases. [Data set is presented in Section D.3.]
Mean
Standard Error
Median
Mode
Standard Deviation
Sample Variance
Range
Minimum
Maximum
No. Observations
Confidence Level
(95.0%)
Kd (ml/g)
2,635
530
247
40
7055
49,781,885
51,999
0.6
52,000
177
1,046.6
Clay
_(%)
.30
3.8
42
42
15
226
38
4
42
15
8.3
Mica
(%)
5.5
0.7
4
4
4.4
20.0
13
2
15
41
1.4
pH
7.4
0.1
8.2
8.2
1.7
2.8
7.8
2.4
10.2
139
0.3
CEC
(meq/100 g)
30.4
3.7
4.8
1.8
37.4
1,396.9
129.9
0.00098
130
103
7.3
Surface Area
(m2/g)
141.3
29.7
31.2
17.7
230.4
53,106
638
8
646
60
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 et 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. Pata set is presented
in Section D.4.]
Mean
Standard Error
Median
Mode
Standard Deviation
Sample Variance
Range
Minimum
Maximum
No. Observations
Confidence Level (95%)
Cesium
K,
(ml/g)
651
188
89
22-
1423
2026182
7602
7.1
7610
57
378
Clay
(%)
5
0.6
5.0
NA
1.0
1.0
2.0
7.1
6.0
3
2.5
Mica
(%)
5.6
0.6
4
4
4.3
18.4
13
2
15
45
1.29
pH
6.9
0.3
6.7
4.0
1.9
3.6
7.8
2.4
10.2
55
0.5
CEC
(meq/lOOg)
34
8.9
20
60
29.5
870
57.4
2.6
70.0
11
19.8
Surface Area
(m2/g)
57.5
13.4
60
70
44.6
1986
123.4
6.6
130
11
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]^) 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 Kd
Clay Content
Mica
PH
CEC
Surface Area
[Cs]aq
Cesium
Krf
1.00
0.05
0.29
0.10
0.52s
0.42a
-0.07
Clay
Content
1.00
0.00
-0.11
0.64a
0.35
0.85a
Mica
1.00
0.08
NA
NA
0.29
pH
1.00
0.37
•-0.11
0.13
CEC
1.00
0.47"
-0.17
Surface Area
1.00
-0.15
a Correlation coefficient is significant at the 5% level of significance (P z 0.05).
Table D.4. Correlation coefficients (r) of the soil-only data set. [Data set is
presented in Section D.4.]
Cesium Kd
Clay Content
Mica
pH
CEC
Surface Area
[Cs]aq
Cesium
Krt
1.00
-0.21
0.27
0.11
-0.83
-0.31
0.18
Clay
Content
1.00
0
0.4
NA
NA
NA
Mica
1.00
0.07
0.991
0.991
0.09
pH
1.00
0.05
-0.03
-0.04
CEC
1.00
0.37
0.00
1 Correlation coefficient is significant at >5% level of significance (P <
Surface Area
1.00
0
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:
log (Cs Ka) = 1.2 + 1.0 log (CEC)
(D.I)
A similar regression analysis using the entire data set (mineral, rocks, and soils) is presented in
Figure D.I.
6
5
4
3
0
-1
—i 1 1 1 1 1 r
. y = 2.09 + 0.73x,r = 0.60
1 1 L
-1.5 -1 -0.5 0 0.5 1 1.5 2 2.5
log CEC (meq/100 g)
Figure D.I. Relation between cesium Kd values and CEC.
D.7
-------
By transposing the CEC and cesium K^ data into logarithms, the regression correlation slightly
increases from 0.52 (Table D.3) to 0.60 (Figure D.I). 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
KJ data set only varies 5 orders of magnitude. Thus, the correlation with CEC, although the
strongest of all the independent variables examined, did riot reduce greatly the variability of
possible cesium Kd values.
D.2.3 CEC as a Function of Clay Content andpH
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 Kj values.
0 10 20 30 40 50 60 70
Clay(%)
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 etal., 1989; Cremers etal, 1988; Douglas, 1989;
Evans etal., 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 hi 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
KJ 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
1.67xlO-12M.]
Mineral
Phases
Mica
Vermiculite
Kaolinite
CEC
(meq/100 g)
20
127
11.2
Kd in Water
(ml/g)
26,000
52,000
2,500
Kd in 0.1 M NaCl
(ml/g)
28,600
2,700
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
K-saturated Mica, TxlCT6 M Cs
K-saturated Mica, 2xlO"7 M Cs
Ca-saturated Mica, 7xlO'6 M Cs
Ca-saturated Mica, 2x1 0"7 M Cs
Adsorption
Cesium K,,
2,890
9,000
1,060
600,000
Desorption
Cesium KJ
5,200
11,300
4,600
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)
1 x ID'7
1 xlO"10
5 x 10'5
1 x ID'10
SxlO-9
1 x ID"8
5 x 10'8
5 x 10'7
IxlO-6
1 x 1CT1
<1 x lO"5
<1 x ICT5
<1 x 10's
1 x ID'3
Solid Phase
Itado Tuff
Sandstone
Limestone
Augite Andesite
Olivine Basalt
Rokko Granite
Biotite
Albite
K-Feldspar
Unwashed Kaolinite
Ca Montmorillonite
Na Montmorillonite
Na Kaolinite
Na Montmorillonite
Background
Aqueous Phase
Deionized Water
Deionized Water
Deionized Water
Deionized Water
Deionized Water
Deionized Water
Deionized Water
Deionized Water
Deionized Water
Distilled Water/pH 10
Distilled Water/pH 10
Distilled Water/pH 10
Distilled Water/pH 10
Distilled Water/pH 4
Reference
Akidaetal., 1989
Akid&etal., 1989
Akidaetal., 1989
Akidaetal., 1989
Akidaetal., 1989
Akidaetal., 1989
Akidaetal., 1989
Akidaetal., 1989
Akidaetal., 1989
Adeleye ef a/., 1994
Adeleye ef a/., 1994
Adeleyeetal., 1994
Adeleye era/., 1994
Adeleyee/a/., 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 et al, 1988; Lieser
and Staunton, 1994; Steinkopff, 1989; Torstenfelt et al., 1982). Most investigators have used a
Freundlich equation to describe this relationship (Adeleye et al, 1994; Konishi et al, 1988; Shiao
etal, 1979; Staunton, 1994; Torstenfelte/a/., 1982). The Freundlich equation is
— a (Cssolution)
(D.2)
where Csabsotbed and Cs^^o,, 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.
105
IV
104
t 1Q3
•e
^ 102
o
101
in0
1U
11111111
• o ]
r 8 8 .,
GO 0
• |8 :
rlSo o 1
1 ft wo
• I ° •
O O v w
o o
1 1 1 1 1 1 1 1
0 0.2 0.4 0.6 0.8 1 1.2 1.4 1.
Solution Cs (]umol/l)
6
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.
a1
1.7
3,300
260
16
12.2
6,070
1,290
163
1.23
0.63
427
1.5
48.1
17
5.22
4.4
0.22
0.017
0.13
0.048
S.lOxlO-4
S.OOxlO'3
1.30xlO-s
2.30xlO's
b1
0.677
0.909
0.841
0.749
0.745
0.899
0.849
0.815
0.657
0.659
0.814
0.599
0.754
0.739
0.702
0.716
1.1
0.53
1
0.67
0.21
0.48
0.013
0.38
Range of Solution Cs
Concentration (M)
Ixl0-8tolxl0-12
lxlO'8 to IxlO'12
lxlO'8 to lxlO'12
IxlO'8 to IxlO-12
IxlO'8 to IxlO'12
IxlO"8 to IxlO'12
IxlO'8 to IxlO'12
IxlO'8 to IxlO"12
Ixl0-9tol.5xl0-2
Ixl0-9tol.5xl0-2
Ixl0-9tol.5xl0-2
Ixl0-9tol.5xl0-2
Ixl0-9tol.5xl0-2
Ixl0-9tol.5xl0-2
Ixl0-9tol.5xl0-2
Ixl0-9tol.5xl0-2
Experimental
Water/Batcombe Sediment
Water/Denchworth Sediment
Water/Tedburn Sediment
Water/Teigngrace Sediment
Water/Batcombe Sediment
Water/Denchworth Sediment
Water/Tedbum Sediment
Water/Teigngrace Sediment
CaCl2/Batcombe Sediment
CaCl2/Batcombe Sediment
CaCl2/Denchworth Sediment
CaCl2/Denchworth Sediment
CaCl2/Tedburn Sediment
CaCl/Tedburn Sediment
CaCl/Teigngrace Sediment
CaCl2/Teigngrace Sediment
Bentonite/Water
Bentonite/Water
Bentonite/Groundwater
Bentonite/Groundwater
Xakadata Loam/Water
Xakadata Loam/Qroundwater
Hachinohe Loaitn/Water
Hachinohe Loam/Groundwater
Ref.2
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
1
2
2
2
2
2
2
2
2
D.13
-------
a1
2:70xlO-4
5.20X1CT4
2.04X10"3
2.27xlO'3
5.04xlCT2
3. 49x1 0'2
0.235
3.03xlO'2
0.135
0.247
8.71X10"3
1. 02x10"*
l.OSxlO'2
3.17xlO'2
0.224
0.241
0.481
1.84
0.274
3.40xlO'2
4.90xlO"2
4.00X10'2
b1
0.546
0.543
0.588
0.586
0.723
0.703
0.821
0.804
0.845
0.881
0.694
0.503
0.709
0.755
0.815
0.839
0.897
0.938
0.82
0.51
0.5
0.5
Range of Solution Cs
Concentration (M)
Ixl0-8tolxl0-2
IxlO'8 to IxlO'2
Ixl0-8tolxl0-2
IxlO'8 to IxlO'2
IxlO-8 to lxlO'2
IxlO'8 to IxlO'2
lxlO'8 to IxlO-2
IxlO-8 to IxlO'2
IxlO'8 to IxlO'2
lxlO'8 to lxlO'2
IxlO-8 to lxlO'2
Ixl0-8tolxl0-2
lxlO'8 to IxlO-2
IxlO-8 to lxlO'2
IxlO'8 to IxlO'2
IxlO'8 to IxlO"2
IxlO'8 to IxlO'2
IxlO'8 to IxlO'2
IxlO-8 to lxlO'2
lxlO'7 to IxlO"3
lxlO'7 to lxlO'3
Experimental
Unwashed/Kaolinite/pH 2
Unwashed/Kaolinite/pH 4
Unwashed/Kaolinite/pH 10
Sodium/Kaolinite/pH 2
Sodium/Kaolinite/pH 4
Na/Kaolinite/pH 7
Na/Kaolinite/pH 10
Ca/Kaolinite/pH 2
Ca/Kaolinite/pH 4
Ca/Kaolinite/pH 7
Ca/Kaolinite/pH 10
Na/Montmorillonite/pH 2
Na/Montmorillonite/pH 4
Na/Montmorillonite./pH 7
Na/Montmorillonite/pH 10
Ca/Montmorillonite/pH 2
Ca/Montmorillonite/pH 4
Ca/Montmorillonite/pH 7
Ca/Montmorillonite/pH 10
Granite/pH 8.2
Granite/pH 8.2
Ref.2
3
3
3
3
3
3
3
3
3
3
3
3
3
3
3
3
3
3
3
4
4
5
1 Parameters "a" and "b" are fitting parameters in the Freundlich equation.
2 References: 1 = Fukui, 1990; 2 = Konishi et al, 1988; 3 = Adeleye et al, 1994; 4 = Seme et
al, 1993; 5 = Shiao etal, 1979.
D.14
-------
Table D.9. Descriptive statistics of the cesium Freundlich equations (Table D.8)
reported in the literature.
Statistic
Mean
Standard Error
Median
Mode
Standard Deviation
Sample Variance
Range
Minimum
Maximum
95% Confidence Level
a
252
150.2
0.222
NA
1019
1038711
6070
0.000013
6070
302
b
0.696
0.029
0.720
0.815
0.198
0.039
1.087
0.013
1.1
0.059
Using the medians of the a and b parameters from the literature, we come up with the overall
equation:
(D.3)
This equation is plotted in Figure D.4. Using Csadsolbed and Cssolutlon from equation D.3, a Kd value
can be calculated according to equations D.4,
(D.4)
Cesium K, values calculated from Equations D.3 and D.4 are presented in Figure D.5.
D.15
-------
-7
•SP
1
•3
03
0
io
10'8
10
-
10
-12
I I I I II I
10-
1S
Solution Cs (mol/1)
Figure D.4. Generalized cesium Freundlich equation
(Equation D.3) derived from the literature.
1000 F-T
10'10 IO* ID'8 10'7 1Q-6 10'3
Solution Cs (moLT)
D.16
-------
D. 2.6 Approach to Selecting Kd Values for Look-up Table
i
Linear regression analyses were conducted with data collected from the literature. These analyses
were used as guidance for selecting appropriate Kd values for the Hook-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 Kj 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.I 1) were assigned by multiplying the central estimates from
Table D.10 by a factor of 2.5. The 2.5 sealer 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.I.
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/I), no organic chelates (such as EDTA), and oxidizing
conditions]
Kd (ml/g)
Central
Minimum
Maximum
CEC (meq/100 g) / Clay Content (wt.%)
<3/<4
200
10
3,500
3-10/4-20
500
30
9,000
10-50/20-60
1,500
80
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.]
KdOnl/g)
Central
Minimum
Maximum
CEC (meq/100 g) / Clay Content (wt.%)
<3/<4
500
30
9,000
3-10/4-20
1250
70
22,000
10 - 50 / 20 - 60
3750
210
66,700
D.18
-------
Table D.12. Calculations for values used in look-up table.
Mica
Concentration
in Clay Fraction
(%)
<5
<5
<5
>5
>5
>5
Clay
Content
(wt%)
<4
4-20
20-60
<4
4-20
20-60
CE1
(ml/g)
200
500
1,500
500
1,250
3,750
Logarithm Scale
LogCE
2.301
2.699
3.176
2.699
3.097
3.574
Lower Limit
(LogCE)/2
1.151
1.349
1.588
1.349
1.548
1.787
Base-10 Scale
Lower Limit
10aogCE>/2(ml/g)
14
22
39
22
35
61
Upper Limit
IfllogCE+OogCEVZ^l/g)
2,828
11,180
58,095
11,180
44,194
229,640
1 CE = Central Estimate
D.19
-------
D.3.0 KJ Data Set for Soils and Pure Mineral Phases
Table D.I3 lists the available cesium K,, 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)
247
62
22
16
12
167
1
1500
160
1100
4100
1400
1100
280
237
8220
325
22100
35800
42600
205
Clay
(wt.%
)
Mica
(%)
pH
6.2
6.2
6.2
6.2
6.2
8.1
7.8
9.3
2.4
9.3
6.1
7.7
6.6
8.3
8.2
8.2
8.2
8.2
8.2
8.2
8.2
CEC"
(meq/100 g)
60
60
60
20
20
70
70
2
109
6
51
107
107
4
SA1
(mVg)
189
113
70
70
70
130
130
60
60
22
103
43
19
Aqueous Cs
(uM)
1.90xlO-2
1.42x10-'
5.94X10'1
1.05
1.53
5.20xlO-3
5.20xlO'3
l.OOxlO-1
l.OOxlO'1
l.OOxlO'1
l.OOxlO-1
l.OOxlO-1
l.OOxlO-1
l.OOxlO-1
l.OOxlO-3
l.OOxlO-3
l.OOxlO-3
l.OOxlO-3
l.OOxlO"3
l.OOxlO-3
l.OOxlO-3
Background
Aqueous
Gorleben
Groundwater
Groundwater-1
Groundwater-2
Water pH 9.3
Groundwater
pH2.4
Groundwater
pH9.3
Water pH 6.1
Groundwater
pH7.7
Water pH 6.6
Groundwater
pH8.3
Soil and Mineral
Phase ID and
Information
Gorleben Sediment
Gorleben Sediment
Gorleben Sediment
Gorleben Sediment
Gorleben Sediment
SI: Quartz,
Kaolinite,
Plagioclase
S2:Quartz,
Kaolinite, Dolomite
Bentonite
Bentonite
Bentonite
Takadate loam
Takadate loam
Hachinohe loam
Hachinohe loam
ym-22
ym-38
ym-45
ym-48
ym-49
ym-49
ym-54
Ref2
1
1
1
1
1
2
2
3
3
3
3
3
3
3
4
4
4
4
4
4
4
D.20
-------
Cesium
Kd
(ml/g)
15200
8440
143
73
1390
757
95
120
130
130
150
160
72
79
75
98
83
33
37
40
39
50
27
25
26
26
38
39
88
92
93
Clay
(wtVe
)
Mica
(%)
15
15
15
15
15
15
3
3
3
3
3
4
4
4
4
4
2
2
2
2
2
2
4
4
4
pH
8.4
8.3
8.2
8.5
8.4
8.5
4
5.5
6.7
7
8.5
10.2
4
5.5
6.7
7
8.5
4
5.5
7
8.5
10.2
4
5.5
6.7
7
8.5
10.2
4
5.5
6.7
CEC"
(meq/100 g)
SA1
(mVg)
31
31
8
8
100
100
Aqueous Cs
(pM)
l.OOxlO'3
l.OOxlO'3
l.OOxlO-3
l.OOxlO-3
l.OOxlO-3
l.OOxlO-3
4.20x10-"
4.20x10"
4.20x10""
4.20x10""
4.20X10-4
4.20X10-4
4.20x10"
4.20X10-4
4.20x10"
4.20X10"4
4.20x10-"
4.20x10"
4.20x1 0-4
4.20x10"
4.20x10-"
4.20x10-"
4.20x10-"
4.20x10-"
4.20x10""
4.20x10-"
4.20x10-"
4.20x10-"
4.20x10"
4.20x10-"
4.20x10""
Background
Aqueous
low salts
hi salts
low salts
hi salts
low salts
hi salts
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
Soil and Mineral
Phase ID and
Information
JA-18
JA-18
JA-32
JA-32
JA-37
JA-37
Savannah River
Savannah River
Savannah River
Savannah River
Savannah River
Savannah River
•4-Mile Creek
4-Mile Creek
4-Mile Creek
4-Mile Creek
4-Mile Creek
Par Pond Soil
Par Pond Soil
Par Pond Soil
Par Pond Soil
Par Pond Soil
Steel Creek Soil
Steel Creek Soil
Steel Creek Soil
Steel Creek Soil
Steel Creek Soil
Steel Creek Soil
Lower 3 Runs Soil
Lower 3 Runs Soil
Lower 3 Runs Soil
Ref2
4
4
4
4
4
4
5
5
5
5
5
5
5
5
5
5
5
5
5
5
5
5
5
5
5
5
.5
5
5
5
5
D.21
-------
Cesium
Kd
(irf/g)
85
94
101
88
89
90
84
101
22
31
37
40
78
27
329
960
1088
1084
28
289
951
1022
1025
18
189
Clay
(wt%
)
Mica
(%)
4
4
4
5
5
5
5
5
2
2
2
2
2
PH
7
8.5
10.2
4
5.5
6.7
7
10.2
4
5.5
6.7
7
10.2
8.25
8.25
8.25
8.25
8.25
8.6
8.6
8.6
8.6
8.6
8.2
8.2
CEC"
(meq/lOOg)
1.83
1.83
1.83
1.83
1.83
1.83
1.83
1.83
1.83
1.83
1.5
1.5
SA1
(mVg)
17.7
17.7
17.7
17.7
17.7
17.7
17.7
17.7
17.7
17.7
10.3
10.3
Aqueous Cs
G«M)
4.20x10-"
4.20x10-"
4.20x10-"
4.20x10-"
4.20x10-"
4.20X10"
4.20x10-"
4.20x10-"
4.20x10""
4.20X10"1
4.20x10""
4.20x10-"
4.20x10"*
2.72xl02
2.90x10-'
l.OSxlO'3
9.11X10-6
1.87X10-6
2.63xl02
3.31X10'1
l.OSxlO-3
9.77x10-*
1.95X10-6
3.61xl02
S.OOxlO'1
Background
Aqueous
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.002 M
Groundwater
0.002 M
Groundwater
0.002 M
Groundwater
0.002 M .
Groundwater
0.002 M
Groundwater
0.013 M
Groundwater
0.013 M
Groundwater
0.013 M
Groundwater
0.013 M
Groundwater
0.013 M
Groundwater
0.002 M
Groundwater
0.002 M
Groundwater
Soil and Mineral .
Phase ID and
Information
Lower 3 Runs Soil
Lower 3 Runs Soil
Lower 3 Runs Soil
Pen Branch Soil
Pen Branch Soil
Pen Branch Soil
Pen Branch Soil
Pen Branch Soil
Upper 3 Runs Soil
Upper 3 Runs Soil
Upper 3 Runs Soil
Upper 3 Runs Soil
Upper 3 Runs Soil
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Flow E Basalt
Flow E Basalt
Ref2
5
5
5
5
5
5
5
5
5
5
5
5
5
6
6
6
6
6
6
6
6
6
6
6
6
D.22
-------
Cesium
Kd
(W/g)
418
450
487
20
214
488
549
617
48
460
1111
1466
1281
56
389
853
952
908
212
1080
Clay
(wt%
)
Mica
(%)
PH
8.2
8.2
8.2
8.7
8.7
8.7
8.7
8.7
8.3
8.3
8.3
8.3
8.3
8.55
8.55
8.55
8.55
8.55
8.3
8.3
CEC°
(meq/100 g)
1.5
1.5
1.5
1.5
1.5
1.5
1.5
1.5
4.84
4.84
4.84
4.84
4.84
4.84
4.84
4.84
4.84
4.84
71
71
SA1
(mVg)
10.3
10.3
10.3
10.3
10.3
10.3
10.3
10.3
31.2
31.2
31.2
31.2
31.2
31.2
31.2
31.2
31.2
31.2
646
646
Aqueous Cs
(MM)
2.34x10-'
2.17X10'5
3.98xlO-«
3.39xl02
4.47x10-'
2.00xlO-3
1.78X10'5
3.24X10-6
1.71xl02
2.13X10'1
S.SOxlO-4
6.37x1 0-6
1.39X10-6
l.SlxlO2
2.57x10-'
1.17xlO'3
l.OSxlO'5
1.74X10-6
4.50x10'
9.17x10-'
Background
Aqueous
0.002 M
Groundwater
0.002 M
Groundwater
0.002 M
Groundwater
0.013 M
Groundwater
0.013 M
Groundwater
0.013 M
Groundwater
0.013 M
Groundwater
0.013 M
Groundwater
0.002 M
Groundwater
0.002 M
Groundwater
0.002 M
Groundwater
0.002 M
Groundwater
0.002 M
Groundwater
0.013 M
Groundwater
0.013 M
Groundwater
0.013 M
Groundwater
0.013 M
Groundwater
0.013 M
Groundwater
0.002 M
Groundwater
0.002 M
Groundwater
Soil and Mineral
Phase ID and
Information
How E Basalt
Flow E Basalt
Flow E Basalt
Flow E Basalt
Flow E Basalt
Flow E Basalt
Flow E Basalt
Flow E Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Smectite
Smectite
-
Ref2
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
D.23
-------
Cesium
Kd
-------
Cesium
Kd
(ml/g)
900
260
80
2200
1800
630
420
460
30
89
31
1
3
6
13
16
' 200
631
794
100
16
158
562
900
790
700
2
4
8
40
100
Clay
(wt%
)
Mica
(%)
pH
5
7
9
5
7
9
5
7
CEC"
(meq/100 g)
0.54
0.35
0.033
1.2
0.93
0.33
0.11
0.0067
0.0034
0.0032
0.00098
0.15849
0.19953
1.58489
1.77828
5.62341
7.94328
39.8107
63.0957
4.46684
6.30957
10
11.2202
SA1
(m2/g)
Aqueous Cs
(uM)
l.OOxlO'7
l.OOxlO'7
l.OOxlO-7
l.OOxlO-7
l.OOxlO-7
l.OOxlO-7
l.OOxlO'7
l.OOxlO-7
l.OOxlO-7
l.OOxlO-7
l.OOxlO'7
l.OOxlO-1
l.OOxlO-1
l.OOxlO-1
l.OOxlO-1
l.OOxlO-1
l.OOxlO-1
l.OOxlO-1
l.OOxlO-1
l.OOxlO"1
l.OOxlO-1
l.OOxlO-1
l.OOxlO-1
l.OOxlO-1
l.OOxlO-1
l.OOxlO-1
l.OOxlO-1
l.OOxlO'1
l.OOxlO"1
l.OOxlO-1
l.OOxlO'1
Background
Aqueous
Water ,
Water
Water
Water
Water
Water
Water
Water
Water
Water
Water
Soil and Mineral
Phase ID and
Information
lonada Granite
Rokka Granite
Limestone
Biotite
Chlorite
Hornblende
Grossular
Forsterite
K-feldspar
Albite
Quartz
Calcite
Apatite
Hematite
Orthoclase
Serpentine
Hornblende
Biotite
Muscovite
Gneiss
Diabase
Stripa Granite
Finsjo Granite
Biotite
Biotite
Biotite
Hematite
Hematite
Hematite
Hornblende
Hornblende
Ref2
8
8
8
8
8
8
8
8
8
8
8
9
9
9
9
9
9
9
9
9
9
9
9
9
9
9
9
9
9
9
9
D.25
-------
Cesium
Kd
(ml/g)
240
3
5
9
700
810
840
7
14
7
52000
26000
2500
2700
28600
94
7
12
2190
7610
620
Clay
(wt%
)
4
5
6
Mica
(%)
9
12
9
PH
9
5
7
9
5
7
9
5
7
9
7.7
8.2
7.9
CEC"
(meq/100 g)
127
20
11.2
127
20
11.2
SA1
(m2/g)
Aqueous Cs
(uM)
l.OOxlO'1
1.00x10"'
l.OOxlO-1
l.OOxlO-1
l.OOxlO"1
l.OOxlO-1
1.00x10-'
l.OOxlO'1
l.OOxlO'1
l.OOxlO-1
1.67xlO-6
1.67xlO-«
1.67X10-6
1.67X10-6
1.67xlO-«
1.67xlO-«
l.OOxlO-7
l.OOxlO'7
8.40xlO'3
8.40X10'3
8.40x10-'
Background
Aqueous
Deionized Water
Deionized Water
Deionized Water
O.lNNaCl
O.lNNaCl
O.lNNaCl
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Soil and Mineral
Phase ID and
Information
Hornblende
Magnetite
Magnetite
Magnetite
Muscovite
Muscovite
Muscovite
Orthoclase
Orthoclase
Orthoclase
Vermiculite
Illite
Kaolinite
Vermiculite
Dlite
Kaolinite
Hanford Vadose
Sediment
Hanford Vadose
Sediment
Sediment CGS-1
Sediment TBS-1
Sediment Trench-8
Ref2
9
9
9
9
9
9
9
9
9
9
10
10
10
10
10
10
11
11
12
12
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 efa/., 1977; 6 = Ames era/., 1982; 7 = Staunton, 1994; 8 = Akiba etal, 1989; 9 =Torstenfelt e*a/., 1982;
10 = Tamura, 1972; 11 = Routson etal., 1980; 12 = Seme etal, 1993.
D.26
-------
D.4.0 Data Set for Soils
Table D.I4 lists the available cesium Kd values identified for experiments conducted with only
soils.
Table D.14. Cesium Kd data set for soils only.
Cesium
Ka
(ml/g)
247
62
22
4100
1400
1100
280
95
120
130
130
150
160
72
79
75
98
83
33
Clay
(wt%)
Mica
(%)
15
15
15
15
15
15
3
3
3
3
3
4
pH
6.2
6.2
6.2
6.1
7.7
6.6
8.3
4
5.5
6.7
7
8.5
10.2
4
5.5
6.7
7
8.5
4
CEC»
(meq/100
g)
20
20
70
70
SA1
(m2/g)
130
130
60
60
Cs
(MM)
1.90xlO-2
1.42x10-'
5.94X10'1
l.OOxlO'1
l.OOxlO-1
l.OOxlO'1
l.OOxlO-1
4.20x10-*
4.20X10"4
4.20x10""
4.20x10-"
4.20X10-4
4.20x10-"
4.20x10^
4.20x10-"
4.20x10""
4.20x10^
4.20x10^
4.20x10-"
Aqueous
Phase
Gorleben
Groundwater
Water pH 6.1
Groundwater
pH7.7
Water pH 6.6
Groundwater
pH8.3
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
Soil ID
and Information
Gorleben Sediment
Gorleben Sediment
Gorleben Sediment
Takadate Loam
Takadate Loam
Hachinohe Loam
Hachinohe loam
Sav. River Site
Sediment
Sav. River Site
Sediment
Sav. River Site
Sediment
Sav. River Site
Sediment
Sav. River Site
Sediment
Sav. River Site
Sediment
4-Mile Creek Sediment
4-Mile Creek Sediment
4-Mile Creek
Sediment.
4-Mile Creek
Sediment.
4-Itdile Creek
Sediment
Pair Pond Soil
Ref.2
1
1
1
4
4
4
4
6
6
6
6
6
6
6
6
6
6
6
6
D.27
-------
Cesium
Ka
(ml/g)
37
40
39
50
27
25
26
26
38
39
88
92
93
85
94
101
88
89
90
84
101
22
31
37
40
78
7
12
Clay
(wt%)
Mica
(%)
4
4
4
4
2
2
2
2
2
2
4
4
4
4
4
4
5
5
5
5
5
2
2
2
2
2
pH
5.5
7
8.5
10.2
4
5.5
6.7
7
8.5
10.2
4
5.5
6.7
7
8.5
10.2
4
5.5
6.7
7
10.2
4
5.5
6.7
7
10.2
CEC<">
(meq/100
g)
SA1
(mVg)
Cs
(n»r>
4.20x10-"
4.20X10"4
4.20x10""
4.20x10-"
4.20x10-"
4.20x10-"
4.20X10"1
4.20X10"1
4.20X10"1
4.20x10-"
4.20X10"1
4.20x10-"
4.20x10^
4.20X10"1
4.20x10-"
4.20x10""
4.20x10-"
4.20x10-"
4.20x10-"
4.20x10^
4.20x10""
4.20x10""
4.20x10^
4.20x10-"
4.20x10-"
4.20x10""
l.OOxlO-7
l.OOxlO'7
Aqueous
Phase
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
0.005 MNa
Groundwater
Groundwater
Soil ID
and Information
Par Pond Soil
Par Pond Soil
Par Pond Soil
Par Pond Soil
Steel Creek Soil
Steel Creek Soil
Steel Creek Soil
Steel Creek Son
Steel Creek Soil
Steel Creek Soil
Lower 3 Runs Soil
Lower 3 Runs
Sediment
Lower 3 Runs
Sediment
Lower 3 Runs
Sediment
Lower 3 Runs
Sediment
Lower 3 Runs
Sediment
Pen Branch Soil
Pen Branch Soil
Pen Branch Soil
Pen Branch Soil
Pen Branch Soil
Upper 3 Runs Soil
Upper 3 Runs Soil
Upper 3 Runs Soil
Upper 3 Runs Soil
Upper 3 Runs Soil
Hanford Vadose
Sediment
Hanford Vadose
Sediment
Ref.2
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
6
8
8
D.28
-------
Cesium
K,
(ml/g)
3,000
4,800
3,100
3,000
2,190
7,610
620
Clay
(wt%)
4
5
6
Mica
(%)
6
7.5
8
5
9
12
9
pH
7.6
5.9
6.6
8
7.7
8.2
7.9
CEC<*>
(meq/100
g)
3
4.3
4.7
2.6
SA1
(m2/g)
8.6
12.2
14.7
6.6
Cs
(uM)
l.OOxlO'1
l.OOxlO'1
1. OOxlO'1
l.OOxlO-1
8.40xlO'3
8.40x1 0'3
8.40xlO'3
Aqueous
Phase
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Soil ID
and Information
Sediment A
Sediment B
Sediment C
Sediment D
Sediment CGS-1
Sediment TBS-1
Sediment Trench-8
Ref.2
10
10
10
10
11
11
11
1 CEC = Cation exchange capacity; SA = surface area.
2 1 = Lieser and Steinkopff, 1989; 4 = Konishi et al, 1988; 6 = Elprince et al, 1977; 8 = Routson et al., 1980; 10 = Legoux
etal, 1992; 11 = Seme etai, 1993.
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-------
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-------
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Environmental Pollutant Assessment System. PNL-7145, Pacific Northwest National
Laboratory, Richland, Washington.
Tamura, T. 1972. "Sorption Phenomena Significant in Radioactive-Waste Disposal." Am.
Assoc. Pet. Geol Mem., 18:318-33.
Torstenfelt, B. K. Andersson, and B. Allard. 1982. "Sorption of Strontium and Cesium on
Rocks and Minerals." Chemical Geology, 36:128-137.
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 National Laboratory, Los Alamos, New Mexico.
D.32
-------
-------
APPENDIX E
Partition Coefficients For Chromium(VI)
-------
-------
Appendix E
Partition Coefficients For Chromium(VT)
E.1.0 Background
The review of chromium K.J data obtained for a number of soils (summarized in Table E.I)
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 Kj) 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 IQ. Soils containing Mn oxides oxidize Cr(III) into Cr(VI) form thus resulting hi
lower Kd values. The relation between oxide/hydroxide contents of iron and manganese
and their effects on K^ have not been adequately quantified except for a few soils.
• The presence of competing anions reduce Cr(VT) adsorption. The inhibiting effect varies
in the order HPO|% H2PO; »SCt COl'/HCOi Cl% NO3-. 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 and Leckie, 1980; Griffin et al, 1977; Leckie et al, 1980; Rai et al, 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 et al. (1986) investigated the adsorption behavior
of Cr(VI) on amorphous iron oxide surfaces. The experiments were conducted with initial
concentrations of 5x10"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 (SO4: 2.5xlO"3 M, solution in equilibrium with 3.5xlO"3 atm CO2),
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 Kj 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
K^ values for these soils. The data showed that 3 of the 4 soils studied exhibited decreasing Kd
values with increasing pH. The Kj values for these soils were close to 1 ml/g at higher pH values
(>8). At lower pH values (about 4.5) the Kj 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 (K^ < 1 ml/g) for Cr(VI).
i
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 Kj values that ranged from 6 to 263 ml/g (Sheppard et al, 1987).
These measurements showed that coarse-textured soils tend to have lower Ka 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 Kj 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 K^ value. However,
removing iron oxide and hydroxide coatings resulted in a K.J 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 tune 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 IQ
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 K^ 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 Kj values or any correlative relationships between K
-------
ferric hydroxide. Therefore, observed removal of Cr(VI) from solution when contacted with
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)O4 Such complex geochemical behavior chromium in
soils implies that depending on the properties of a soil, the measured Kj 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 HPOl" and H2PO4 ions and to a very small extent by Cl~ and NO3' ions. The data
indicate that Cr(VI) adsorption was inhibited by anions in order of HPO|", HyPO^ » SO|"» Cl%
NO; (Leckie et al, 1980; MacNaughton, 1977; Rai et al, 1986; Rai et al, 1988; Stollenwerk and
Grove, 1985).
E.4
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£.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 et al. (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 K^
values are significantly higher at lower pH values and decline with increasing pH. Also, Kj
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 et al (1980) and Rai et al. (1988) showed
that iron oxides in the soil reduce Cr(VI) to Cr(HI) and precipitate Cr(III) as a (Fe,Cr)(OH)3
mineral. Also, studies conducted by Rai et 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 K7), Cr(VI) adsorption tends to be very low (very low Ka values)
irrespective of DCB extractable iron content. Similarly, soils which contain very low DCB
extractable iron, adsorb very little Cr(VI) (very low Kj values) irrespective of soil pH values.
Additionally, Cr(VI) adsorption studies show that the presence of competing anions such as
HPOJ-, H2PC>4, SO5% CO|% and HCOi will reduce the Ka 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 SO|% CO|% and HCOi (fixed through a single selected partial pressure
of CO2) concentrations (Tables E.4 and E.5). Among these competing anions, SO|" 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) Kj 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 864" concentrations increase from 0 to 2xlO'3 M (or 191.5
mg/1).
E.7
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Figure E.I. Variation of K,j for Cr(VI) as a function of pH and DCS 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 Kj 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
et al., 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 Kj 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 et al. (1988) on effects of competing anions on
Cr(VI) adsorption on Cecil/Pacolet soil.
CrCVI)1
pH
9.26
9.29
8.57
7.80
7.41
7.38
6.99
6.94
6.67
6.49
6.19
6.16
5.89
5.84
5.46
5.49
4.98
4.98
4.49
4.49
-logC
(mol/ra3)
3.05
3.05
3.11
3.30
3.44
3.46
3.66
3.65
3.79
3.79
3.99
3.94
4.08
4.06
4.19
4.21
4.33
4.32
4.52
4.39
-logS
(mol/kg)
5.66
5.88
5.34
5.00
4.89
4.88
4.81
4.81
4.78
4.78
4.75
4.75
4.74
4.74
4.73
4.73
4.72
4.72
4.71
4.72
Ka
(ml/g)
2
1
6
20
35
38
71
69
102
102
174
155
219
209
288
302
407
398
646
468
Cr(VI) + Sulfate1
pH
8.92
8.38
8.38
7.70
7.67
7.37
7.24
6.85
6.76
6.58
6.56
6.15
6.15
5.75
5.79
5.35
5.33
4.68
4.69
-logC
(mol/m3)
3.05
3.07
3.04
3.12
3.12
3.19
3.23
3.34
3.37
3.43
3.34
3.55
3.51
3.58
3.56
3.60
3.59
3.55
3.47
-logS
'mol/kg)
6.27
5.71
5.70
5.28
5.28
5.11
5.09
4.95
4.96
4.92
4.95
4.85
4.88
4.82
4.86
4.83
4.84
4.86
4.86
Ka
(ml/g)
1
2
2
7
7
12
14
24
26
32
25
50
43
58
51
59
57
49
41
Cr(VI) -t- Carbonate1
pH
9.62
9.15
9.01
7.92
7.95
7.53
7.52
7.19
7.31
7.22
6.99
6.70
6.68
5.84
6.08
5.12
5.12
4.76
4.75
4.33
4.34
-logC
(mol/m3)
3.05
3.05
3.06
3.06
3.06
3.08
3.07
3.12
3.10
3.12
3.13
3.22
3.21
3.65
3.54
4.11 ^
4.14
4.20
4.11
4.39
4.37
-logS
[mol/kg)
6.88
6.79
6.35
6.12
6.10
5.85
6.06
5.55
5.67
5.55
5.48
5.21
5.24
4.87
4.91
4.78
4.78
4.78
4.78
4.76
4.77
K<
(ml/g)
0
0
1
1
1
2
1
4
3
4
4
10
9
60
43
214
229
263
214
427
398
1 Cr(VI) concentration: 10"* M, Sulfate Concentration: 1Q-27M, CO2 : 10-'6atm.
E.ll
-------
Table E.5. Data from Rai et al. (1988) on effects of competing
anions on Cr(VI) adsorption on Kenoma soil.
CrifVI) + Sulfate + Carbonate
CrfVI) concentration: 10"6 M, Sulfate Concentration: 10'2
E.12
-------
-------
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 K^.
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 et al. (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
-------
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en
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en
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I
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a
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CO
3
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O O O
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111
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.11
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II
00
F.5
-------
5500
4400
3300
2200
1100
0
s
Figure F.I. Correlative relationship between Kj and pH.
F.6
-------
Figure F.2. Variation of Kj 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 (ug/1)
0.1-0.9
1.0-9.9
10 - 99.9
100 - 200
Kd (ml/g)
Minimum
Maximum
Minimum
Maximum
Minimum
Maximum
Minimum
Maximum
Soil pH
4.0 - 6.3
940
8,650
420
4,000
190
1,850
150
860
6.4 - 8.7
4,360
23,270
1,950
10,760
900
4,970
710
2,300
8.8 - 11.0
11,520
44,580
5,160
20,620
2,380
9,530
1,880
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: IE. XAFS Determination of Pb(II) and Pb(n>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.
i
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." InTheBiogeochemistryofLeadinthe
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, 11:3 59-364.
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. Grouse. 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. Zirnmie, and C. O. Riggs
(eds.), pp 221-242. ASTM STP 746. American Society of Testing Materials. Washington,
D.C.
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. Altaian. 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 and Migration 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 and B. 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 plutonkim 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 hi soils and the measured adsorption of plutonium has not
been quantified.
Because plutonium in nature can exist in multiple oxidation states (EL, 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.
G.2
-------
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 etal. (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.
j
Even though the adsorption behavior of plutonium on soil minerals such as glauconite (Evans,
1956), montmorillonite (Billon, 1982; Bondietti etal, 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 etal., 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 [1.7x10"* 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
Kj 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
G.3
-------
bicarbonate concentrations (8.2 x 10"6 to 2.9 x lO^M)] resulted in distribution coefficients, K.J,
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(TV) sorption on alumina was lowered to
about 30 ml/g at 0.5 M bicarbonate (Charyulu et 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 PuCyxH2O 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(TV) 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 PuCyxH2O. The equilibrating solutions used in these
experiments contained salts such as NaNO3, NaCl, N^SO* and NaHCO3 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 NaNO3 electrolyte medium by Sanchez et
al. (1985) indicated isotherms typical of metal and/or metal-like complex specie adsorption on
substrate (Benjamin and Leckie, 1981). This indicated that Pu(TV) 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
G.4
-------
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(TV) and Pu(V) adsorption data obtained in 0.1 M NaNO3 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(TV) or Pu(V) adsorption occurred on
goethite when the total carbonate concentration approached 1,000 meq/1 (0.5 M CO3). 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(W) 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(TV) 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
G.5
-------
highly alkaline conditions. These data are unreliable because initial plutonium concentration of
6.8xlO~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(HI), Pu(IV), and Pu(VT) 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 IxlO"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 1.68xl05 ml/g. Experiments conducted by Dahlman et al (1976) also showed
exceedingly high Kd value (3xlOs ml/g) for Pu(TV) 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 lO^M) 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(TV) fixed at about 5xlO"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.
G.6
-------
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(TV) 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"n 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(VT) upon adsorption was reduced to the tetravalent state. In these
experiments, the Kd data obtained at lower initial concentrations (~lxlO'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 2.1x10'8 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.
G.7
-------
A number of studies indicate that Kd values for plutonium adsorption on river, oceanic, and lake
sediments range from about IxlO3 to IxlO6 ml/g. Duursma and coworkers calculated that Kd for
marine sediments was about IxlO4 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 (in, 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 4xlOsml/g, and for oceanic particulates ranged from
IxlO5 to 4xl05 ml/g.
G.2.0 Data Set for Soils
The most detailed data set on plutonium Kd measurements! were obtained by Glover et 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 lO^M) 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.I).
G.3.0 Approach and Regression Models
The most detailed data set on plutonium Kd measurements; were obtained by Glover et 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 lO^M) 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.
G.8
-------
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 (1C as percent CaCO3), organic carbon content
(OC as wt.%), and the clay content (CLAY as wt.%).
G.9
-------
Table G.I. 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
CO-A
CO-B
CO-C
ID-A
ID-B
ID-C
ID-D
WA-A
WA-B
SC
NY
NM
AR-A
AR-B
AR-C
IL
K_
(ml/g)
2,200
200
1,900
1,700
320
690
2,100
100
430
280
810
100
710
80
430
230
PH
5.7
5.6
7.9
7.8
8.3
8.0
7.5
8.0
8.2
5.4
5.4
6.4
6.2
4.8
2.3
3.6
CEC1
(meq/lOOg)
20.0
17.5
29.6
15.5
13.8
8.2
17.5
6.4
5.8
2.9
16.0
7.0
34.4
3.8
16.2
17.4
EC1
(mmhos/cm)
3.6
0.4
0.4
0.5
0.8
1.0
1.2
0.9
0.4
0.4
1.2
1.7
0.5
0.4
0.3
0.5
DCARB1
(meq/I)
5.97
0.97
1.98
2.71
2.51
2.52
4.90
2.60
2.30
0.50
1.40
2.80
0.10
0.10
0.10
0.10
1C %'
CaCO3
0.4
0.3
2.4
17.2
7.9
5.2
0.0
0.6
0.0
0.2
s 0.0
0.2
0.9
0.7
0.6
0.7
OC1
(%
mass)
2.4
3.4
0.7
0.8
0.2
0.3
0.1
0.3
0.1
0.7
2.7
0.7
3.2
0.6
2.3
3.6
CLAY1
(%
mass)
36
22
64
34
32
23
3
14
14
20
36
18
56
9
37
16
1 CEC: Cation exchange capacity; EC: Electrical conductivity; DCARB: Dissolved carbonate;
1C: Inorganic carbon; OC: Organic carbon; CLAY: Soil clay content.
G.10
-------
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.
i
The scatter matrix (Figure G.I) 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.
G.ll
-------
KD
PH
CEC
EC
DCARB
1C
OCOG °o o»
oc
CLAY
Figure G.I. Scatter plot matrix of soil properties and the distribution coefficient
plutonium.
of
The variable selection with and without an intercept indicated that the 2 most significant variables
for reliably forecasting the Kd values were the concentrations of dissolved carbonate (DCARB)
and the clay content (CLAY) of soils (Table G.2). Using these 2 independent variables, several
forms of polynomial regression models and a piecewise regression model with a breakpoint were
generated. The results showed that the best regression model among all the models tested was
the piecewise regression model. The relationship between the Kd values and the 2 independent
variables (CLAY and DCARB) is shown as a 3-dimensional surface (Figure G.2). This graph
illustrates that the highest Kd values are encountered under conditions of high clay content and
dissolved carbonate concentrations. In contrast, the low Kd values are encountered in soils
containing low clay content and low dissolved carbonate concentrations.
Using the piecewise regression model, a look-up table (Table G.3) was created for ranges of clay
content and soluble carbonate values which are typically encountered in soils.
G.12
-------
Table G.2. Regression models for plutonium adsorption.
Model Type
Linear Regression
Forward Stepwise
linear Regression
Forward Stepwise
Linear Regression
Backward Stepwise
Linear Regression
Backward Stepwise
Piecewise Linear
Regression
Polynomial
Polynomial
Polynomial
Polynomial
Forecasting Equation
KJ - 284.6 (DCARB) + 27.8 (CLAY) - 594.2
Kj = 488.3 (DCARB) + 29.9 (CLAY) - 1 19. 1 (pH) - 356.8 (EC)
Kj - 284.6 (DCARB) + 27.8 (CLAY) - 594.2
Kj = 351.4 (DCARB)
Kj = 25.7 (DCARB) + 12. 14 (CLAY) + 2.41 for Kd values <767.5
Kj = 286.0 (DCARB) + 21.3(CLAY) - 81.2 for Kj values >767.5
Kj = -156.0 (DCARB) + 15.2 (CLAY) +16.1 (DCARB)2 - 0.04 (CLAY)2 + 11.3 (DCARB)(CLAY) - 87.0
Ki = -171.1(DCARB) + 10.5 (CLAY) -l-17.2(DCARB)2 + 0.02 (CLAY)2 + 11.6 (DCARB)(CLAY)
Kj = -106.1(DCARB) -M 1.2 (CLAY) + 12.5 (DCARBXCLAY) - 72.4
Kj = -137.9 (DCARB) + 9.3 (CLAY) + 13.4 (DCARBXCLAY)
R2
0.7305
0.8930
0.7305
0.7113
0.9730
0.9222
0.9219
0.9194
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, (ml/e)
Minimum
Maximum
Clay Content (wt.%)
0-30
Soluble Carbonate
(meq/1)
0.1-2
5
420
3-4
80
470
5-6
130
520
31-50
Soluble Carbonate
(meq/1)
0.1-2
380
1,560
3-4
1,440
2,130
5-6
2,010
2,700
51-70
Soluble Carbonate
(meq/1)
0.1 • 2
620
1,980
3-4
1,860
2,550
5-6
2,440
3,130
G.13
-------
Figure G.2. Variation of Kj for plutonium as a function of clay content and
dissolved carbonate concentrations.
G.14
-------
G.4.0 References
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G.15
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Waste Components on the Solubility and Sorption of Cobalt, Strontium, Neptunium,
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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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Glover, P. A., F. J. Miner, and W. O. Polzer. 1976. "Plutonium and Americium Behavior in the
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G.16
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Nelson, D. M.5 R P. Larson, and W. R. Penrose. 1987. "Chemical Speciation of Plutonium in
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G.18
-------
-------
APPENDIX H
Partition Coefficients For Strontium
-------
HgMHHMHMMMHHMMMMHMWM|
-------
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 hi
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.l. The 63 strontium Kd values in this data set ranged from
H.2
-------
1.6 ml/g for a measurement made on a sandy soil dominated by quartz (Lieser et al, 1986) to
10,200 ml/g for a measurement made on a tuff1 soil collected at Yucca Mountain, Nevada
(Sample YM-38; Vine et al., 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.
Mean
Standard Error
Median
Mode
Standard Deviation
Kurtosis
Minimum
Maximum
Number of
Observations
SrK,
(ml/g)
355
183
15
21
1,458
34
1.6
10,200
63
Clay
Content
(wt.%)
7.1
1.1
5
5
7.85
10.7
0.5
42.4
48
pH
6.8
0.21
6.7
6.2
1.35
-0.5
3.6
9.2
42
CEC
(meq/100 g)
4.97
1.21
0.9
2
9.66
11.6
0.05
54
63
Surface
Area
(«n2/g)
1.4
0
1.4
1.4
0.00
-3
1.4
1.4
7.00
Ca
(mg/1)
56
23
0
0
134
3.4
0.00
400
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 -mih 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 Ka 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; Seme 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 Kd
Clay Content
pH
CEC
Surface Area
Ca Cone.
Strontium
1C,
1.00
0.821
0.28
0.841
0.00
-0.17
Clay
Content
1.00
0.03
0.911
-1.00
0.00
pH
1.00
0.281
0.00
-0.20
CEC
1.00
l.OO1
0.03
Surface
Area
1.00
—
Ca Cone.
1.00
1 Correlation coefficients significant at or less than the 5% level of probability (P ^ 0.05).
H.4
-------
^58
$f
105
104
1000
100
10
1
n i
aio.(
1 1 '"I 1 1
»
r °i
0 °°
o o° « ® °o o "
o fto s
00
Figure H.I. 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 K^ 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 Kj 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.I is presented as Equation 1 hi 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 K
-------
values at low CEC values, 2 approaches were evaluated. First, the data in Figure H.I 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= 2xlO"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 hi 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.
repelled by the negative charge of permanently charged minerals.
H.6
-------
^
I
•o
Nj
00
140
120
100
80
60
40
i i i i i i i
— -
o
-
_
0
0
-
-o o °
20 || § ° *> o
n IKr^ ' i0i i i i i
0 2 4 6 8 10 12 14
CEC(meq/100g)
Figure H.2. Relation between strontium Kd values for soils with
CEC values less than 15 meq/100 g.
1000
-a 100
^
£ 10
1
: I ill:
: 0 :
O ;
0
b° =
• i°l° •
08
i i i i
* 0 10 20 30 40 50
Clay (%, wt.)
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).
#
1
2
3
4
5
6
7
8
9
10
11
12
Equation
Kd = -272+126(CEC)
Kd=114(CEC)
Kd=10.0 + 4.05(CEC)
Kd = 5.85(CEC)
Kd=-42+14(CEC) +
2.33(pH)
Ka = 3.53(CEC) +
1.67(pH)
CEC = -4.45 +
0.70(clay) + 0.60(pH)
CEC = 0.40(clay) +
0.19(pH)
K,3 = -108+10.5(clay) +
11.2(pH)
Kd = 3.54(clay) +
1.67(pH)
Clay = 3.36 +
1.12(CEC)
Clay = 1.34(CEC)
n2
63
63
57
57
27
25
27
25
27
25
48
48
Data
Range3
All
All
CEC<15
CEC<15
All
CEC<15
All
CEC<15
All
CEC<15
All
All
95% Confidence Limits1
Intercept
Lower
-501
—
3.32
—
-176
—
-10.6
._
-270
—
2.30
—
Upper
-43
—
16.6
—
91
—
1.67
.„
53.3
—
4.41
—
Slope First
Independent
Parameter
Lower
105
95
2.13
4.25
11.3
0.62
0.59
0.24
7.32
0.62
0.97
1.16
Upper
147
134
5.96
7.44
18.3
6.46
0.82
0.56
13.6
6.46
1.26
1.51
Slope Second
Independent
Parameter
Lower
—
—
—
—
-17.7
-0.50
-0.30
-0.01
-12.5
-0.50
—
—
Upper
—
—
—
—
22.4
3.85
1.50
0.40
34.9
3.85
—
—
R24
0.70
0.67
0.25
0.12
0.77
0.34
0.86
0.55
0.67
0.34
0.84
0.69
F Value5
IxlO'17
2xlO'16
9xlO'5
7xlO'3
SxlO"8
9xlO'3
4x10-"
IxlO-4
2X10"6
9xlO'3
IxlO'19
2xlO'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
-------
/oB
^
c
V^-X
13
i.
00
105
I. \J
10"
1000
100
10
1
1 ,
! 1 1 1 1 1 1
r o -
o
o :
r i
O i
O •
1 0 i
CO o :
Oo o oo<*> g«&o° :
0 o 8 o^jj o ° i
i i i ® i ii
54567891
pH
0
Figure H.4. Relation between strontium KJ values and soil pH.
H.2.4 Approach
Two strontium K
-------
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.]
Kd (ml/g)
Minimum
Maximum
CEC (meq/100 g)
3
pH
<5
1
40
5-8
2
60
8-10
3
120
3-10
pH
<5
10
150
5-8
15
200
8-10
20
300
10 - 50
pH
<5
100
1,500
5-8
200
1,600
8-10
300
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.]
Kd (ml/g)
Minimum
Maximum
Clay Content (wt.%)
<4%
pH
<5
1
40
5-8
2
60
8-10
3
120
4 - 20%
pH
<5
10
150
5-8
15
200
8-10
20
300
20 - 60%
pH
<5
100
1,500
5-8
200
1,600
8-10
300
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
12
12
11
11
Y-Intercept
—
—
3.36
3.36
Slope
1.34
1.34
1.1.2
1.12
CEC
(meq/100 g)
3
15
15
50
Clay Content
(%)
4
20
20
59
1 Number of equation in Table H.3.
H.11
-------
H.3.0 IQ 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)
21
19
22
26
24
30
43
21.4
25
12.7
7.9
15.6
9.4
7.6
6.4
7.7
28.1
Clay
Content
(%)
0.8
0.8
0.8
0.8
0.8
0.8
0.8
5
5
5
5
5
5
5
5
5
5
pH
5.2
5.6
6.2
6.45
6.6
8.4
9.2
CEC
(meq/
100 g)
0.9
0.9
0.9
0.9
0.9
0.9
0.9
0.47
0.83
0.39
0.46
0.81
0.21
0.25
0.24
0.26
0.76
Surface
Area
(mVg)
1.4
1.4
1.4
1.4
1.4
1.4
1.4
[Ca]
ppm
0
0
0
0
0
0
0
[Sr]
*
*
*
*
*
*
*
Background
Solution
NaClO4
NaC104
NaClO4
NaClO4
NaClO4
NaClO4
NaClO4
(Sroundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
SoU
ID
Soil A
Soil A
Soil A
Soil A
Soil A
Soil A
Soil A
Reference1, Comments
1, * = 4.4e2Bq/ml85-Srin
2.4xlO-8MSrCl2
1, * = 4.4e2Bq/ml85-Srin
2.4x10-" MSrCl2
1, * = 4.4e2Bq/ml85-Srin
2.4xlO-8MSrCl2
15 *=4.4e2Bq/ml85-Srin
2.4xlO-8MSrCl2
1, * = 4.4e2Bq/ml85-Srin
2.4xlO-8MSrCl2
1, *=4.4e2Bq/ml85-Srin
2.4x1 0-8MSrCl2
1, *=4.4e2Bq/ml85-Srin
2.4xlO-8MSrCl2
2
2, CEC was estimated by
adding exch. Ca,Mg,K
2,GW = 7.4Ca,1.7Mg,
2.2Na,5.6Cl, 18ppmSO4
2, Aquifer sediments
Chalk River Nat'l Lab,
Ottawa, Canada
2, Described as sand texture
2, Assumed 5% clay, mean
[clay] in sandy soils
2
2
2
H.12
-------
SrK,,
(ml/g)
7.63
11.4
20.1
13
9.8
11
13
7.8
3.8
3
2.5
4
15
21
24
3
4.5
53.
5.7
3.5
4.6
5.8
6.1
Clay
Content
(%)
5
5
5
5
5
5
5
5
5
5
5
10
10
10
10
10
10
10
10
pH
4
5
6
7.4
3.6
5.2
6.8
7.9
5.2
5.6
5.8
5.9
CEC
(meq/
100 g)
0.26
0.41
0.44
0.25
0.29
0.22
0.39
0.2
0.1
0.1
0.13
5.5
5.5
5.5
5.5
5.5
5.5
5.5
5.5
2
2
2
2
Surface
Area
(mVg)
[Ca]
ppm
0
0
0
0
400
400
400
400
0
0
0
0
[Sr]
IxlO-'M
IxlO-SM
lxlO-*M
IxlO-SM
IxlO-'M
1x10-^
1x10-%!
IxlO-SM
lxlO-'°M
lxlO-10M
lxlO-'°M
lxlO-10M
Background
Solution
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
O.OlMNaCl
O.OlMNaCl
O.OlMNaCl
O.OlMNaCl
O.OlMCaCl
O.OlMCaCl
O.OlMCaCl
O.OlMCaCl
NaOH/HCl
NaOH/HCl
NaOH/HCl
NaOH/HCl
Soil
m
Puye
soil-Na
Puye
soil-Na
Puye
soil-Na
Puye
soil-Na
Puye
soil-Ca
Puye
soil-Ca
Puye
soil-Ca
Puye
soil-Ca
Hanford
soil
Hanford
soil
Hanford
soil
Hanford
soil
Reference *, Comments
2
2
2
2
2
2
2
2
2
2
2
3
3, Noncalcareous soils
3
3
3
3
3
3
4
4, Carbonate system
4
4
H.13
-------
SrK,,
(ml/g)
8.3
17
21
27
47
81
19.1
21.5
232
48.5
10,200
2,500
3,790
3,820
1.6
2.6
3.4
4.6
6.7
400
Clay
Content
(%)
4
6
5
0.5
3
5
8
13
42.4
PH
6
7.4
7.6
7.8
8.4
9.1
7.66
7.87
8.17
8.24
8.17
8.13
8.24
8.24
6.2
6.2
6.2
6.2
6.2
7.2
CEC
(meq/
100 g)
2
2
2
2
2
2
10.4
5.9
4.57
3
54
21
27
27
0.05
0.3
0.5
0.8
1.3
34
Surface
Area
(nrVg)
[Ca]
ppm
0
0
0
0
0
0
129
58.5
35.1
0
[Sr]
lxlO-IOM
lxlO-1
-------
SrlQ
(ml/g)
135
600
70
Clay
Content
(%)
26.9
33.5
3.5
pH
8.3
6.5
8.3
CEC
(meq/
100 g)
13.6
26.3
5.8
Surface
Area
(m2/g)
[Ca]
ppm
0
0
0
[Sr]
Background
Solution
Water
Water
Water
Soil
ED
Bowdoin
Soil
Hall soil
Composite
Soil
Reference1, Comments
8, soil from Montana
8, soil from Nebraska
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 et al, 1980; 7 = Lieserand Steinkopff, 1989; 8 = McHenry, 1958
H.15
-------
H.4.0 IQ 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)
21
19
22
26
24
30
43
0
290
140
17
37
8
6
Clay
Conten
t(%)
0.8
0.8
0.8
0.8
0.8
0.8
0.8
pH
5.2
5.6
6.2
6.45
6.6
8.4
9.2
5.5
5.5
5.5
5.5
5.5
5.5
5.5
CEC
(meq/
100 g)
0.9
0.9
0.9
0.9
0.9
0.9
0.9
3.3
3.6
0.6
1.9
0.5
0.5 •
Surface
Area
(mVg)
1.4
1.4
1.4
1.4
1.4
1.4
1.4
26.4
43.9
1.4
2.2
0.7
[Ca]
(ppm)
0
0
0
0
0
0
0
0
0
0
0
0
0
[Sr]
*
*
*
*
*
*
*
*
*
*
*
*
*
#
Background
Solution
NaClO,
NaC104
NaClO,
NaClO,,
NaClO4
NaClO4
NaClO4
Soil ID
SoU A
Soil A
SoU A
Soil A
SoU A
Soil A
Soil A
Quartz
Kaolinite
Halloysite
Chlorite
Sericite
Oligoclase
Hornblend
Reference1
and Comments
1, Ohnuki, 1994
I,*=4.4xl02 Bq/ml 85-
Srin2.4xlO-8MSrCl2
I,* = 4.4xl02 Bq/ml 85-
Sr in 2.4xlO-sM SrCl2
I,*=4.4xl02 Bq/ml 85-
Srin2.4xlO-8MSrCl2
1, *=4.4xl02 Bq/ml 85-
Srin2.4xlO-sMSrCl2
1, * =4.4xl02 Bq/ml 85-
Srin2.4xlO-8MSrCl2
1; *=4.4xl02 Bq/ml 85-
Sr in 2.4xlO'sM SrCl2
1, *=4.4xl02 Bq/ml 85-
Srin2.4xlO-sMSrCl2
1, * = 4.4xl02 Bq/ml 85-
Srin2.4xlO-*MSrCl2
1, * = 4.4xl02 Bq/ml 85-
Srin2.4xlO-8MSrCl2
1, *= 4.4x1 02 Bq/ml 85-
Srin2.4xlO'8MSrCl2
1, *= 4.4x1 02 Bq/ml 85-
Srm2.4xlO-sMSrCl2
1, * = 4.4xl02 Bq/ml 85-
Srin2.4xlO-sMSrCl2
1, * = 4.4xl02 Bq/ml 85-
Srin2.4xlO-8MSrCl2
H.16
-------
SrIQ
(ml/g)
16
110
7.7
9.9
12.6
13.7
10.1
15.8
13.8
11
142.
6
7.5
6.9
8.3
8
6.7
6.8
4.9
5.1
8.5
8.8
5.6
5.3
7.2
5.1
6.5
Clay
Conten
t(%)
pH
5.5
5.5
5.8
6.1
6.1
5.8
6
5.8
5.8
5.9
5.6
5.8
5.9
5.9
6.1
6.2
6.2
6.2
6.2
6.2
6.2
6.2
6.3
6.4
6.4
6.3
6.4
CEC
(meq/
100 g)
0.7
8.5
Surface
Area
(mVg)
19.3
[Ca]
(ppm)
0
0
24
25
23
22
24
21
27
21
21
24
21
17
24
21
28
18
19
17
18
20
16
18
18
17
[Sr]
*
*
113uCi/l
105 jxCi/l
105 n(M
123 jiCi/1
99]iCi/l
143 nCi/1
113 uCi/1
114uCi/l
124 fiCifl
115 uCi/1
117uCi/l
108 uCi/1
68 (iCi/1
71 uCi/1
72p.Ci/L
84uCi/l
84|iCi/l
87uCi/l
88uCi/l
90nCi/l
77uCi/l
79(iCi/l
65nCi/l
72uCiyi
75uCi/l
Background
Solution
Groundwater
Groundwater
Groimdwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Soil ID
Pyroxene
Mn02
AA45/1
AA45/3
AA45/4
AA45/5
AA45/7
AA38/1
AA38/2
AA38/3
AA38/4
AA38/5
AA38/6
AA38/8
AA27/1
AA27/2
AA27/3
AA27/4
AA27/5
AA27/6
AA27/7
AA27/8
AA34/1
AA34/2
AA34/3
AA34/4
AA34/5
Reference1
and Comments
U*=4.4xl02Bq/ml85-
Sr in 2.4xlO'8M SrCl2
1, *= 4.4x1 (FBq/ml 85-
Sr in 2.4x1 0-^SrClj
2 Jackson and Inch, 1989
2,Ka =-.38Ca + 0.82.r2
= 0.19
2, Ca not important to Sr
IM
2
2
2
2
2
2
2
2
2
2
2
2
2
2
2
2
2
2
2
2
2
2
H.17
-------
SrKfl
(ml/g)
6
6.5
7.6
21.4
25
12.7
7.9
15.6
9.4
7.6
6.4
7.7
28.1
7.63
11.4
20.1
13
9.8
11
13
7.8
3.8
3
2.5
4
15
Clay
Conten
t(%)
10
10
PH
6.2
6.2
6.2
4
5
CEC
(meq/
100 g)
0.47
0.83
0.39
0.46
0.81
0.21
0.25
0.24
0.26
0.76
0.26
0.41
0.44
0.25
0.29
0.22
0.39
0.2
0.1
0.1
0.13
5.5 •
5.5
Surface
Area
(nrVg)
[Ca]
(ppm)
14
15
17
0
0
[Sr]
79p.Ci/l
107 uCi/1
107 uCi/l
IxlO-'M
IxlO-'M
Background
Solution
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
.OlMNaCl
.OlMNaCl
Soil ID
AA34/6
AA34/7
AA34/8
Puye
soil-Na
Reference1
and Comments
2
2
2
3 Patterson and Spoel,
1981
3, CEC was
approximated by adding
exch. Ca,Mg,K
3, Groundwater =7.4
ppm Ca, 1 .7 ppm Mg, 2.2
ppm Na, 5.6 ppm Cl, 18
ppm SCX,
3
3
3
3
3
3
3
3
3
3
3
3
3
3
3
3
3
3
4
4, Noncalcareous soils
H.18
-------
SrKa
(ml/g)
21
24
3
4.5
5.2
5.7
7.2
12.7
14.9
12.9
25.1
40.6
48.6
35
39.2
252
16.4
10.3
8.2
7.6
7.8
112
10.5
3.7
3.5
4.6
Clay
Conten
t(%)
10
10
10
10
10
10
pH
6
7.4
3.6
5.2
6.8
7.9
3
5
7
9
11
5.2
5.6
CEC
(meq/
100 g)
5.5
5.5
5.5
5.5
5.5
5.5
2 .
2
Surface
Area
(mVg)
0.98
0.96
0.88
0.8
0.73
0.39
0.36
0.32
0.25
0.51
0.38
0.34
0.34
[Ca]
(ppm)
0
0
400
400
400
400
0
0
0
0
0
0
0
[Sr]
lxW-*M
IXIO-SM
IxlO-^
1x10-^
IxlO'8!^
IxlO^E
0.1 ppm
0.1 ppm
0.1 ppm
0.1 ppm
0.1 ppm
lxlO-10M
lxlO-'°M
Background
Solution
.OlMNaCl
.OlMNaCl
.01M CaCl2
.OlMCaCk
.01MCaCl2
.01MCaCl2
2,000 ppm
Na
2,000 ppm
Na
2,000 ppm
Na
2,000 ppm
Na
2,000 ppm
Na
NaOH/HCl
NaOH/HCl
Soil ID
Puye
soE-Ca
Hanford Soil
Hanford Soil
Hanford Soil
Hanford Soil
Hanford Soil
C-27
C-27
C-97
C-55
C-81
C-62
C-71
C-85
C-77
MK-4
TK3
RJK2
NK2
Hanford soil
Hanford soil
Reference1
and Comments
4
4
4
4
4
4
5
5
5
5
5
6
6
6
6
6
6
6
6
6
6
6
6
6
7
7
H.19
-------
SrlQ
(ml/g)
5.8
6.1
8.3
17
21
27
47
81
140
160
1500
1100
1800
950
550
260
19.1
21.5
23.2
48.5
10200
2500
3790
Clay
Conten
t(%)
70
70
70
70
10
10
10
10
4
6
5
0
0
0
0
pH
5.8
5.9
6
7.4
7.6
7.8
8.4
9.1
2.4
2.4
9.3
9.3
6.1
8
6.5
8.2
7.66
7.87
8.17
8.24
8.17
8.13
8.24
CEC
(meq/
100 g)
2
2
2
2
2
2
2
2
10.4
5.9
4.57
3
54
21
27
Surface
Area
(«n2/g)
70
70
70
70
130
130
60
60
[Ca]
(ppm)
0
0
0
0
0
0
0
0
0
0.
0
0
129
58.5
35.1
[Sr]
lxlO-10M
lxlO-10M
lxlO-'°M
lxlO-'°M
lxlO-IOM
lxlO-10M
lxlO-10M
lxlO-10M
IxlO^
IxlO-'M
IxlO-^
IxlO'^
IxlO-^
IxlO-8]^
IxlO"8!^
IXIO-SM
100 nCifl
100 fiCi/1
100 nCi/1
3.8x10-^
3.8x10-^
S.SxlO-8!^
S.SxlO^M
Background
Solution
NaOH/HCl
NaOH/HCl
NaOH/HCl
NaOH/HCl
NaOH/HCl
NaOH/HCl
NaOH/HCl
NaOH/HCl
Water
Groundwater
Water
Groundwater
Water
Groundwater
Water
Groundwater
Hanford
loroundwater
Hanford
Groundwater
Hanford
iGroundwater
Yucca
iGroundwater
Yucca
Groundwater
Yucca
Groundwater
Yucca
Groundwater
Soil ID
Hanford soil
Hanford soil
Hanford soil
Hanford soil
Hanford soil
Hanford soil
Hanford soil
Hanford soil
Bentonite
Bentonite
Bentonite
Bentonite
Takadate Loam
Takadate Loam
Hacbinohe
Loam
Hachinohe
Loam
cgs-1
trench-8
tbs-1
YM-22
YM-38
YM48
YM-49
Reference1
and Comments
7
7
7
7
7
7
7
7
8
8
8
8
8, hydrohalloysite=10%,
70% silt
8, hydrohalloysite=10%,
70% silt
8, hydrohalloysite = 10%,
90% silt
8, hydrohalloysite = 10%,
90% silt
9
9, Groundwater pH = 8.3
9
10, Los Alamos, New
Mexico
10, Yucca Mt tuff
sediments
10, Approximate initial
pH, final pH are
presented
10, Final pH 8.1- 8.5
H.20
-------
SrK,
(ml/g)
3820
27000
4850
85
17.7
385
149
25000
530
71,000
1.6
2.6
3.4
4.6
6.7
17,000
150
780
95
440
39
Clay
Content
t(%)
0
0
0
0
0
0
0
0.5
3
5
8
13
PH
8.24
8.4
8.63
8.25
8.5
8.39
8.45
12
12
12
6.2
6.2
6.2
6.2
6.2
CEC
(meq/
100 g)
27
0.05
0.3
0.5
0.8
1.3
97
3.4
2.4
1.9
1.9
1.2
Surface
Area
(m2/g)
31
31
8
8
105
105
[Ca]
(ppm)
10
50
10
50
10
50
[Sr]
S.SxlO-'M
S-SxlO-'M
3.8x10-^
S.SXIO-SM
S.SxlO'8!^
S.SxlO-SM
S.SXIO-SM
10 nCi/ml
10 nCi/ml
10 nCi/ml
10x10"^
10xlO-«M
lOxlO"6]^
lOxlO-^I
10xlO-«M
lxlO-'°M
lxlO'10M
lxlO-10M
lxlO-'°M
lxlO-10M
lxlO-I(>M
Background
Solution
Yucca
Groundwater
Yucca
Groundwater
Yucca
Groundwater
Yucca
Groundwater
Yucca
Groundwater
Yucca
Groundwater
Yucca
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Soil ID
YM-50
JA-18
JA-19
JA-32
JA-33
JA-37
JA-38
kaolinite
chlorite
FeOOH
Sediments
Sediments
Sediments
Sediments
Sediments
Ohyatuff
Pyrophyllite
Sandstone
Shale
Augite
Andesite
Plagiorhyolite
Reference1
and Comments
10, Sediments = 106-500
um fractions
10
10
10
10
10
10
13
13
13
14
14, Added Kaolinite to
sand
14, CEC estimated based
on kaolinite = 10
meq/100 g
14
14
14, Akiba and
Hashimoto, 1990
14, log Kd = log CEC +
constant for trace [Sr]
14, pH not held constant,
ranged from 6 to 9.
14, Ig solid:50ml
sol'n,centrifuged,32-
60mesh
14,CECofCsandKd of
Sr
14
_
H.21
-------
SrK,,
(ml/g)
380
50
82
22
1.3
2,000
140
40
20
71
150
0.92
14
30
3
23
400
135
600
70
2.4
4.7
6
2.3
5.5
4.8
Clay
Conten
t(%)
42.4
26.9
33.5
3.5
pH
7.2
8.3
6.5
8.3
4
5
7
4
5
7
CEC
(meq/
100 g)
0.75
0.57
0.54
0.35
0.033
2
0.93
0.36
0.33
0.11
0.07
0.067
0.034
0.032
0.022
0.0098
34
13.6
26.3
5.8
Surface
Area
(m2/g)
[Ca]
(ppm)
0
0
0
0
[Sr]
lxlO-10M
lxlO-:°M
lxlO'10M
lxlO-10M
lxlO-IOM
lxlO-IOM
lxlO-10M
lxlO-10M
lxlO-10M
lxlO-10M
lxlO-10M
lxlO-IOM
lxlO-IOM
lxlO'10M
lxlO-IOM
lxlO-IOM
Background
Solution
Water
Water
Water
Water
Groundwater
SoU ID
Olivine Basalt
Vitric Massive
Tuff
Inada granite
Rokko Granite
Limestone
Muscovite
Chlorite
Hedenbergite
Hornblende
Grossular
Microcline
Forsterite
K-Feldspar
Albite
Epidote
Quartz
Ringhold Soil
Bowdoin Soil
Hall Soil
Composite Soil
Eolian Sand
Eolian Sand
Eolian Sand
Mol White
Sand
Mol White
Sand
Mol White
Sand
Reference1
and Comments
14
14
14
14
14
14
14
14
14
14
14
14
14
14
14
14
1 1 , Soil from Richland
WA
1 1 , from Montana
1 1 , from Nebraska
1 1, from Hanford Site
12
12, Belgian soils
12, Composition of
Groundwater was not
given
12, Compared static vs.
dynamic Kd
12
12
H.22
-------
SrK<
(nU/g)
2.6
5.3
7.2
Clay
Conten
t(%)
pH
4
5
7
CEC
(meq/
100 g)
Surface
Area
(mVg)
[Ca]
(ppm)
[Sr]
Background
Solution
Soil ID
Mol Lignitic
Sand
Mol Lignitic
Said
Mol Lignitic
Sand
Reference1
and Comments
12
12
12
1 References: 1 = Ohnuki, 1994; 2 = Jackson and Inch ,1989; 3 =Patterson and Spoel ,1981; 4 = Keren and O'Connor, 1983; 5 Nelson.
1 959; 6 = Inch and Kffley, 1 987; 7 = Rhodes and Nelson, 1 957; 8 = Konishi et al. , 1 988; 9 = Seme at al. , 1 993; 1 0 = Vine et al. , 1 980;
11 - 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, 27:275-279.
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. "Clmoptilolite 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."
Radiochimica Acta, 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.5 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 at Hanford 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.
H.25
-------
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.
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, andK. 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. The
1.2
-------
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.
Mean
Standard Error
Median
Mode
Standard Deviation
Sample Variance
Minimum
Maximum
No. Observations
Thorium K,,
(ml/g)
54,000
29,944
5,000
100,000
123,465
l.SxlO10
100
500,000
17
Clay
Content
(wt.%)
26.8
6.3
30
40
.14.1
199.2
12
40
5
pH
6.1
0.4
6
6
1.5
2.1
4
10
17
CEC
(meq/100 g)
13.7
11.2
2.9
2.9
29.8
886.2
1.7
81.2
7
Calcite
(wt.%)
29
13.4
25
0
30.1
905
0
60
5
Al/Fe-
Oxides
(wt%)
-
—
—
—
—
-
—
-
0
Organic
Matter
(wt.%)
-
-
-
-
-
~
-
—
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 L2. Correlation coefficients (r) of the thorium Kd value data set presented in
Section 1.3.
~
Thorium Kd
Clay Content
pH
CEC
Calcite
Thorium K,,
1
-0.79
0.58 2
(0.14)3
-0.15
0.76
Clay Content
1
-0.84 1
—
-0.998 2
pH
1
-0.21
0.85 1
CEC
1
—
1>2 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 IQ values at the highest pH value (pH 10), the con-elation is 0.58 (n = 16). Including this Kj
value, the correlation coefficient decreases to 0.14.
1.2.2 Thorium Kd Values as a Function ofpH
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
-------
6
-*. 5
SP
y =-0.13+ 0.69x;r = 0.71
3456789 10 11
pH
Figure 1.1. Linear regression between thorium K
-------
The regression equation between the pH range of 4 to 8 that is shown in Figure 1.1 is
log (Th Kd) = -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 2:0.05 indicates that the intercept is not significantly (P z 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.]
Intercept Coefficient
Slope Coefficient
Coefficients
2.22
0.57
Standard
Error
1.06
0.18
t- Statistic
0.47
3.24
P-value
0.64
0.006
Lower
95%
-1.77
0.19
Upper
95%
2.76
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 a!., 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 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 I.I 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 Kj) = -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" categoiy, the lower and upper
log (Th Kj) 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 10'26 M for pH < 5) were
assigned a Kd of 300,000 ml/g.
Table L5. 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 rng/1), no organic
chelates (such as EDTA), and oxidizing conditions.]
IMinl/g)
Minimum
Maximum
pH
3-5
Dissolved Th (M)
<10-2.6
62
6,200
>10-2.6
300,000
300,000
5-8
Dissolved Th (M)
<109
1,700
170,000
>109
300,000
300,000
8-10
Dissolved Th (M)
<109
20
2,000
>109
300,000
300,000
1.7
-------
7
6
V 5
£ 4
s
3
2
1
i i i i
s
0 ""
^yV^
xX £'X° '
x x %4^ *'*'
- ftX*^v *'
y
t O
- »' •
t
t
t
i i i i
2 46 8 10 12
PH
Figure 1.2. Linear regression between thorium K^ values
and pH for the pH Range 4 to 8. [Values ±1
logarithmic unit from the regression line are
also identified. The single K^ value at pH 10
is identified by the filled circle)].
1.3.0 K Data Set for Soils
The data set of thorium Ka 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
Kd
(ml/g)
10,0000
500,000
1,000
100,000
150,000
100
24,000
5,800
1,028.6
1,271
5,000
10,000
15,000
1,578
1,862.5
1,153.7
206.9
pH
7.6
6
4
8
7
10
6
6
5.1
52
4.5
5.8
7
5.2
5.1
5.2
62.
Clay
(wt%)
40
40
12
30
12
CEC1
(meq/
lOOg)
3
2.9
2.1
81.2
2.9
2.1
1.7
OMP
(wt.%)
Fe-
Oxides
(wt%)
Th
(M)
Calcite
(wt.%)
0
0
60
25
60
Solution
Chemistry
Synthetic GW,
pH6.6
Syn-GW,232!!!
Competing Ion
Syn-GW,232^
Competing Ion
Syn.GW,232^
Competing Ion
Syn-GW,232!!!
Competing Ion
Syn-GW,232!!!
Competing Ion
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Groundwater
Soil ID and
Characteristics
Soil A
Silt+Qtz Sed., Schist soil
Silt+Qtz Sed., Schist soil
SilHQtz+OM+calcite,
Schist Soil
Cadarache Sed.
SilW-QtzHOM+calcite,
Schist Soil
Glacial till, Clay
Fine Coarse Sand
Gleyed Dystric Brunisol, Ae
Horizon 4-1 5 cm
Gleyed Dystric Brunisol, Bf
Horizonl 5-45 cm
Jefferson City, Wyoming,
Fine Sandstone and Silty
Clay
Jefferson City, Wyoming,
Fine Sandstone and Silty
Clay
Jefferson City, Wyoming,
Fine Sandstone and Silty
Clay
Gleyed Dystric Brunisol, Ah
Horizon
Gleyed Dystric Brunisol, Ae
Horizon
Gleyed Dystric Brunisol, Bf
Horizon
Gleyed Dystric Brunisol, C
Horizon
Ref
1
2
2
2
2
2
3
3
4
' 4
5
5
5
6
6
6
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 et al, 1987; 5 = Haji-Djafari et al.,
1981; 6 = Thibaultef 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. Grouse. 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
-------
Rai, D., A. R. Felmy, D. A. Moore, and M. J. Mason. 1995. "The Solubility of Th(IV) and
U(TV) Hydrous Oxides in Concentrated NaHCO3 and Na^Os Solutions." In Scientific Basis
for Nuclear Waste Management XHII, Part 2, T. Murakami and R. 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 ofRadionuclides 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(TV) 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, K^for Use in Environmental Assessments.
AECL-10125, Whiteshell Nuclear Research Establishment, Atomic Energy of Canada
Limited, Pinawa, Canada.
1.11
-------
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(VT). 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 carbonate complexes on the adsorption and solubility of
dissolved U(VT). This complexation especially affects the adsorption behavior of U(VT) 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 IQ 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
-------
J.2.1 Sources of Error and Variability
The Kj 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 Kj) or up to more
than 95 percent (very large Ka) 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 hi
terms of "uranium," and do not distinguish between the different valence states of uranium [U(VI)
and U(TV)] present. In most studies, it is fair for the reader to assume that the authors are
referring to U(VT) 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(TV) by reaction with ferrous iron [Fe(n)] exposed on fresh mineral surfaces.
Because a major decrease of dissolved uranium typically results from this reduction due to
precipitation of U(EV) hydrous-oxide solids (i.e., lower solubility), the measured Kd values can be
too large as a measure of U(VT) sorption. This scenario is possible when one considers the
geochemical processes associated with some in situ remediation technologies currently under
development For example, Fruchter et 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(ni) present in the sediment minerals to Fe(TT). Laboratory experiments have shown
that dissolved U(VT) will accumulate, via reduction of U(V1) to U(IV) and subsequent
precipitation as a U(IV) solid, when it contacts such treated sediments.
J.3
-------
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 et al. (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 etal. (1982, Table HI) 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 Cl :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 Cl: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 da}/s and an initial concentration of
dissolved uranium of approximately 3. IxlO'8 mg/ml. The uranium Ka values measured at pH
values of 2.8 and 7.1 by Erickson (1980) are listed in Table J.5.
J.4
-------
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 |jg/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-SO4) containing 100 ug/1 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(VT) 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(VT) was determined to be constant for
concentrations between 3.3 and 100 ug/1 UOf" 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(VT)-containing solids occurred, which resulted in apparent Kd values greater than 400 ml/g.
Kaplan et al. (1996) measured the Kd values for U(VT) 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 absorption 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 etal. [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 ug/1 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(VT) 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(VT) 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
(1^0.1). The aqueous solutions consisted of a sample of uncontaminated groundwater from the
Hanford Site. The Kd values listed in Kaplan and Seme (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, 26.1
J.6
-------
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
K,j 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 deteimined for initial
concentrations of l.OxlQ-4, l.OxlO'5, LOxlO"6, l.OxlO'7, and 2.15xlO-8M uranium. TheKd
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 73>U and ^U for 48 soil/porewater
samples. The Kd values were determined from analyses of a8U and S5U in soil samples and
associated porewaters taken from contaminated zones downgradient of the seepage basins. It
should be noted that the mass concentration of ^U is significantly less than (e.g., <1 percent) the
concentration of a8U 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 ^U 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. 1). The reader should note that the field-derived Kd
values in Figures J.I, 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,
J.7
-------
and then decreases with increasing pH over the pH range from 5.2 to 6.7. Serkiz and Johnson
found that the field-derived Kd values for ^U and ^U 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.
10,000
-a . 1,000
100
10
5
pH
Figure J.I. Field-derived Kj values for **U and S5U 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 ^U and 235U,
respectively. Solid symbols represent minimum Kd values for
23SU and 23SU that were based on minimum detection limit
values for the concentrations for the respective uranium
isotopes in porewaters associated with the soil sample.]
18
-------
100,000
10,000
•SJ 1,000
"5
« 100
10
1
(
n • d'-' d
DD DD •
D "P, *
n rf** • *••*• * •
§b lefe n n
2 «•© D Q *
j
o §o
* Ps
^
3 10 20 30 40 5
Clay-Size Particle Content (wt%)
0
Figure J.2. Field-derived Kd values for ^U and ^U 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 ^"U and ^U, respectively. Solid symbols
represent minimum Kd values for ^U and ^U 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
-------
1 00 000
10,000
1
* 100
10
1
1
c
n " 'n * n B
n D •
n • H • D
«t,Q« n B ^ Q
D
: R D
-. Q o o
L|^eg a
) 5 10 15 20 2
CEC (meq/kg)
5
Figure J.3. Field-derived Kd values for ^U and asU 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
^U and 235U, respectively. Solid symbols represent minimum Kd values for
^U 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.]
Seme 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 Seme et
al. (1993) are given in Table 15.
110
-------
Sheppard and Thibault (1988) investigated the migration of several radionudides, 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 Kj 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 etal (1976), Haji-Djafari etal (1981), Neiheisel (1983),
Rancon (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.
Waraeckeandcoworkers(Warneckee*a/., 1984, 1986, 1988, 1994; Warnecke and Hild, 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 n 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.ll
-------
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 etal. (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 n 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 13) for sediments and different water types
for the Konrad site are: 4 (Quaternary fresh water), 6 (Turanian 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 HI) list minimum and maximum uranium Kd values (0.54-15.2
ml/g) for 26 rock samples from the Asse n 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.I. Uranium Kd values (ml/g) listed by Warnecke et al. (1994, Table 1).
Sediment
Type
Sandy
Clayish
Fresh Water
Typical
Kd Value
27
17
Minimum
K,, Value
0.8
8.6
Maximum
Kd Value
332
100
Saline Water
Typical
K* Value
1
14 - 1,400
Minimum
Kj Value
0.3
14.1
Maximum
Kd Value
1.6
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
-------
Wamecke 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 Konracl site sediments and
waters duplicate those presented Warnecke et al. (1994) and Warnecke and Hild (1988).
Wamecke 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 Wamecke et
al (1984, Table IH) 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(VT) on clay-mineral separates from subsurface
soils from 3 DOE sites. The materials included the clay separates (<2 pm 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 nmol/1)
uranium and Ca(ClO4)2 or NaClO4 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 ofU(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
-------
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 lO'MO'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 etal (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 NaHCO3) using 4 initial
uranium concentrations. The initial uranium concentrations used for the 25°C experiments
included l.OxlO"4, l.OxlO"5, 1.4xlO"6, and 4.4xlO'7 mol/1 uranium. The batch experiments were
conducted under oxidizing conditions at 5, 25, and 65°C hi an environmental chamber. Solutions
were equilibrated with the mineral solids in a ratio of 10 inl/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 etal. (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 NaHCO3) using 4 initial uranium concentrations. The initial uranium concentrations used
for the 25°C experiments included l.OlxlQ-4, l.OSxlO"5, LOSxlO"6, and 4.89xlQ-7 mol/1 uranium
for the 0.01 MNaCl solution, and l.OlxlQ-4, l.OSxlO'5, 1.53x10-*, and 5.46xlO"7 mol/1 uranium
for the 0.01 M NaHCO3 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 hi 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 IT) 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
MNaHCO3 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
-------
Borovec (1981) investigated the adsorption of U(VT) 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"* 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 (a™, in meq/100 g) on clays at 20°C and
pH 6.0 with respect to CEC (in meq/100 g),
a» = 0.90 CEC + 1.56 (r = 0.99522) ,
and specific surface (A, in m2/g) of clays,
8,1 = 0.11 A + 2.05 (r = 0.97232).
J.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 et 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
K,, 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 (101-6) ml/g, and
a range for its Kd values of 0.1 to 1,000,000 ml/g. Looney et al note that Iheir 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 Eire 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
-------
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 hi
Table J.2. The reader should refer to sources cited hi 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
NirexLtd.); 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); RTVM [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
-------
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.
Orpsnwstxon
AECL
GSF
IAEA
KBS-3
NAGRA
NIREX
NRC
NRPB
PAGIS
PAGIS SAFIR
PSE
RIVM
SKI
TVO
UKDoE
Argillaceous (Clay)
Sorb ing
Material
Bentonite-Sand
Sediment
Pelagic Clay
Bentonite
Bentonite
Clay
Clay Mudstone
Clay, Soil Shale
Clay
Bentonite
Subseabed
Clay
Sediment
Sandy Clay
Bentonite
Bentonite
Baltic Sea
Sediment
Ocean Sediment
Lake Sediment
Clay
Coastal Marine
Water
K,,
(ml/g)
100
2
500
120
1,000
5,000
10
20
300
90
100
600
0.02
10
200
90
500
500
500
200
1000
Crystalline Rock
Sorbing
Material
Granite
Granite
Granite
Granite
Basalt
Tuff
Granite
Crystalline
Rock, Reducing
Crystalline
Rock, Real.
Ka
(ml/g)
5
5,000
1,000
5
4
300
5,000
200
5
Soil/Soil ||
Sorbing
Material
Soil/Sediment
Soil/Sediment
Soil/Sediment
Soil/Sediment
Soil/Sediment
Soil/Sediment
Soil/Sediment
K-
(ml/R)
20
20
100
300
1,700
500
50
J.17
-------
In 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 in
crystalline rock/water systems of 1,000 (NAGRA), 5,000 [Svensk Karnbranslehantering AB
(Nuclear Fuel and Waste Management Company), Sweden; 8KB], 1000 (TVO), and 6 (Canadian
Nuclear Fuel Waste Management Programme, CNFWM). For bentonite/groundwater systems,
they list 5,000 (NAGRA), 3,000 (8KB), 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 assessment(s) 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
etal. (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 3:35 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.
SoU Type
Sand
Loam
Clay
Organic
Geometric
MeanK,,
Values (ml/g)
35
15
1,600
410
Observed Range of
Kd Values (ml/g)
0.03 - 2,200
0.2 - 4,500
46 - 395,100
33 - 7,350
Number of
K,, Values
24
8
7
6
J.18
-------
J.3.0 Approach in Developing Kj 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 hi 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 ffpH
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 K<, values
reported by Serkiz and Johnson (1994) (see Figure J.I), 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 andLangmuir, 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(V1), especially near and above neutral
pH conditions where dissolved U(VT) 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 et al. (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 at
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
-------
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 Kj 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 etal. (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(V1) through the
formation of strong anionic carbonate complexes. In turn, this complexation affects the
adsorption behavior of U(VT) at alkaline pH conditions. Even differences in partial pressures of
CO2 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 CO2 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 CO2 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
J.21
-------
discussed, and one would have to make assumptions about possible equilibrium between the
solutions and atmospheric or soil-related partial pressures of CO2 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, Serlciz 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 etal (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 KdLook-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)
Minimum
Maximum
pH
3
<1
32
4
0.4
5,000
5
25
160,000
6
100
1,000,000
7
63
630,000
8
0.4
250,000
9
<1
7,900
10
<1
5
J.22
-------
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 et al.
(1992) and the Kd values reported by Kaplan et al. (1996, 1998), Lindenmeirer et al. (1995), and
Seme 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(YI)/U(TV) redox/precipitation reactions resulting from contact of dissolved U(VT) with Fe(n) •
exposed on the fresh mineral surfaces.
J.4.0 Use of Surface Complexation Models to Predict Uranium K,, 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 Kj 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 etal, 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, K^) was derived for uranium sorption using the
MENEQL geochemical code and site-specific geochemical data for soil CEC values, groundwater
compositions, and measured uranium Kd values. Quartz (SiO^ 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
[cc-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(VT) 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 behavior
associated with disposal of high-level radioactive waste in a geologic repository. The parameter
J.24
-------
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 et
al (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., Kj) 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 appendk are provided exactly as
given by Ames and Rai (1978).
J.25
-------
• Manskaya et al. (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
-------
powder diffraction (XRD) indicated the formation of uranium-containing phases accompanied by
unreacted zeolite. The products of the reactions involving Na- and K-A zeolites contained a
phase similar to compreignacite (K2O6UO3-11H2O). Those experiments conducted with Ca-A
zeolite contained a phase similar to becquerelite (CaO6UO3-11H2O).
Ho and coworkers studied the adsorption of U(VT) on a well-characterized, synthetic hematite
(a-Fe^) 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 pm, surface area of 34
mVg, isoelectric point2 of pH 7.6, and composition of >98 percent a-Fe;,O3 and <2 percent
p-FeO(OH). Ho and Doem (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(VT) 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(VT) was measured.
The adsorption of U(VT) 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(VT) concentration over the pH range from 6.5 to 9.1 using the
hematite sol described previously. Their experimental procedure was simikr to that described by
Ho and Doern, except that the measurements were completed using a 1x10"3 mol/1 NaHCO3
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 SxlO"6 mol/1, the
reported percentages of U(VT) 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 arid 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(VT) 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(VT) 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
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
-------
concentrations of humic acid, there is a change 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(VT) by Wyoming montmorillonite. The
experiments consisted of reacting, at room temperature, the <2-um size fraction of
montmorillonite saturated with Na+, K+, Mg2+, Ca2+, and Ba2+ with U(VT) nitrate solutions
containing 1 to 300 ppm U(VT). 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(VT)] and systems with variable volumes and fixed amounts of U(VT) [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 (UO2+) was strongly preferred
over Na+ and K+ by the clay, and less strongly preferred versus Mg2+, Ca2+, and Ba2+.
Vochten etal (1990) investigated the adsorption of U(VT) hydrolytic complexes on well-
characterized samples of natural zeolites in relation to the double-layer potential of the minerals.
The zeolite samples included chabazite (CaAl2Si4O12-6H2O), heulandite
[(Ca,Na2)Al2Si1On-6H.2O], scolecite (CaAl2Si3O10-3H2OX and stilbite
[(Ca,Na2,K2)Al2Si7Olg-7H2O]. 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 jjm) zeolite added to a 50 ml solution of
2xlO'5 mol/1 U(VT). The suspension was shaken for 1 week in a nitrogen atmosphere to avoid the
formation of U(VT) carbonate complexes. Given the relatively small dimension of the channels in
the zeolite crystal structure and ionic diameter of the non-hydrated UOf" 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(VT) 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(VT) 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 mat Vochten et al. (1990) may have
exceeded the solubility of U(VI) above pH 5 and precipitated a U(VI) solid, such schoepite
(UO3-2H2O), 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
8.3
8.3
8.3
83
8.3
83
83
83
83
83
83
83
83
83
83
8.3
83
83
83
83
UKd
(ml/g)
1.98
0.49
2.81
0.62
0.45
0.54
0.62
0.40
0.10
0.08
2.0
0.5
2.7
1.0
0.5
0.2
1.1
1.1
0.6
0.6
Clay
Cont
(wt%)
CEC
(meq/lOOg)
5.2
5.2
5.2
5.2
5.2
5.2
5.2
5.2
5.2
5.2
Surface
Area
(mVg)
Solution
lanford Groundwater
ianford Groundwater
lanford Groundwater
rlanford Groundwater
rlanford Groundwater
Hanford Groundwater
Stanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Soil Identification
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Trench 8 Loamy Sand
Reference / Comments
Caplan and Serne (1995,
Part Sat Column, 40%)
Kaplan and Seme (1995,
Part. Sat Column, 40%)
Kaplan and Serne (1995,
Part Sat Column, 38%)
Kaplan and Seme (1995,
Part Sat Column, 22%)
Kaplan and Serne (1995,
Part Sat. Column, 30%)
Kaplan and Seme (1995,
Part Sat. Column, 23%)
Kaplan and Serne (1995,
Part Sat. Column, 25%)
Kaplan and Serne (1995,
Part Sat Column, 17%)
Kaplan and Seme (1995,
Part Sat Column, 7%)
Kaplan and Serne (1995,
Part Sat Column, 7%)
Lindenmeir et al. (1995,
Saturated Column 1)
Lindenmeir et al. (1995,
Saturated Column 1)
Lindenmeir eraZ. (1995,
Saturated Column 1)
Lindenmeir et al. (1995,
Unsat Column 1, 65%)
Lindenmeir et al. (1995,
Unsat UFA 1,70%)
Lindenmeir etal. (1995,
Unsat UFA 2, 24%)
Lindenmeir et al. (1995,
U nsat Column 1, 63%)
Lindenmeir etal. (1995,
Unsat Column 2, 43%)
Lindenmeir etal. (1995,
Unsat UFA 1A, 29%)
Lindenmeir et al. (1995,
Unsat. UFA 1C, 29%)
J.29
-------
pH
8.4
8.4
8.4
8.4
8.4
7.92
8.05
7.99
7.99
7.98
7.97
8.48
8.26
8.44
9.12
8.46
8.46
8.46
8.46
8.46
8.46
8.46
8.46
8.46
8.46
8.46
8.46
8.46
8.46
8.46
8.46
8.46
UKd
(ml/g)
0.20
0.15
0.09
0.15
0.14
1.99
1.92
1.91
2.10
2.25
2.44
1.07
1.46
1.37
2.12
0.90
1.70
1.00
1.10
3.50
2.10
0.24
0.64
0.51
0.46
0.35
0.53
0.23
0.15
0.1
0.16
0.12
Clay
Cent
(wt%)
CEC
(meq/lOOg)
5.3
5.3
5.3
5.3
5.3
6.4
6.4
6.4
6.4
6.4
6.4
6.4
6.4
6.4
6.4
6.4
5.3
6.0
6.4
5.3
6.0
6.4
6.4
6.4
6.4
6.4
6.4
5.3
5.3
5.3
5.3
5.3
Surface
Area
(mVg)
6.3
6.3
6.3
6.3
6.3
14.8
14.8
14.8
14.8
14.8
14.8
14.8
14.8
14.8
14.8
14.8
6.3
6.3
14.8
6.3
6.3
14.8
14.8
14.8
14.8
14.8
14.8
6.3
6.3
6.3
6.3
6.3
Solution *
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
lanford Groundwater
ianford Groundwater
ianford Groundwater
SoU Identification
Trench 94
Trench 94
Trench 94
Trench 94
Trench 94
Trench AE-3
Trench AE-3
Trench AE-3
Trench AE-3
Trench AE-3
Trench AE-3
Trench AE-3
Trench AE-3
Trench AE-3
Trench AE-3
Trench AE-3
Trench 94
TSB-1
Trench AE-3
Trench 94
TSB-1
Trench AE-3
Trench AE-3
Trench AE-3
Trench AE-3
Trench AE-3 •
Trench AE-3
Trench 94
Trench 94
Trench 94
Trench 94
Trench 94
Reference / Comments
Kaplan etal. (1998, Batch)
Kaplan etal. (1998, Batch)
Kaplan etal. (1998, Batch)
Kaplan et al. (1998, Batch)
Kaplan etal. (1998, Batch)
Kaplan etal. (1998, Batch)
Kaplan et al. (1998, Batch)
Kaplan et al. (1998, Batch)
Kaplan et al. (1998, Batch)
Kaplan etal. (1998, Batch)
Kaplan et al. (1998, Batch)
Kaplan etal. (1998, Batch)
Kaplan etal. (1998, Batch)
Kaplan etal. (1998, Batch)
Kaplan et al. (1998, Batch)
Kaplan et al. (1996, 100%
Unsaturated Batch)
Kaplan etal. (1996, 100%
Unsaturated Batch)
Kaplan etal. (1996, 100%
Unsaturated Batch)
Kaplan etal. (1996, Batch)
Kaplan etal. (1996, Batch)
Kaplan et al. (1996, Batch)
Kaplan etal. (1996)
Kaplan et al. (1996)
Kaplan etal. (1996)
Kaplan et al. (1996)
Kaplan etal. (1996)
Kaplan et al. (1996)
Kaplan etal. (1996)
Kaplan et al. (1996)
Kaplan et al. (1996)
Kaplan etal. (1996)
Kaplan et al. (1996)
J.30
-------
pH
2.00
4.50
5.75
7.00
5.6
5.6
5310
5.10
5.20
6.20
7.00
UKd
(ml/g)
2
1
3
750
770
550
100
200
1,000
2,000
25,000
250
58.4
294.9
160
45.4
450
day
Cent
(wt.%)
8
7
15
36
21
19
36
CEC
(meq/lOOg)
28.0
Surface
Area
(mVg)
Solution
SoU Identification
Sand
Sand
Sand
Clayey Sand
Clayey Sand
Clayey Sand
?ine Sandstone and
Silty Sand
Fine Sandstone and
Silty Sand
Fine Sandstone and
Silty Sand
Fine Sandstone and
Silty Sand
Red-Brown Clayey
Red-Brown Clayey
Silty Loam Clay
Reference / Comments
Neiheisel [1983, as listed
in Thibault et al. (1990)]
Neiheisel [1983, as listed
in Thibault ef a/. (1990)]
Neiheisel [1983, as listed
in Thibault etal. (1990)]
Neiheisel [1983, as listed
in Thibault et al, (1990)]
Neiheisel [1983, as listed
in Thibault etal. (1990)]
Neiheisel [1983, as listed
inlhibaultefa/. (1990)]
Haji-Djafari etal. [1981, as
listed in Thibault et al.
(1990)]
Haji-Djafari et al. [1981, as
listed in Thibault et al.
(1990)]
Haji-Djafari et al. [1981, as
listed in Thibault et al.
(1990)]
Haji-Djafari etal. [1981, as
listed in Thibault et al.
(1990)]
Seeley and Kelmers [1984, as
listed in Thibault et al.
(1990)]
Seeley and Kelmers [1984, as
listed in Thibault et al.
(1990)]
Thibault et al. (1990, values
determined by coworkers)
Thibault et al. (1990, values
determined by coworkers)
Thibault et al. (1990, values
determined by coworkers)
Thibault et al. (1990, values
determined by coworkers)
Thibault et al. (1990, values
determined by coworkers)
J.31
-------
pH
7.30
4.90
5.50
7.40
7.40
6.60
6.50
7.10
7.00
7.00
6.50
2.80
7.10
8.3
8.3
8.3
8.00
8.00
8.00
8.00
8.00
8.00
UKd
(ml/g)
1.2
0.03
2900
1.9
2.4
590
4500
15
16
33
4400
200
790,000
1.70
2.30
79.30
56.0
7.5
13.2
17.8
20.2
13.0
Clay
Cent
(wt%)
15
2
1
10
11
10
10
12
CEC
(meq/lOOg)
17.0
5.8
120.0
9.1
8.7
10.8
12.6
13.4
2.6
5.2
6.0
Surface
Area
(m2/g)
Solution !"
~
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater
Hanford Groundwater, GR-1
Hanford Groundwater, GR-1
Hanford Groundwater, GR-1
Hanford Groundwater, GR-1
Hanford Groundwater, GR-1
Hanford Groundwater, GR-1
Soil Identification
Loam
Medium Sand
Organic
Fine Sandy Loam
Fine Sandy Loam
Fine Sandy Loam
Fine Sandy Loam
Fine Sandy Loam
Sand
Organic Peat
Clay Fraction
Abyssal Red Clay
Abyssal Red Clay
CGS-1 sand (coarse
gravel sand)
Trench 8 Loamy Sand
(medium/coarse sand)
TBS-1 Loamy Sand
(Touchet Bed sand)
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Flow E Basalt
Reference / Comments
Thibault et al. (1990, values
determined by coworkers)
Thibault et al. (1990, values
determined by coworkers)
Thibault et al. (1990, values
determined by coworkers)
Thibault et al. (1990, values
determined by coworkers)
Thibault et al. (1990, values
determined by coworkers)
Thibault et al. (1990, values
determined by coworkers)
Thibault et al. (1990, values
determined by coworkers)
Thibault etal (1990, values
determined by coworkers)
Rancon [1973, as listed in
Thibault et al. (1990)]
Rancon [1973, as listed in
Thibault etal. (1990)]
Dahlman etal. [1976, as listed
in Thibault et al. (1990)]
Erickson (1980)
Erickson(1980)
Serne etal. (1993, Batch)
Serne etal. (1993, Batch)
Seme etal. (1993, Batch)
Salter etal. (1981)
Salter etal. (1981)
Salter ef a/. (1981)
Salter etal. (1981)
Salter etal. (1981)
Salter etal. (1981)
J.32
-------
pH
8.00
8.00
8.00
8.00
8.00
8.00
8.00
8.00
8.00
10.00
10.00
10.00
10.00
10.00
10.00
10.00
10.00
10.00
10.00
10.00
10.00
10.00
10.00
10.00
7.66
7.66
7.66
7.66
8.38
8.38
8.38
838
7.65
7.65
UKd
(ml/g)
2.7
2.2
3.2
2.9
16.0
2.2
3.5
5.2
5.8
2.8
23
2.8
2.8
2.5
1.0
0.5
0.4
0.8
0.2
0.9
0.6
0.8
0.5
0.4
7.5
13
18-
20
2.4
2.9
2.9
2.5
2.7
2.2
Clay
Cont
(wt%)
CEC
(meq/lOOg)
1.83
1.83
1.83
1.83
1.83
1.83
1.83
1.83
1.5
1.5
Surface
Area
(mVg)
17.7
17.7
17.7
17.7
17.7
17.7
17.7
17.7
10.3
10.3
Solution
Hanford Groundwater, GR-1
Hanford Groundwater, GR-1
Hanford Groundwaler, GR-1
Hanford Groundwater, GR-1
Hanford Groundwater,GR-l
Hanford Groundwaler.GR-1
Hanford Groundwater,GR-l
Hanford Groundwater,GR-l
Hanford Groundwater,GR-l
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwaler,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwaler,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-l
Hanford Groundwater,GR-l
Hanford Groundwater,GR-l
Hanford Groundwater,GR-l
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-l
Hanford Groundwater,GR-l
Soil Identification
Flow E Basalt
Flow E Basalt
Flow E Basalt
Flow E Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Flow E Basalt
Flow E Basalt
Flow E Basalt
Flow E Basalt
Flow E Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Umtanum Basalt
Flow E Basalt
Flow E Basalt
Reference / Comments
Salter etal. (1981)
Sailer era/. (1981)
Salter etal. (1981)
Salter et al. (1981)
Salter etal. (1981)
Salter etal. (1981)
Salter et al. (1981)
Salter et al. (1981)
Salter etal. (1981)
Salter etal. (1981)
Salter etal. (1981)
Salter etal. (1981)
Salter et al. (1981)
Salter etal. (1981)
Salter etal. (1981)
Salter etal. (1981)
Salter etal. (1981)
Salter etal. (1981)
Salter etal. (1981)
Salter et al. (1981)
Salter et al. (1981)
Salter etal. (1981)
Salter etal. (1981)
Salter etal (1981)
Ames etal (1982)
Ames etal. (1982)
Ames etal. (1982)
Ames etal. (1982)
Ames etal. (1982)
Ames et al. (1982)
Ames etal. (1982)
Ames et al. (1982)
Ames etal. (1982)
Ames et al. (1982)
J.33
-------
pH
7.65
7.65
8.38
8.38
838
8.38
7.90
7.90
7.90
7.90
8.48
8.48
8.48
8.48
7.7
7.7
7.7
7.7
7.7
7.7
7.7
7.7
6.85
6.80
6.90
6.90
8.60
8.65
8.65
8.80
UKd
(ml/g)
3.2
2.9
0.55
0.38
0.78
0.19
2.2
3.5
5.2
5.8
0.57
0.83
0.47
0.42
27
. 39
127
76
12
42
48
22
477,285
818,221
1,739,87
7
1,690,52
2
4313
14,098
21,362
26,269
Clay
Cont
(wt%)
CEC
(meq/lOOg)
1.5
1.5
1.5
1.5
1.5
1.5
4.84
4.84
4.84
4.84
4.84
4.84
4.84
4.84
71.66
4.84
4.84
4.84
4.84
4.84
4.84
4.84
Surface
Area
(mVg)
10.3
10.3
10.3
10.3
10.3
10.3
31.2
31.2
31.2
31.2
31.2
31.2
31.2
31.2
646
31.2
31.2
31.2
31.2
31.2
31.2
31.2
Solution f
Hanford Groundwaler,GR-l
Hanford Groundwater,GR-l
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-l
Hanford Groundwater,GR-l
Hanford Groundwater,GR-l
Hanford Groundwater,GR-l
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-l
Hanford Groundwater,GR-l
Hanford Groundwater,GR-l
Hanford Groundwater,GR-l
Hanford Groundwater,GR-2
Hanford Groundwaler,GR-2
Hanford Groundwater,GR-2
Hanford Groundwater,GR-2
0.01 NaCl
0.01 NaCl
0.01 NaCl
0.01 NaCl
0.01 NaHCCv
0.01 NaHCO3
O.OlNaHCOj
0.01NaHCO3
SoU Identification
Flow E Basalt
Flow E Basalt
Flow E Basalt
Flow E Basalt
Flow E Basalt
Flow E Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Pomona Basalt
Smectite, secondary
Smectite, secondary
Smectite, secondary
Smectite, secondary
Smectite, secondary
Smectite, secondary
Smectite, secondary
Smectite, secondary
AmorFe(ni)
Hydroxide
AmorFe(nr)
Hydroxide
AmorFeCffl)
Hydroxide
AmorFeCIH)
Hydroxide
AmorFeCffl)
Hydroxide
AmorFe(ffl)
Hydroxide
AmorfefOI)
Hydroxide
AmorFe(IH)
Hydroxide
Reference / Comments
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames era/. (1982)
Ames er a/. (1982)
Ames era/. (1982)
Ames era/. (1983c)
Ames er a/. (1983c)
Ames era/. (1983c)
Amesera/.(1983c)
Ames era/. (1983c)
Ames era/. (1983c)
Ames era/. (1983c)
Ames era/. (1983c)
J.34
-------
pH
7.15
7.15
7.15
7.15
7.15
7.15
7.15
7.15
7.15
7.15
7.15
8.65
8.65
8.65
8.65
8.65
8.65
8.65
8.65
8.65
8.65
8.65
7
8.5
7
8.5
7
8.5
7
UKd
(ml/g)
8.4
43.9
253.5
544.3
113.7
251.0
459.7
68.2
67.9
85.4
95.4
0.9
3.4
23.0
80.8
2.2
26.9
602.5
3489.6
0.6
1.1
0.6
544.5
90.5
657.8
400.8
542.0
1.8
299.9
Clay
Cont
(wt%)
CEC
(meq/lOOg)
15.3
15.3
15.3
15.3
0.95
0.95
0.95
0.95
1.17
1.17
1.17
15.3
15.3
15.3
15.3
0.95
0.95
0.95
0.95
1.17
1.17
1.17
25
25
12.2
12.2
120
120
95
Surface
Area
(mVg)
1.59
1.59
1.59
1.59
1.88
1.88
1.88
1.88
1.22
1.22
1.22
1.59
1.59
1.59
1.59
1.88
1.88
1.88
1.88
1.22
1.22
1.22
116.1
116.1
68.3
68.3
747
747
861
Solution
0.01 NaCl
0.01 NaCl
0.01 NaCl
0.01 NaCl
0.01 NaCl
0.01 NaCl
0.01 NaCl
0.01 NaCl
0.01 NaCl
0.01 NaCl
0.01 NaCl
0.01NaHCO3
0.01 NaHC03
0.01 NaHCO3
0.01NaHCO3
0.01 NaHC03
0.01 NaHCO3
0.01 NaHC03
0.01 NaHCO3
0.01 NaHCO3
0.01 NaHC03
0.01 NaHC03
0.01 NaCl
0.01 NaHCOj
0.01 NaCl
0.01 NaHCO3
0.01 NaCl
0.01 NaHCO3
0.01 NaCl
Soil Identification
Biotite
Biotite
Biotite
Biotite
Muscovite
Muscovite
Muscovite
Muscovite
Fhlogopite
Phlogopite
Phlogopite
Biotite
Biotite
Biotite
Biotite
Muscovite
Muscovite
Muscovite
Muscovite
Phlogopite
Phlogopite
Phlogopite
niHe, only lowest U
cone
Illite, only lowest U
cone
Kaolinite, only lowest
Uconc
Kaolinite, only lowest
Uconc
Montmorillonite, only
lowest U cone
Montmorillonite, only
lowest Uconc
Nontronite, only lowest
Uconc
Reference / Comments
Amesetal. (1983b)
Amesetal. (1983b)
Amesetal. (1983b)
Amesetal. (1983b)
Amesetal. (1983b)
Amesetal. (1983b)
Amesetal. (1983b)
Amesetal. (1983b)
Amesetal. (1983b)
Amesetal. (1983b)
Aiassetal. (1983b)
Amesefa/.(1983b)
Amesetal. (1983b)
Amesetal. (1983b)
Ames efoZ. (1983b)
Amesefo/.(1983b)
AmesefaZ. (1983b)
Amesetal. (1983b)
Amesetal. (1983b)
Amesetal. (1983b)
Amesetal. (1983b)
Amesetal. (1983b)
Amesetal. (1983a)
Amesetal. (1983a)
Amesefat(1983a)
Amesetal. (1983a)
Ameserai(1983a)
Ames et al. (1983a)
Amesetal. (1983a)
J.35
-------
pH
8.5
7
8.5
7
8.5
7
8.5
7
8.5
7.3
6.2
8.0
6.8
5.6
8.0
UKd
(ml/g)
4.1
138.0
114.2
66.5
0.6
225.7
1.7
300.5
639.9
4200.0
136.0
44
4360
328
54
39
33
16
270
0
7
139
Clay
Cont
(wt%)
CEC
(meq/lOOg)
95
16.03
16.03
140.2
140.2
3.18
3.18
2.79
2.79
4.36
1.29
9.30
4.36
1.29
9.30
Surface
Area
(mz/g)
861
137.3
137.3
20
20
46.8
46.8
626.3
626.3
Solution '
0.01 NaHCO3
0.01 NaCl
0.01 NaHCO3
0.01 NaCl
0.01 NaHCO3
0.01 NaCl
0.01 NaHCO3
0.01 NaCl
0.01 NaHCO3
Soil Identification
Nontronite, only lowest
Uconc
Glauconite, only lowest
Uconc
Glauconite, only lowest
Uconc
Clinoptilolite, only
lowest U cone
Clinoptilolite, only
lowest Uconc
Opal, only lowest U
cone
Opal, only lowest U
cone
Silica Gel,, only lowest
Uconc
Silica Gel,, only lowest
Uconc
Spesutie (silt loam)
Transonic (silt loam)
Yuma (sandy loam)
Spesutie (silt loam)
Transonic (silt loam)
Yuma (sandy loam)
River Sediment
(Quartz, clay, calcite,
organic matter)
River Peat
River Sediment
(Quartz, clay, calcite)
Soil (Quartz and Clay,
from Altered Schist)
Quartz
Calcite
niite
Reference / Comments
Am&setal. (1983a)
Amesetal. (1983a)
Amesetal. (1983a)
Ameserai (1983a)
Amesetal. (1983a)
Amesetal. (1983a)
Amesefa/.(1983a)
Ames et al. (1983a)
Amesetal. (1983a)
Erikson etal.(1993)
Erikson era/. (1993)
Erikson efai(1993)
Erikson efrf. (1993)
Erikson et al. (199.3)
Erikson eraZ. (1993)
Rancon (1973) as cited
by Ames and Raj (1978)
Rancon (1973) as cited
by Ames and Rai (1978)
Rancon (1973) as cited
by Ames and Rai (1978)
Rancon (1973) as cited
by Ames and Rai (1978)
Rancon (1973) as cited
by Ames and Rai (1978)
Rancon (1973) as cited
by Ames and Rai (1978)
Rancon (1973) as cited
by Ames and Rai (1978)
J.36
-------
pH
3.83
350
3.94
3.96
4.03
4.13
4.28
433
UKd
(ml/s)
27
(0.8-
332)
1
(0.3-1.6)
17
(8.5-
100)
14-1,400
4
6
6
20
1.4
2.6
3
3
310
235
741
211
694
720
898
630
Clay
Cent
(wt%)
CEC
(meq/lOOg)
Surface
Area
(mVg)
Solution
FreshWater
Fresh Water
Saline Water
Saline Water
Quaternary fresh water
Turonian fresh water
Cenomanian saline water
Albian (Hauterivain) saline
water
Albian (Hils) saline water
Kimmeridgian saline water
Oxfordian saline water
Bajocian (Dogger) saline
water
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Soil Identification
Goiieben Salt Dome,
Sandy Sediment
Goiieben Salt Dome,
Sandy Sediment
Gorleben Salt Dome,
Clayish Sediment
Gorleben Salt Dome,
Clayish Sediment
Former Konrad Iron
Ore Mine
Former Konrad Iron
Ore Mine
Former Konrad Iron
Ore Mine
Former Konrad Iron
Ore Mine
Former Konrad Iron
Ore Mine
Former Konrad Iron
Ore Mine
Former Konrad Iron
Ore Mine
Former Konrad Iron
Ore Mine
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Reference / Comments
Wameckeef a/. (1984, 1986,
1994), Wamecke and Hild
(1988)
Wamecke etal (1984, 1986,
1994), Wamecke and Hild
(1988)
Wameckeef al. (1984, 1986,
1994), Wamecke and Hild
(1988)
Wamecke etal. (1984, 1986,
1994), Wamecke and Hild
(1988)
Wamecke etal. (1986),
Wamecke and Hild (1988)
Wamecke etal. (1986),
Wamecke and Hild (1988)
Wamecke etal. (1986),
Wamecke and Hild (1988)
Wamecke et al. (1986),
Wamecke and Hild (1988)
Wamecke et al. (1986),
Wamecke and Hild (1988)
Wamecke etal. (1986),
Wamecke and Hild (1988)
Wameckeef al. (1986),
Wamecke and Hild (1988)
Wamecke etal. (1986),
Wamecke and Hild (1988)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
J.37
-------
pH
4.36
4.53
4.58
4.61
4.71
4.81
4.95
4.84
5.00
5.10
5.11
5.19
5.52
5.15
5.24
5.16
5.28
5.52
5.44
5.54
5.58
5.85
5.45
UKd
(ml/g)
247
264
903
324
522
1,216
1,185
3,381
2,561
2,635
3,807
4,293
4,483
4,574
5,745
7,423
3,214
5,564
6,687
6,185
6,615
7,124
8,146
Clay
Cont
(wt%)
CEC
(meq/lOOg)
—
Surface
Area
(m2/g)
Solution '
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
fiinction of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Soil Identification
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Reference / Comments
Giblin (1980)
Giblin(1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
J.38
-------
pH
5.56
5.74
5.50
5.69
5.54
5.66
5.81
5.86
5.75
6.01
6.20
5.95
6.35
6.40
6.35
6.46
6.13
6.26
6.80
6.86
6.81
7.10
7.85
UKd
(ml/g)
8,506
9,332
10,462
10,681
11,770
13,616
14,675
14,417
20,628
24,082
22,471
26,354
26,078
25,601
27,671
30,529
31,477
33,305
37,129
37,657
32312
29,390
33,583
Clay
Cent
(wt%)
CEC
(meq/lOOg)
Surface
Area
(mz/g)
Solution
Synthetic Groundwaler,
unction of pH
Synthetic Groundwaler,
junction of pH
Synthetic Groundwaler,
junction of pH
Synthetic Groundwaler,
junction of pH
Synthetic Groundwater,
[unction of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwaler,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Soil Identification
rCaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Reference / Comments
Giblin(1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
J.39
-------
pH
7.67
8.40
8.51
9.45
9.80
9.90
3.8
3.5
3.7
3.7
4.0
6.4
6.5
6.6
7.7
8.0
8.3
8.6
9.0
3.4
4.4
UKd
(ml/g)
26,518
30,523
19,632
23,177
17,763
14,499
2
5
8
69
116
1,216
1,824
2,679
7,379
2,506
21,979
3,999
14,689
27
326
Clay
Cont,
(wt%)
CEC
(meq/lOOg)
Surface
Area
(m2/g)
Solution
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Soil Identification
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Kaolinite
Quartz
Quartz
Quartz
Quartz
Quartz
Quartz
Quartz
Quartz
Quartz
Quartz
Quartz
Quartz
Quartz
Biotite
Biotite
Reference / Comments
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Giblin (1980)
Anderssonef al. (1982)
Anderssonef al. (1982)
Andersson et al. (1982)
Anderssonef orZ. (1982)
Andersson et al. (1982)
Anderssonef al. (1982)
Andersson etal. (1982)
Anderssonef al. (1982)
Andersson et al (1982) '
Andersson etal. (1982)
Andersson etal. (1982)
Anderssonef al. (1982)
Anderssonef al. (1982)
Anderssonef al (1982)
Andersson et al (1982)
J.40
-------
pH
4.4
4.7
5.1
5.2
6.4
7.3
73
7.4
8.1
9.0
3.3
3.8
4.0
4.0
4.4
4.5
5.0
5.3
6.0
7.7
3.6
UKd
(ml/g)
522
418
1,489
2,512
2,812
7,228
16,634
9,840
4,732
8,337
207
324
726
668
3,767
4,732
16,218
8,241
140,605
24,660
460
day
Cent
(wt%)
CEC
(meq/lOOg)
Surface
Area
(mVg)
Solution
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Sofl Identification
Biotite
Biotite
Biotite
Biotite
Biotite
Biotite
Biotite
Biotite
Biotite
Biotite
Apatite
Apatite
Apatite
Apatite
Apatite
Apatite
Apatite
Apatite
Apatite
Apatite
Attapulgite
(Palygorskite)
Reference / Comments
Andersson et al. (1982)
Andersson et al. (1982)
Andersson et al. (1982)
Andersson et al. (1982)
Andersson et al. (1982)
Andersson et al. (1982)
Andersson et al. (1982)
Andersson et al (1982)
Andersson et al. (1982)
Andersson etal. (1982)
Andersson et al. (1982)
Andersson et al. (1982)
Andersson etal. (1982)
Andersson et al. (1982)
Andersson etal. (1982)
Andersson etal. (1982)
Andersson et al. (1982)
Andersson etal. (1982)
Andersson etal. (1982)
Andersson et al. (1982)
Andersson etal. (1982)
J.41
-------
PH
4.1
4.2
4.5
4.7
5.1
5.9
6.4
7.3
7.8
8.7
3.2
4.4
6.6
7.0
7.0
73
8.2
8.4
9.0
5.1
5.0
UKd
(ml/g)
1,514
7,194
6,471
4,753
23,335
12,531
266,686
645,654
82,224
46,132
1,175
12,503
3,917
10,139
28,054
10,715
21,528
20,370
18,621
7,391
1,177
Clay
Cont
(wt%)
CEC
(meq/lOOg)
45
45
Surface
Area
(mVg)
99
99
Solution *
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
[unction of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Synthetic Groundwater,
function of pH
Ca Electrolyte, COZ Free
Ca Electrolyte, COZ Free
Soil Identification
Attapulgite
(Palygorskite)
Attapulgite
(Palygorskite)
Attapulgite
(Palygorskite)
Attapulgite
(Palygorskite)
Attapulgite
(Palygorskite)
Attapulgite
(Palygorskite)
Attapulgite
(Palygorskite)
Attapulgite
(Palygorskite)
Attapulgite
(Palygorskite)
Attapulgite
(Palygorskite)
Montimorillonite
Montimorillonite
Montimorillonite
Montimorillonite
Montimorillonite
Montimorillonite
Montimorillonite
Montimorillonite
Montimorillonite
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Reference / Comments
Andersson era/. (1982)
Andersson et al. (1982)
Andersson et al. (1982)
Andersson et al. (1982)
Andersson et al. (1982)
Andersson etal. (1982)
Andersson etal. (1982)
Andersson etal. (1982)
Andersson etal. (1982)
Andersson et al. (1982)
Andersson et al. (1982)
Andersson et al. (1982)
Andersson et al. (1982)
Andersson etal. (1982)
Andersson etal. (1982)
Andersson et al. (1982)
Andersson et al. (1982)
Andersson et al. (1982)
Andersson et al. (1982)
Zachara et al. (1992, Fig 6)
Zachara et al. (1992, Fig 6)
J.42
-------
PH
5,1
5.4
5.3
5.5
5.5
5.8
5.8
4.7
4.8
5.0
5.0
5.4
5.7
5.6
5.9
5.9
6.0
6.1
63
6.3
6.4
6.5
UKd
(rnl/g)
2,180
3,680
4,437
7,265
7,108
23,603
22,948
176
176
283
297
708
1,961
2^67
4,283
4,936
7,936
8,586
17,631
19,553
30,963
43,756
Clay
Cent
(wt%)
CEC
(meq/lOOg)
45
45
45
45
45
45
45
45
45
45
, 45
45
45
45
45
45
45
45
45
45
45
45
Surface
Area
(mVg)
99
99
99
99
99
99
99
99
99
99
99
99
99
99
99
99
99
99
99
99
99
99
Solution
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Soil Identification
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Kenoma Clay, <2um
fraction
Reference / Comments
Zachara et al. (1992, Fig 6)
Zachara et al. (1992, Fig 6)
Zachara et al. (1992, Fig 6)
Zachara. et al. (1992, Fig 6)
Zachara et al. (1992, Fig 6)
Zachara et al. (1992, Fig 6)
Zachara et al (1992, Fig 6)
Zachara. et al. (1992, Fig 6)
Zachara et al. (1992, Fig 6)
Zachara et al. (1992, Fig 6)
Zachara el al. (1992, Fig 6)
Zachara et al. (1992, Fig 6)
Zachara et al. (1992, Fig 6)
Zachara et al. (1992, Fig 6)
Zachara et al. (1992, Fig 6)
Zachara et al. (1992, Fig 6)
Zachara et al. (1992, Fig 6)
Zachara et al. (1992, Fig 6)
Zachara etal. (1992, Fig 6)
Zachara et al (1992, Fig 6)
Zachara et al. (1992, Fig 6)
Zachara et al. (1992, Fig 6)
J.43
-------
PH
5.1
5.2
5.2
5.4
5.4
5.6
5.7
5.8
5.9
5.9
6.0
6.1
4.8
4.8
5.1
5.0
5.5
5.5
5.8
6.1
6.1
6.3
6.4
UKd
(ml/g)
508
554
676
874
1,136
1,136
2,143
2,363
9,829
11,966
33,266
37,596
377
399
620
637
1,476
1,603
3,091
6,047
5,823
13,713
13,341
Clay
Cont
(wt%)
CEC
(meq/lOOg)
59
59
59
59
59
59
59
59
59
59
59
59
59
59
59
59
59
59
59
59
59
59
59
Surface
Area
(mVg)
112
112
112
112
112
112
112
112
112
112
112
112
112
112
112
112
112
112
112
112
112
112
112
Solution "
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Ca Electrolyte, CO, Free
Soil Identification
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Reference / Comments
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara eta/. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara etal. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al (1992, Fig 7)
Zachara et al. (1992, Fig 7)
J.44
-------
pH
4.9
5.1
5.1
5.6
5.7
6.7
6.8
7.0
7.0
7.0
7.1
7,1
7.4
7.4
7.7
6.28
6.28
6.28
6.28
6.28
6.09
6.09
UKd
(m!/g)
918
1,168
1,251
2,719
2,928
14,848
13,036
13,827
18,042
19,150
21,771
18,097
26,008
19,488
31,032
3,400
2,800
3,000
11,600
18,600
3,200
8,900
Clay
Cont
(wt%)
CEC
(meq/lOOg)
59
59
59
59
59
59
59
59
59
59
59
59
59
59
Surface
Area
(mz/g)
112
112
112
112
112
112
112
112
112
112
112
112
112
112
Solution
Ca Electrolyte, COZ Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, CO2 Free
Ca Electrolyte, COZ Free
Ca Electrolyte, CO2 Free
Reducing Conditions
Reducing Conditions
Reducing Conditions
Reducing Conditions
Reducing Conditions
Reducing Conditions
Reducing Conditions
Soil Identification
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
Ringold Clay Isolate,
<2um Fraction
PCE Surface Core, 0-8
cm
PCE Surface Core,
9-16 cm
PCE Surface Core,
17-24 cm
PCE Surface Core,
25-32 cm
PCE Surface Core,
33-40 cm
PCE Deep Core, 9-16
cm
PCE Deep Core, 17-24
cm
Reference / Comments
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara. et at. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Zachara et al. (1992, Fig 7)
Sheppard and Thibault
(1988, In Situ)
Sheppard and Thibault
(1988, In Situ)
Sheppard and Thibault
(1988, In Situ)
Sheppard and Thibault
(1988, In Shu)
Sheppard and Thibault
(1988, In Situ)
Sheppard and Thibault
(1988, In Situ)
Sheppard and Thibault
(1988, In Shu)
J.45
-------
PH
6.09
6.09
5.94
6.82
7.28
7.28
7.28
5.7
5.7
5.7
5.7
5.7
4.16
4.99
3.42
3.19
UKd
(ml/g)
9,400
12,500
3,000
8,800
2,600
1,700
700
1,300
2,100
2,000
2,900
870
46
46
900
2,200
560
85.0
170.0
5.3
2.1
day
Cent.
(wt%)
0.5
3.3
3
1.5
CEC
(meq/lOOg)
2.3
3.0
. 2.7
2.9
0.8
1.11
1.82
3.74
1.39
Surface
Area
(mVg)
Solution
Reducing Conditions
Reducing Conditions
Reducing Conditions
Reducing Conditions
Reducing Conditions
Reducing Conditions
Reducing Conditions
Reducing Conditions
Reducing Conditions
Reducing Conditions
Reducing Conditions
Reducing Conditions
Site Borehole Groundwater
Site Borehole Groundwater
Site Borehole Groundwater
She Borehole Groundwata:
Site Borehole Groundwater
Soil Identification
PCE Deep Core, 25-32
cm
PCE Deep Core, 33-40
cm
SCE Surface Core, 0-5
cm
SCE Surface Core,
6-20 cm
SCE Surface Core,
21-25 cm
SCE Surface Core,
26-30 cm
SCE Surface Core,
31-40 cm
PCE Surface Core,
0-40 cm
PCE Deep Core, 40-80
cm
SCE Surface Core,
1-10 cm
SCE Surface Core,
10-30 cm
SCE Surface Core,
30-40 cm
Clay (Glacial Till, Less
Than 5 mm)
Cl:2 (Brown, Slightly
Silly, Less Than 5 mm)
C3 (Dark Brown
Coarse Granular
Deposit, Less Than 5
mm)
C6 (Brown Coarse
Granular Deposit, Less
Than 5 mm)
Sand (Light Brown
Coarse Granular
Deposit, Less Than 5
mm)
A12
A13
A13R
A22
Reference / Comments
Sheppard and Thibault
(1988, In Situ)
Sheppard and Thibault
(1988, In Situ)
Sheppard and Thibault
(1988, In Situ)
Sheppard and Thibault
(1988, In Situ)
Sheppard and Thibault
(1988, In Situ)
Sheppard and Thibault
(1988, In Situ)
Sheppard and Thibault
(1988, In Situ)
Sheppard and Thibault
(1988, Batch)
Sheppard and Thibault
(1988, Batch)
Sheppard and Thibault
(1988, Batch)
Sheppard and Thibault
(1988, Batch)
Sheppard and Thibault
(1988, Batch)
Bell and Bates (1988)
Bell and Bates (1988)
Bell and Bates (1988)
Bell and Bates (1988)
Bell and Bates (1988)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
J.46
-------
pH
3.01
3.19
3.5
3.29
5.42
3.72
3.24
3.93
3.86
4.02
3.83
4.62
4.64
4.67
3.66
4.09
3.61
4.69
3.68
3.75
3.96
4.17
5.53
4.64
5.27
4.51
6.78
4.14
4
4.04
5.85
432
3.87
4.27
4.05
UKd
(ml/g)
1.7
3.7
1.4
1.2
2,200.0
2.3
2.7
8.5
10.1
52
14.0
390.0
180.0
190.0
6.4
39.0
5.3
530.0
6.4
23.0
30.0
980.0
3,600.0
6,300.0
14,000.0
13,000.0
11,000.0
13.0
9.3
320.0
310.0
2,700.0
980.0
290.0
1,500.0
380.0
Clay
Cent
(wt%)
4.5
4.4
3.1
4.7
2.5
2
2.8
3.9
4.9
2.5
7.5
6.2
5.5
12.6
1.2
8.2
3.3
6.4
6.4
5.5
6.1
7.9
3
5.3
2
10.5
4.5
6.4
3.9
73
6.5
3.7
CEC
(meq/lOOg)
1.4
7.92
1
2.1
0.68
0.42
4.71
3.06
3.8
5.69
2.5
8.42
21.4
3.02
15.1
2.39
1.28
6.12
2.54
8.54
11.4
5.04
1.96
2.55
11.4
8.5
15.5
13.3
13.8
11.5
10.5
Surface
Area
(mVg)
Solution
Soil Identification
A23
A31
A32
A42
A52
A53
B13
B14
B23
B23R
B24
B32
B33
B42
B43
B51
B52
B52R
CIS
C14
C22
C23
C32
C33
C42
C43
D13
D13RA
D13RB
E13
E14
E23
E23R
E24
E33
E34
Reference / Comments
Serkiz and Johnson (1994)
Seridz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Seridz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
J.47
-------
pH
5.27
4.87
4.3
4.9
4.69
6.48
4.85
4.77
5.2
4.12
5.91
5.63
4.16
4.99
3.42
3.19
3.01
3.19
3,5
3.29
5.42
3.72
3.24
3.93
3.86
4.02
3.83
4.62
4.64
4.67
3.66
4.09
3.61
4.69
3.68
3.75
UKd
(ml/g)
16,000.0
18,000.0
7,500.0
830.0
160.0
16,000.0
8,700.0
2,900.0
34,000.0
330.0
5,500.0
27,000.0
139.0
361.0
9.46
3.79
1.55
4.43
1.38
1.19
160.0
16.0
2.0
10.4
10.7
4.0
11.3
332.0
212.0
180.0
7.1
20.8
2.6
180.0
5.6
28.3
Clay
Cont
(wt%)
31.8
14.5
15.5
8.1
13
14.2
18.3
17.2
14.2
42.2
16.3
0.5
3.3
3
1.5
4.5
4.4
3.1
4.7
2.5
2
2.8
3.9
4.9
2.5
7.5
6.2
5.5
12.6
1.2
8.2
3.3
6.4
CEC
(meq/lOOg)
20.6
20.6
16.1
8.51
7.48
11.6
15.1
13.6
11.8
19.9
13.3
1.11
1.82
3.74
1.39
1.4
7.92
1
2.1
0.68
0.42
4.71
3.06
3.8
5.69
2.5
8.42 -
21.4
3.02
15.1
2.39
Surface
Area
(mVg)
Solution i
Soil Identification
E41
E42
F12
F13
F22
F23
F32
F33
F42
F43
F52
F53
A12
A13
A13R
A22
A23
A31
A32
A42
A52
A53
B13
B14
B23
B23R
B24
B32
B33
B42
B43
B51
B52
B52R
C13
C14
Reference / Comments
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
J.48
-------
pH
3.96
4.17
5.53
4.64
5.27
4.51
6.78
4.14
4
4.04
4.32
3.87
4.27
4.05
5.27
4.87
4.3
4.9
4.69
6.4S
4.85
4.77
5.2
4.12
5.91
5.63
UKd
(ml/g)
27.4
823.0
540.0
690.0
1,400.0
460.0
690.0
26.6
22.6
650.0
190.0
310.0
360.0
470.0
270.0
870.0
630.0
690.0
2^00.0
1,200.0
950.0
660.0
220.0
910.0
700.0
600.0
960.0
day
Cent
(wt%)
6.4
5.5
6.1
7.9
3
5.3
2
10.5
4.5
3.9
7.3
6.5
3.7
31.8
14.5
15.5
8.1
13
14.2
18.3
17.2
14.2
42.2
16.3
CEC
(meq/lOOg)
1.28
6.12
2.54
8.54
11.4
5.04
1.96
2.55
11.4
8.5
13.3
13.8
11.5
10.5
20.6
20.6
16.1
8.51
7.48
11.6
15.1
13.6
11.8
19.9
13.3
Surface
Area
(mVg)
Solution
Son Identification
C22
C23
C32
C33
C42
C43
D13
D13RA
D13RB
E13
E14
E23R
E24
E33
E34
E41
E42
F12
F13
F22
F23
F32
F33
F42
F43
F52
F53
Reference / Comments
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Seridz and Johnson (1994)
Serkiz and Johnson (1994)
Seridz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Seridz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
Serkiz and Johnson (1994)
J.49
-------
J.6.0 References
I*
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:i
Kaplan, R, J. Seme, A. T. Owen, J. Conca, T. W. Wietsma, and T. L. Gervais. 1996.
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