UNDERSTANDING VARIATION IN PARTITION COEFFICIENT, Kd ...

United States Office of Air and Radiation EPA 402-R-99-004B 

Environmental Protection August 1999 

Agency 

UNDERSTANDING VARIATION IN 

PARTITION COEFFICIENT, K d, VALUES 

Volume II: 

Review of Geochemistry and Available K d Values 

for Cadmium, Cesium, Chromium, Lead, Plutonium, 

Radon, Strontium, Thorium, Tritium ( 3 H), and Uranium

UNDERSTANDING VARIATION IN 

PARTITION COEFFICIENT, K d, VALUES 

Volume II: 

Review of Geochemistry and Available K d Values 

for Cadmium, Cesium, Chromium, Lead, Plutonium, 

Radon, Strontium, Thorium, Tritium ( 3 H), 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

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. 

ii

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 (K d) 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 ( 3 H), and uranium. This two-volume 

report also addresses a void that has existed on this subject in both this Agency and in the user 

community. 

It is important to note that soil scientists and geochemists knowledgeable of sorption 

processes in natural environments have long known that generic or default partition coefficient 

values found in the literature can result in significant errors when used to predict the absolute 

impacts of contaminant migration or site-remediation options. Accordingly, one of the major 

recommendations of this report is that for site-specific calculations, partition coefficient values 

measured at site-specific conditions are absolutely essential. 

For those cases when the partition coefficient parameter is not or cannot be measured, 

Volume II of this document: (1) provides a “thumb-nail sketch” of the key geochemical processes 

affecting the sorption of the selected contaminants; (2) provides references to related key 

experimental and review articles for further reading; (3) identifies the important aqueous- and 

solid-phase parameters controlling the sorption of these contaminants in the subsurface 

environment under oxidizing conditions; and (4) identifies, when possible, minimum and 

maximum conservative partition coefficient values for each contaminant as a function of the key 

geochemical processes affecting their sorption. 

iii

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. 

iv 

Stephen D. Page, Director 

Office of Radiation and Indoor Air

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 

DW89937220-01-03. Lynnette Downing served as the Department of Energy’s Project Officer 

for this Interagency Agreement. PNNL authors involved in this project include: 

Kenneth M. Krupka 

Daniel I. Kaplan 

Gene Whelan 

R. Jeffrey Serne 

Shas V. Mattigod 

v

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 (K d) Values 

401 M Street, SW (6602J) 

Washington, DC 20460 

or 

webmaster.oria@epa.gov 

vi

ABSTRACT 

This two-volume report describes the conceptualization, measurement, and use of the partition (or 

distribution) coefficient, K d, 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 K d values and its use in 

formulating the retardation factor, R f. The K d concept and methods for measurement of K d values 

are discussed in detail in Volume I. Particular attention is directed at providing an understanding 

of: (1) the use of K d values in formulating R f, (2) the difference between the original 

thermodynamic K d 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 K d parameter in contaminant transport codes. A conceptual overview of chemical reaction 

models and their use in addressing technical defensibility issues associated with data from K d 

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 K d 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 ( 3 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. 

vii

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 The K d Model ......................................................... 2.1 

3.0 Methods, Issues, and Criteria for Measuring K d Values .......................... 3.1 

viii 

Page 

3.1 Laboratory Batch Methods ............................................ 3.1 

3.2 Laboratory Flow-Through Method ...................................... 3.1 

3.3 Other Methods ..................................................... 3.2 

3.4 Issues ............................................................ 3.2 

4.0 Application of Chemical Reaction Models .................................... 4.1 

5.0 Contaminant Geochemistry and K d Values ................................... 5.1 

5.1 General ........................................................... 5.1 

5.2 Cadmium Geochemistry and K d 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, K d , Values .................................. 5.10 

5.2.6.1 General Availability of K d Values .............................. 5.10 

5.2.6.2 Look-Up Tables .......................................... 5.11 

5.2.6.2.1 Limits of K d Values with Aluminum/Iron-Oxide Concentrations ..... 5.11 

5.2.6.2.2 Limits of K d Values with Respect to CEC ...................... 5.12 

5.2.6.2.3 Limits of K d Values with Respect to Clay Concentrations .......... 5.12 

5.2.6.2.4 Limits of K d Values with Respect to Concentration of 

Organic Matter .......................................... 5.12

5.2.6.2.5 Limits of K d Values with Respect to Dissolved Calcium, 

Magnesium, and Sulfide Concentrations, and Redox Conditions ..... 5.12 

5.3 Cesium Geochemistry and K d Values .................................... 5.13 


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 


5.3.6.1 General Availability of K d Data ............................... 5.15 

5.3.6.2 Look-Up Tables .......................................... 5.16 

5.3.6.2.1 Limits of K d with Respect to pH ......................... 5.18 

5.3.6.2.2 Limits of K d with Respect to Potassium, Ammonium, 

and Aluminum/Iron-Oxide Concentrations ................. 5.18 

5.4 Chromium Geochemistry and K d Values ................................. 5.18 


Controlling Retardation ........................................... 5.18 






5.4.6.1 General Availability of K d Data ................................ 5.21 

5.4.6.2 Look-Up Tables ........................................... 5.22 


5.4.6.2.2 Limits of K d with Respect to Extractable Iron Content ......... 5.23 

5.4.6.2.3 Limits of K d with Respect to Competing Anion Concentrations .. 5.23 

5.5 Lead Geochemistry and K d Values ..................................... 5.25 


Controlling Retardation ........................................... 5.25 






5.5.6.1 General Availability of K d Data ................................ 5.31 

5.5.6.2 K d Look-Up Tables ........................................ 5.33 


5.5.6.2.2 Limits of K d with Respect to Equilibrium Lead 

ix

Concentrations Extractable Iron Content ........................ 5.34 

5.6 Plutonium Geochemistry and K d Values ................................. 5.34 









5.6.6.2 K d Look-Up Tables ....................................... 5.43 

5.6.6.2.1 Limits of K d with Respect to Clay Content .................. 5.43 

5.6.6.2.2 Limits of K d with Respect to Dissolved Carbonate 

Concentrations ....................................... 5.44 

5.7 Radon Geochemistry and K d Values .................................... 5.44 








5.8 Strontium Geochemistry and K d Values .................................. 5.46 









5.8.6.2 Look-Up Tables .......................................... 5.51 

5.8.6.2.1 Limits of K d with Respect to pH, CEC, and 

Clay Concentrations Values ............................ 5.52 

5.8.6.2.2 Limits of K d with Respect to Dissolved Calcium 

Concentrations ...................................... 5.52 

5.8.6.2.3 Limits of K d with Respect to Dissolved Stable 

Strontium and Carbonate Concentrations .................. 5.53 

5.9 Thorium Geochemistry and K d Values ................................... 5.53 

x







5.9.6 Partition Coefficient, K d, Values ................................... 5.61 


5.9.6.2 Look-Up Tables .......................................... 5.62 

5.9.6.2.1 Limits of K d with Respect to Organic Matter and 

Aluminum/Iron-Oxide Concentrations .................... 5.63 

5.9.6.2.2 Limits of K d with Respect to Dissolved Carbonate 

Concentrations ....................................... 5.63 

5.10 Tritium Geochemistry and K d Values ................................... 5.64 


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, K d , Values ................................. 5.65 

5.11 Uranium Geochemistry and K d Values .................................. 5.65 


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, K d , Values ................................. 5.74 

5.11.6.1 General Availability of K d Data .............................. 5.74 

5.11.6.2 Look-Up Table .......................................... 5.74 

5.11.6.2.1 Limits K d Values with Respect to Dissolved 

Carbonate Concentrations ............................. 5.75 

5.11.6.2.2 Limits of K d Values with Respect to Clay Content and CEC ... 5.76 

5.11.6.2.3 Use of Surface Complexation Models to Predict 

Uranium K d Values .................................. 5.76 

5.12 Conclusions ..................................................... 5.77 

6.0 References ........................................................... 6.1 

xi

Appendix A - Acronyms and Abbreviations ......................................A.1 

Appendix B - Definitions .................................................... B.1 

Appendix C - Partition Coefficients for Cadmium ................................. C.1 

Appendix D - Partition Coefficients for Cesium ...................................D.1 

Appendix E - Partition Coefficients for Chromium ................................. E.1 

Appendix F - Partition Coefficients for Lead ..................................... F.1 

Appendix G - Partition Coefficients for Plutonium ................................G.1 

Appendix H - Partition Coefficients for Strontium .................................H.1 

Appendix I - Partition Coefficients for Thorium ................................... I.1 

Appendix J - Partition Coefficients for Uranium ................................... J.1 

xii

LIST OF FIGURES 

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 µg/l total dissolved U(VI) .............................. 5.70 

Figure 5.6b. Calculated distribution of U(VI) hydrolytic species as a function of pH 

at 1,000 µg/l total dissolved U(VI) ............................... 5.71 

Figure 5.7. Calculated distribution of U(VI) aqueous species as a function of pH 

for the water composition in Table 5.1 ............................ 5.72 

Figure C.1. Relation between cadmium K d values and pH in soils ................... C.5 

Figure D.1. Relation between cesium K d values and CEC .........................D.7 

Figure D.2. Relation between CEC and clay content ............................D.8 

Figure D.3. K d 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 K d values calculated from generalized Fruendlich equation 

(Equations D.3 and D.4) derived from the literature ..................D.16 

xiii 

Page

Figure E.1. Variation of K d for Cr(VI) as a function of pH and DCB extractable 

Iron content without the presence of competing anions ................ E.10 

Figure F.1. Correlative relationship between K d and pH .......................... F.6 

Figure F.2. Variation of K d 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 d) of plutonium ....................................G.12 

Figure G.2. Variation of K d for plutonium as a function of clay content and 

dissolved carbonate concentrations. ...............................G.14 

Figure H.1. Relation between strontium K d values and CEC in soils. ................H.5 

Figure H.2. Relation between strontium K d values for soils with CEC 

values less than 15 meq/100 g. ...................................H.7 

Figure H.3. Relation between strontium K d values and soil clay content ..............H.7 

Figure H.4. Relation between strontium K d values and soil pH .....................H.9 

Figure I.1. Linear regression between thorium K d values and pH for the pH 

range from 4 to 8 ............................................. I.5 

Figure I.2. Linear regression between thorium K d values and pH for the pH 

range from 4 to 8 ............................................. I.8 

Figure J.1. Field-derived K d values for 238 U and 235 U from Serkiz and Johnson (1994) 

plotted as a function of porewater pH for contaminated 

soil/porewater samples ......................................... J.8 


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 K d values for 238 U and 235 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 K d values used for development of K d look-up table ........... J.19 

xiv

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 K d values for cadmium as a function of pH. ............. 5.11 

Table 5.5. Estimated range of K d values (ml/g) for cesium based on CEC 

or clay content for systems containing

Table 5.16. Uranium(VI) aqueous species included in the speciation calculations. ........ 5.69 

Table 5.17. Look-up table for estimated range of K d 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.1. Descriptive statistics of the cadmium K d data set for soils ................... C.3 

Table C.2. Correlation coefficients (r) of the cadmium K d data set for soils .............. C.4 

Table C.3. Look-up table for estimated range of K d values for cadmium based on pH ...... C.5 

Table C.4. Cadmium K d data set for soils. ....................................... C.6 

Table D.1. Descriptive statistics of cesium K d 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 K d 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 K d 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

Table D.10. Estimated range of K d values (ml/g) for cesium based on CEC 

or clay content for systems containing

Table H.3. Simple and multiple regression analysis results involving 

strontium K d values, CEC (meq/100 g), pH, and clay content (percent). ........H.8 

Table H.4. Look-up table for estimated range of K d values for strontium based 

on CEC and pH. .................................................H.10 

Table H.5. Look-up table for estimated range of K d 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 K d data set for soils. .....................................H.12 

Table H.8. Strontium K d data set for pure mineral phases and soils ...................H.16 

Table I.1. Descriptive statistics of thorium K d value data set presented in Section I.3. ...... I.3 

Table I.2. Correlation coefficients (r) of the thorium K d value data set presented 

in Section I.3. .................................................... I.4 

Table I.3. Calculated aqueous speciation of thorium as a function of pH. ............... I.5 

Table I.4. Regression coefficient and their statistics relating thorium K d values and pH. ..... I.6 

Table I.5. Look-up table for thorium K d values (ml/g) based on pH and 

dissolved thorium concentrations. ..................................... I.7 

Table I.6. Data set containing thorium K d values. ................................. I.9 

Table J.1. Uranium K d values (ml/g) listed by Warnecke et al. (1994, Table 1). ......... J.12 

Table J.2. Uranium K d values listed by McKinley and Scholtis (1993, Tables 1, 2, 

and 4) from sorption databases used by different international organizations for 

performance assessments of repositories for radioactive wastes. ............ J.17 

xviii

Table J.3. Geometric mean uranium K d 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 K d values for uranium based on pH. ...... J.22 

Table J.5. Uranium K d values selected from literature for development 

of look-up table. ................................................. J.29 

xix

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, 

K d, 1,2 values and their use in formulating the retardation factor, R f. The contaminant retardation 

factor (R f) 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 soils 3 ). 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 K d 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 K d 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, “K d’s are the coin of the realm in this business.” For better or worse, the K d model is 

integral part of current methodologies for modeling contaminant and radionuclide transport and 

risk analysis. 

The K d concept, its use in fate and transport computer codes, and the methods for the 

measurement of K d values are discussed in detail in Volume I and briefly introduced in Chapters 2 

and 3 in Volume II. Particular attention is directed at providing an understanding of: (1) the use 

of K d values in formulating R f, (2) the difference between the original thermodynamic K d 

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 K d 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. 

1.1

transport codes, and (3) the explicit and implicit assumptions underlying the use of the K d 

parameter in contaminant transport codes. 

The K d 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 K d 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 K d 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 K d 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 K d 

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 10 2 to 10 6 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 

K d 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 K d value even when selected within the range of technically-defensible, limiting 

minimum and maximum K d values. For those situations that are sensitive to the selected K d value, 

site-specific K d values are essential. 

The K d is usually a measured parameter that is obtained from laboratory experiments. The 

5 general methods used to measure K d values are reviewed. These methods include the batch 

laboratory method, the column laboratory method, field-batch method, field modeling method, 

and K oc 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 K d values and modeling of adsorption processes is discussed in detail in Volume I 

and briefly described in Chapter 4 of Volume II. 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 K d 

construct. The use of geochemical modeling to address quality assurance and technical 

defensibility issues concerning available K d data and the measurement of K d values is also 

1.2

discussed. The geochemical modeling review includes a brief description of the EPA’s 

MINTEQA2 geochemical code and a summary of the types of conceptual models it contains to 

quantify adsorption reactions. The status of radionuclide thermodynamic and contaminant 

adsorption model databases for the MINTEQA2 code is also reviewed. 

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 K d 

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 ( 3 H), 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: 

C Cations - cadmium, cesium, plutonium, strontium, thorium, and uranium(VI). 

C Anions - chromium(VI) (as chromate) and uranium(VI). 

C Radionuclides - cesium, plutonium, radon, strontium, thorium, tritium ( 3 H), and uranium. 

C Conservatively transported contaminants - tritium ( 3 H) and radon. 

C Nonconservatively transported contaminants - other than tritium ( 3 H) and radon. 

C Redox sensitive elements - chromium, plutonium, and uranium. 

The general geochemical behaviors discussed in this report can be used by analogy to estimate the 

geochemical interactions of similar elements for which data are not available. For example, 

contaminants present primarily in anionic form, such as Cr(VI), tend to adsorb to a limited extent 

to soils. Thus, one might generalize that other anions, such as nitrate, chloride, and 

U(VI)-anionic complexes, would also adsorb to a limited extent. Literature on the adsorption of 

these 3 solutes show no or very little adsorption. 

The concentration of contaminants in groundwater is controlled primarily by the amount of 

contaminant present at the source; rate of release from the source; hydrologic factors such as 

dispersion, advection, and dilution; and a number of geochemical processes including aqueous 

geochemical processes, adsorption/desorption, precipitation, and diffusion. To accurately predict 

contaminant transport through the subsurface, it is essential that the important geochemical 

processes affecting contaminant transport be identified and, perhaps more importantly, accurately 

described in a mathematically and scientifically defensible manner. Dissolution/precipitation and 

adsorption/desorption are usually the most important processes affecting contaminant interaction 

with soils. Dissolution/precipitation is more likely to be the key process where chemical 

nonequilibium exists, such as at a point source, an area where high contaminant concentrations 

1.3

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., Fick’s law). Diffusion is a dominant transport mechanism when 

advection is insignificant, and is usually a negligible transport mechanism when water is being 

advected in response to various forces. 

1.4

2.0 The K d Model 

The simplest and most common method of estimating contaminant retardation is based on the 

partition (or distribution) coefficient, K d. The K d 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 K d 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 K d 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 K d 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 K d values from the literature for use in transport codes. 

The partition coefficient, K d, is defined as the ratio of the quantity of the adsorbate adsorbed per 

mass of solid to the amount of the adsorbate remaining in solution at equilibrium. For the 

reaction 

the mass action expression for K d is 

K d ' 

A % C i ' A i , (2.1) 

Mass of Adsorbate Sorbed 

Mass of Adsorbate in Solution ' A i 

C i 

where A = free or unoccupied surface adsorption sites 

C i = total dissolved adsorbate remaining in solution at equilibrium 

A i = amount of adsorbate on the solid at equilibrium. 

The K d is typically given in units of ml/g. Describing the K d in terms of this simple reaction 

assumes that A is in great excess with respect to C i and that the activity of A i is equal to 1. 

Chemical retardation, R f, is defined as, 

R f ' v p 

v c 

where v p = velocity of the water through a control volume 

v c = velocity of contaminant through a control volume. 

2.1 

(2.2) 

, (2.3) 

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

etardation 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 K d 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 R f is defined as 

R f ' 1 % D b 

n e 

where D b = porous media bulk density (mass/length 3 ) 

n e = effective porosity of the media at saturation. 

The K d 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 K d values should be used for site-specific contaminant and risk 

assessment calculations. Ideally, site-specific K d values should be measured for the range of 

aqueous and geological conditions in the system to be modeled. However, literature-derived K d 

values are commonly used for screening calculations. Suitable selection and use of literaturederived 

K d values for use in screening calculations of contaminant transport is not a trivial matter. 

Among the assumptions implicit with the K d 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 K d values should be used only to predict transport in systems chemically identical to 

those used in the laboratory measurement of the K d. Variation in either the soil or aqueous 

chemistry of a system can result in extremely large differences in K d values. 

A more robust approach than using a single K d to describe the partitioning of contaminants 

between the aqueous and solid phases is the parametric-K d model. This model varies the K d value 

according to the chemistry and mineralogy of the system at the node being modeled. The 

parametric-K d value, unlike the constant-K d 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-K d 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-K d 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 K d value. Thus, the code must 

2.2 

K d 

(2.4)

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 K d 

value and/or retardation factor, R f. Generally, computer codes that can accommodate the 

parametric-K d model use a chemical subroutine to update the K d value used to determine the R F, 

when called by the main transport code. The added complexity in solving the transport equation 

with the parametric-K d sorption model and its empirical nature may be the reasons this approach 

has been used sparingly. 

Mechanistic models explicitly accommodate for the dependency of K d values on contaminant concentration, 

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. 

2.3

3.0 Methods, Issues, and Criteria for Measuring K d Values 

There are 5 general methods used to measure K d values: the batch laboratory method, laboratory 

flow-through (or column) method, field-batch method, field modeling method, and K oc 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 K d values from experimental data. Consequently, it 

is not only common, but expected that K d values measured by different methods will produce 

different values. 

3.1 Laboratory Batch Method 

Batch tests are commonly used to measure K d 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 K d 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, K d 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 K d 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 K d 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 K d 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

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 K d value can be calculated from the retardation 

factor. It is frequently useful to compare the back-calculated K d 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 K d value needs to 

be calculated from the retardation factor. Another advantage is that the column experiment 

provides a much closer approximation of the physical conditions and chemical processes 

occurring in the field site than a batch sorption experiment. Column experiments permit the 

investigation of the influence of limited spatial and temporal (nonequilibium) contact between 

solute and solid have on contaminant retardation. Additionally, the influence of mobile colloid 

facilitated transport and partial saturation can be investigated. A third advantage is that both 

adsorption or desorption reactions can be studied. The predominance of 1 mechanism of 

adsorption or desorption over another cannot be predicted a priori and therefore generalizing the 

results from 1 set of laboratory experimental conditions to field conditions is never without some 

uncertainty. Ideally, flow-through column experiments would be used exclusively for determining 

K d 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 K oc method, in-situ batch method, and the field 

modeling method. The K oc method is a very effective indirect method of calculating K d 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 K d 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 K d values uses 

groundwater monitoring data and source term data to calculate a K d value. One key drawback to 

this technique is that it is very model dependent. Because the calculated K d value are model 

dependent and highly site specific, the K d values must be used for contaminant transport 

calculations at other sites. 

3.4 Issues 

A number of issues exist concerning the measurement of K d values and the selection of K d values 

from the literature. These issues include: using simple versus complex systems to measure K d 

values, field variability, the “gravel issue,” and the “colloid issue.” Soils are a complex mixture 

3.2

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 K d 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 K d values based 

on experiments conducted with the 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 K d values 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 K d value is almost never sufficient for an entire 

study site and should change as chemically important environmental conditions change. Three 

approaches used to vary K d values in transport codes are the K d look-up table approach, the 

parametric-K d approach, and the mechanistic K d approach. The extent to which these approaches 

are presently used and the ease of incorporating them into a flow model varies greatly. 

Parametric-K d values typically have limited environmental ranges of application. Mechanistic K d 

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-K d model interfaced with 

transport codes is the look-up table. In K d look-up tables, separate K d 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 K d 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 K d values can result in appreciable different K d 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

(e.g., soil). An increasing body of evidence indicates that under some subsurface conditions, 

components of the solid phase may exist as colloids 1 that may be transported with the flowing 

water. Subsurface mobile colloids originate from (1) the dispersion of surface or subsurface soils, 

(2) decementation of secondary mineral phases, and (3) homogeneous precipitation of groundwater 

constituents. Association of contaminants with this additional mobile phase may enhance 

not only the amount of contaminant that is transported, but also the rate of contaminant transport. 

Most current approaches to predicting contaminant transport ignore this mechanism not because 

it is obscure or because the mathematical algorithms have not been developed, but because little 

information is available on the occurrence, the mineralogical properties, the physicochemical 

properties, or the conditions conducive to the generation of mobile colloids. There are 2 primary 

problems associated with studying colloid-facilitated transport of contaminants under natural 

conditions. First, it is difficult to collect colloids from the subsurface in a manner which 

minimizes or eliminates sampling artifacts. Secondly, it is difficult to unambiguously delineate 

between the contaminants in the mobile-aqueous and mobile-solid phases. 

Often K d values used in transport models are selected to provide a conservative estimate of 

contaminant migration or health effects. However, the same K d 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 K d from the upper range of literature values. 

It is incumbent upon the transport modeler to understand the strengths and weaknesses of the 

different K d methods, and perhaps more importantly, the underlying assumption of the methods in 

order to properly select K d values from the literature. The K d 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 K d 

value(s) for contaminant transport modeling. 

1 A colloid is any fine-grained material, sometimes limited to the particle-size range of 

4.0 Application of Chemical Reaction Models 

Computerized chemical reaction models based on thermodynamic principles may be used to 

calculate processes such as aqueous complexation, oxidation/reduction, adsorption/desorption, 

and mineral precipitation/dissolution for contaminants in soil-water systems. The capabilities of a 

chemical reaction model depend on the models incorporated into its computer code and the 

availability of thermodynamic and/or adsorption data for aqueous and mineral constituents of 

interest. Chemical reaction models, their utility to understanding the solution chemistry of 

contaminants, and the MINTEQA2 model in particular are described in detail in Chapter 5 of 

Volume I. 

The MINTEQA2 computer code is an equilibrium chemical reaction model. It was developed 

with EPA funding by originally combining the mathematical structure of the 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, gasphase 

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, soilwater 

systems. Of the contaminants selected for consideration in this project [chromium, 

cadmium, cesium, tritium ( 3 H), lead, plutonium, radon, strontium, thorium, and uranium], the 

MINTEQA2 thermodynamic database contains speciation and solubility reactions for chromium, 

including the valence states Cr(II), Cr(III), and Cr(VI); cadmium; lead; strontium; and uranium, 

including the valence states U(III), U(IV), U(V), and U(VI). Some of the thermodynamic data in 

the EPA version have been superseded in other users’ databases by more recently published data. 

The MINTEQA2 code includes 7 adsorption model options. The non-electrostatic adsorption 

act 

models include the activity Kd , 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, K d, 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 K d value. The user however must make assumptions concerning the composition and 

mass of the dominant sorptive substrate, and supply the adsorption parameters for surface- 

4.1

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 K d 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 K d values for barium, beryllium, cadmium, Cr(III), 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 K d study. If the appropriate data are reported, it is likely that a user 

could hand calculate a composition-based K d 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 K d 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 K d measurements to determine 

if the measurement had been affected by processes such as mineral precipitation which might have 

compromised the derived K d values. Although chemical reaction models cannot be used to 

predict K d values, they can provide aqueous speciation and solubility information that is 

exceedingly valuable in the evaluation of K d values selected from the literature and/or measured in 

the laboratory. 

4.2

5.0 Contaminant Geochemistry and K d Values 

The important geochemical factors affecting the sorption 1 of cadmium (Cd), cesium (Cs), 

chromium (Cr), lead (Pb), plutonium (Pu), radon (Rn), strontium (Sr), thorium (Th), tritium ( 3 H), 

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 K d 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 groundwaters, 

will likely be more reducing that surface water. As explained later in this chapter, the adsorption 

and solubility of certain contaminants and radionuclides may be significantly different under 

reducing groundwater conditions compared to oxidizing conditions. However, it was necessary 

to limit the scope of this review to oxidizing conditions. Use of the water composition in 

Table 5.1 does not invalidate the aqueous speciation calculations discussed later in this chapter 

relative to the behavior of the selected contaminants in oxidizing and transitional groundwater 

systems. The calculations demonstrate what complexes might exist for a given contaminant in any 

oxidizing water as a function of pH and the specified concentrations of each inorganic ligand. If 

the concentration of a complexing ligand, such as phosphate, is less for a site-specific 

groundwater compared to that used for our calculations, then aqueous complexes containing that 

contaminant and ligand may be less important for that water. 

Importantly, water composition in Table 5.1 has a low ionic strength and contains no natural (e.g., 

humic or fulvic acids 1 ) 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 0 (gas), and tritium as tritiated water, HTO (T = tritium, 3 H). 

Throughout this chapter, particular attention will be directed at identifying the important aqueousand 

solid-phase parameters controlling retardation 2 of contaminants by sorption in soil. This 

information was used to guide the review and discussion of published K d 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 alkalinesoluble 

portion which remains in solution at low pH and is of lower molecular weight” (Gascoyne, 

1982). 

2 Retarded or attenuated (i.e., nonconservative) transport means that the contaminant moves 

slower than water through geologic material. Nonretarded or nonattenuated (i.e., conservative) 

transport means that the contaminant moves at the same rate as water. 

5.2

Table 5.1. Estimated mean composition of river 

water of the world from Hem (1985). 1 

Dissolved Constituent 

5.3 

Total Concentration 

mg/l mol/l 

Silica, as H 4SiO 4 20.8 2.16 x 10 -4 

Ca 15 3.7 x 10 -4 

Mg 4.1 1.7 x 10 -4 

Na 6.3 2.7 x 10 -4 

K 2.3 5.9 x 10 -5 

Inorganic Carbon, as CO 3 57 9.5 x 10 -4 

SO 4 11 1.1 x 10 -4 

Cl 7.8 2.2 x 10 -4 

F 1 5 x 10 -5 

NO 3 1 2 x 10 -5 

PO 4 0.0767 8.08 x 10 -7 

1 Most values from this table were taken from Hem (1985: Table 3, 

Column 3). Mean concentrations of total dissolved fluoride and 

phosphate are not listed in Hem (1985, Table 3). The concentration of 

dissolved fluoride was taken from Hem (1985, p. 120) who states that 

the concentration of total dissolved fluoride is generally less than 

1.0 mg/l for most natural waters. Hem (1985, p. 128) lists 25 µg/l for 

average concentration of total dissolved phosphorous in river water 

estimated by Meybeck (1982). This concentration of total phosphorus 

was converted to total phosphate (PO 4) listed above.

Element 

Table 5.2. Concentrations of contaminants used in the aqueous 

species distribution calculations. 

Total Conc. 

(µg/l) 

Reference for Concentration of Contaminant 

Used in Aqueous Speciation Calculations 

Cd 1.0 Hem (1985, p. 142) lists this value as a median concentration of dissolved 

cadmium based on the reconnaissance study of Duram et al. (1971) of metal 

concentrations in surface waters in the United States. 

Cs -- Distribution of aqueous species was not modeled, because mobility of dissolved 

cesium is not significantly affected by complexation (see Section 5.3). 

Cr 1.4 Hem (1985, p. 138) lists this value as an average concentration estimated by 

Kharkar et al. (1968) for chromium in river waters. 

Pb 1.0 Hem (1985, p. 144) lists this value as an average concentration estimated by 

Duram et al. (1971) for lead in surface-water samples from north- and southeastern 

sections of the United States. 

Pu 3.2 x 10 -7 This concentration is based on the maximum activity of 239,240 Pu measured by 

Simpson et al. (1984) in 33 water samples taken from the highly alkaline Mono 

Lake in California. 

Rn -- Aqueous speciation was not calculated, because radon migrates as a dissolved gas 

and is not affected by complexation (see Section 5.7). 

Sr 110 Hem (1985, p. 135) lists this value as the median concentration of strontium for 

larger United States public water supplies based on analyses reported by Skougstad 

and Horr (1963). 

Th 1.0 Hem (1985, p. 150) gives 0.01 to 1 µg/l as the range expected for thorium 

concentrations in fresh waters. 

3 H -- Aqueous speciation was not calculated, because tritium ( 3 H) migrates as tritiated 

U 0.1 and 

1,000 

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 µg/l) of total dissolved 

uranium were used to model the species distributions. Hem (1985, p. 148) gives 

0.1 to 10 µg/l 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 µg/l. 

5.4

5.2 Cadmium Geochemistry and K d Values 


Controlling Retardation 

The dominant cadmium aqueous species in groundwater at pH values less than 8.2 and containing 

moderate to low concentrations of sulfate (

5.2.3 Aqueous Speciation 

Cadmium forms soluble complexes with inorganic and organic ligands resulting in an increase of 

cadmium mobility in soils (McLean and Bledsoe, 1992). The distribution of cadmium aqueous 

species was calculated using the water composition described in Table 5.1 and a concentration of 

1 µg/l 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 MINTEQA2 calculations 

indicate that cadmium speciation is relatively simple. In groundwaters of pH values less than 6, 

essentially all of the dissolved cadmium is expected to exist as the uncomplexed Cd2+ ion 

(Figure 5.1). The aqueous species included in the MINTEQA2 calculations are listed in 

+ 

Table 5.3. As the pH increases between 6 and 8.2, cadmium carbonate species [CdHCO3 and 

" 

CdCO3 (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 CdCO3 (aq). The species 

" + + + 

CdSO4 (aq), CdHCO3, 

CdCl , and CdOH are also present, but at much lower concentrations. 

The species distribution illustrated in Figure 5.1 does not change if the concentration of total 

dissolved cadmium is increased from 1 to 1,000 µg/l. 

Table 5.3. Cadmium aqueous species included 

in the speciation calculations. 

Aqueous Species 

Cd 2+ 

CdOH + " - 2- 

, Cd(OH) 2 (aq), Cd(OH)3, 

Cd(OH)4 , Cd2OH 3+ 

+ " 4- 

CdHCO3, CdCO3 (aq), Cd(CO3) 3 

" 2- 

CdSO4 (aq), Cd(SO4) 2 

+ 

CdNO3 CdCl + " - " 

, CdCl2 (aq), CdCl3, 

CdOHCl (aq) 

CdF + " 

, CdF2 (aq) 

5.6

Percent Distribution 

100 

90 

80 

70 

60 

50 

40 

30 

20 

10 

0 

CdCl + 

+ 

CdHCO3 

3 4 5 6 7 8 9 10 

5.7 

Cd 2+ 

pH 

o 

CdCO3 (aq) 

o 

CdSO4 (aq) 

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 µg/l 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.

5.2.4 Dissolution/Precipitation/Coprecipitation 

Lindsay (1979) calculated the relative stability of cadmium compounds. His calculations show 

that at pH values less than 7.5, most cadmium minerals are more soluble than cadmium 

concentrations found in oxic soils (10 -7 M), indicating that cadmium at these concentrations is not 

likely to precipitate. At pH levels greater than 7.5, the solubilities of Cd 3(PO 4) 2 or CdCO 3 may 

control the concentrations of cadmium in soils. Cavallaro and McBride (1978) and McBride 

(1980) demonstrated that otavite, CdCO 3, 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 CdCO 3. 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)CO 3. 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). Serne (1977) studied the effect of 

oxidized and reduced sediment conditions on the release of cadmium from dredged sediments 

collected from the San Francisco Bay. Greater than 90 percent of the cadmium in the reduced 

sediment [sediment incubated in the presence of low O 2 levels (Eh350 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 or sediments and gets oxidized 

when exposed to air. 1 The pyrite oxidizes to form FeSO 4, 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/l), the adsorption of cadmium often 

correlates with the CEC of the soil (John, 1971; Levi-Minzi et al., 1976; McBride et al., 1981; 

Navrot et al., 1978; Petruzelli et al., 1978). During cation exchange, cadmium generally 

exchanges with adsorbed calcium and magnesium (McBride et al., 1982). The ionic radius of 

Cd 2+ is comparable to that of Ca 2+ and, to a lesser extent, Mg 2+ . 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 Cd 2+ 

and possibly CdOH + may adsorb to aluminum- and iron-oxide minerals (Balistrieri and Murray, 

1981; Davis and Leckie, 1978). 

As with other cationic metals, cadmium adsorption exhibits pH dependency. The effect of pH on 

cadmium adsorption by soils (Huang 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 

al., 1976), this effect is probably due to the CEC of the organic material rather than to 

complexation by organic ligands (Singh and Sekhon, 1977). In fact, removal of organic material 

from soils does not markedly reduce cadmium adsorption and may enhance adsorption (Petruzelli 

et al., 1978). Clay minerals with adsorbed humic acids (organo-clay complexes) do not adsorb 

cadmium in excess of that expected for clay minerals alone (Levy and Francis, 1976). 

5.2.6 Partition Coefficient, K d , Values 

5.2.6.1 General Availability of K d Data 

A total of 174 cadmium K d values were found in the literature and included in the data base used 

to create the look-up tables. 1 The cadmium K d values as well as the ancillary experimental data 

are presented in Appendix C. Data included in this table were from studies that reported K d 

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 (

5.2.6.2 Look-Up Tables 

One cadmium K d 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 K d, a correlation coefficient value of 0.75. A 

detailed explanation of the approach used in selecting the K d values used in the table is presented 

in Appendix C. Briefly, it involved conducting a regression analysis between pH and K d 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 K d values for each of the 3 pH categories. 

The cause for this is likely that there are several other soil parameters influencing the K d in 

addition to pH. Unfortunately, the correlations between the cadmium K d values and the other soil 

parameters in this data set were not significant (Appendix C). 

5.2.6.2.1 Limits of K d Values With Respect to Aluminum/Iron-Oxide Concentrations 

The effect of iron-oxide concentrations on cadmium K d values was evaluated using the data 

presented in Appendix C. Of the 174 cadmium K d 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 K d 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 K d values 

because oxide minerals provide adsorption (surface complexation) sites. 

Table 5.4. Estimated range of K d 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 (

5.2.6.2.2 Limits of K d Values with Respect to CEC 

The effect of CEC on cadmium K d values was evaluated using the data presented in Appendix 

C. Of the 174 cadmium K d 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 K d 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 K d values because cadmium can adsorb to minerals via cation exchange. 

5.2.6.2.3 Limits of K d Values with Respect to Clay Content 

The effect of clay content on cadmium K d values was evaluated using the data presented in 

Appendix C. Of the 174 cadmium K d 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 K d 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 K d 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 K d Values with Respect to Concentration of Organic Matter 

The effect of organic matter concentration, as approximated by total organic carbon, on 

cadmium K d values was evaluated using the data presented in Appendix C. Of the 174 

cadmium K d 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 K d 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 K d 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 K d 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 K d values. It 

was anticipated that sulfides would induce cadmium precipitation, thereby increasing cadmium K d 

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 and K d Values 



The aqueous speciation of cesium in groundwater is among the simplest of the contaminants being 

considered in this study. Cesium forms few stable complexes and is likely to exist in groundwater 

as the uncomplexed Cs + ion, which adsorbs rather strongly to most minerals, especially mica-like 

clay minerals. The extent to which adsorption will occur will depend on (1) the concentration of 

mica-like clays in the soil, and (2) the concentration of major cations, such as K + which has a 

small ionic radius as Cs + , that can effectively compete with Cs + for adsorption sites. 

5.3.2 General Geochemistry 

Cesium (Cs) exists in the environment in the +1 oxidation state. Stable cesium is ubiquitous in the 

environment with concentrations in soils ranging between 0.3 and 25 mg/kg (Lindsay, 1979). The 

only stable isotope of cesium is 133 Cs. Fission products include 4 main cesium isotopes. Of these, 

only 134 Cs [half life (t ½) = 2.05 y], 135 Cs (t ½ = 3 x 10 6 y), and 137 Cs (t ½ = 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 134 Cs, 135 Cs and/or 137 Cs has been 

identified at 9 of the 45 Superfund National Priorities List (NPL) sites. 


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., Ca 2+ ) 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


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. 


In general, most soils sorb cesium rather strongly (Ames and Rai, 1978). Some mica-like 

minerals, such as illite {(K,H 3O)(Al,Mg,Fe) 2(Si,Al) 4O 10[(OH) 2,H 2O]} and vermiculite 

[(Mg,Fe,Al) 3(Si,Al) 4O 10(OH) 2·4H 2O], 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 + , Ca 2+ , and Mg 2+ , 

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

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. 



Three generalized, simplifying assumptions were established for the selection of cesium K d 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: 

C 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. 

C Cesium adsorption into mica-like minerals occurs much more readily than desorption. 

Thus, K d values, which are essentially always derived from adsorption studies, will greatly 

overestimate the degree to which cesium will desorb from these surfaces. 

C 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 K d 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 K d 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 (

organic chelates (e.g., EDTA). Initially, attempts were made to include in the K d 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, micalike 

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 K d values. The descriptive statistics of the cesium K d 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 K d values for the look-up table. The K d 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 K d 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 K d values and the solid phase parameters. Thus, the K d 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 K d 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 K d 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

Table 5.5. Estimated range of K d values (ml/g) for cesium based 

on CEC or clay content for systems containing 

5.3.6.2.1 Limits of K d Values with Respect to pH 

Of the 177 cesium K d 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 K d 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 K d values 

were not held constant throughout this data set. 

5.3.6.2.2 Limits of K d 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 K d data set (Appendix D). It was 

anticipated that dissolved potassium and ammonium would compete with cesium for adsorption 

sites, thereby decreasing K d values. The presence of aluminum and/or iron oxides in the solid 

phase was expected to increase cesium K d values. 

5.4 Chromium Geochemistry and K d Values 



A plume containing high concentrations of chromium is more likely to be composed of Cr(VI) 

than Cr(III) because the former is less likely to adsorb or precipitate to the solid phase. 

Chromium(VI) is also appreciably more toxic than Cr(III). It exhibits significant subsurface 

mobility in neutral and basic pH environments. In acid environments, Cr(VI) may be moderately 

adsorbed by pH-dependent charge minerals, such as iron- and aluminum-oxide minerals. The 

reduction of Cr(VI) to Cr(III) by ferrous iron, organic matter, and microbes is generally quite 

rapid whereas the oxidation of Cr(III) to Cr(VI) by soil manganese oxides or dissolved oxygen is 

kinetically slower. The most important aqueous- and solid-phase parameters controlling 

retardation of chromium include redox status, pH, and the concentrations of aluminum- and ironoxide 

minerals and organic matter. 


Chromium is found in the environment primarily in the +3 and +6 oxidation states. The 

geochemical behavior and biological toxicity of chromium in these 2 oxidation states are 

profoundly different. Chromium(VI) tends to be soluble, forms anionic or neutral dissolved 

species, can be very mobile, and is acutely toxic (Nriagu and Nieboer, 1988). In contrast, Cr(III) 

tends to precipitate, forms cationic dissolved species, is immobile under moderately alkaline to 

slightly acidic conditions, and is relatively nontoxic. The primary human activities leading to the 

introduction of chromium into the environment are ore processing, plating operations, and 

5.18

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). 


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(III) 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(III) exists predominantly as Cr3+ below pH 3.5 in a Cr(III)-H2O system. 

With increasing pH, hydrolysis of Cr3+ yields CrOH2+ + " - 

, Cr(OH) 2, 

Cr(OH)3(aq), 

and Cr(OH)4, 

4+ 5+ 

Cr2(OH) 2 , and Cr3(OH) 4 . At higher chromium concentrations, polynuclear species, such as 

4+ 5+ " 

Cr2(OH) 2 and Cr3(OH) 4 , can form slowly at 25 C (Baes and Mesmer, 1976). Chromium(VI) 

hydrolyses extensively, forming primarily anionic species. These species are HCrO4 2- 2- 

CrO4 (chromate), and Cr2O7 (dichromate) (Baes and Mesmer, 1976; Palmer and Wittbrodt, 

1991; Richard and Bourg, 1991). Palmer and Puls (1994) presented some Cr(VI) speciation 

2- 

diagrams representative of groundwater conditions. They showed that above pH 6.5, CrO4 - 

generally dominates. Below pH 6.5, HCrO4 dominates when the total concentration of dissolved 

2- 

Cr(VI) is low (

(tarapacaite) in chromium sludge from a plating facility. They also reported that BaCrO 4 formed 

a complete solid solution with BaSO 4. They concluded that these solid solutions can be a major 

impediment to the remediation of chromium-contaminated sites by pump-and-treat technologies. 

Chromium(VI) is a strong oxidant and is rapidly reduced in the presence of such common electron 

donors as aqueous Fe(II), ferrous iron minerals, reduced sulfur, microbes, and organic matter 

(Bartlett and Kimble, 1976; Nakayama et al., 1981). Studies indicate that Cr(VI) can be reduced 

to Cr(III) by ferrous iron derived from magnetite (Fe 3O 4) and ilmenite (FeTiO 3) (White and 

Hochella, 1989), hematite (Fe 2O 3) (Eary and Rai, 1989), 1 and pyrite (FeS 2) (Blowes and Ptacek, 

1992). 

The reduction of Cr(VI) by Fe(II) is very rapid. The reaction can go to completion in a matter of 

minutes (Eary and Rai, 1989). The rate of reduction of Cr(VI) increases with decreasing pH and 

increasing initial Cr(VI) and reductant concentrations (Palmer and 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(III) can precipitate with Fe(III) to form a solid solution with 

the general composition Cr xFe 1-x(OH) 3 (Sass and Rai, 1987). The solubility of chromium in this 

solid solution decreases as the mole fraction of Fe(III) increases. The oxidation reaction proceeds 

much more slowly than the reduction reaction; the former reaction requires months for 

completion (Eary and Rai, 1987; Palmer and Puls, 1994). Only 2 constituents in the environment 

are known to oxidize Cr(III): dissolved oxygen and manganese-dioxide minerals [e.g., pyrolusite 

($-MnO 2)]. Eary and Rai (1987) reported that the rate of Cr(III) oxidation was much greater in 

the presence of manganese-dioxide minerals than dissolved oxygen. 


The extent to which Cr(III) sorbs to soils is appreciably greater than that of Cr(VI) because the 

former exists in groundwater as a cation, primarily as Cr 3+ (and its complexed species), whereas 

2- - 

the latter exists as an anion, primarily as CrO4 or HCrO4. 

Most information on Cr(VI) 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 ("-Al2O3) and 

amorphous iron oxide [Fe2O3·H2O(am)] at low to medium pH values (pH 2 to 7) and adsorbs 

weakly to silica (SiO2) 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 (IEP) 2 of 

these minerals. When the pH of the system is greater than the isoelectric point, the mineral has a 

net negative charge. When the pH is below the isoelectric point, the mineral has a net positive 

1 Eary and Rai (1989) attributed the reduction of Cr(VI) to Cr(III) by hematite (Fe2O 3) as 

containing having trace quantities of Fe(II). 

2 The isoelectric point (IEP) of a mineral is the pH at which it has a net surface charge of zero. 

More precisely, it is the pH at which the particle is electrokinetically uncharged. 

5.20

charge. Hence, anion adsorption generally increases as the pH becomes progressively lower than 

the isoelectric point. The isoelectric point of gibbsite ("-Al 2O 3) is 9.1, amorphous iron oxide 

[Fe 2O 3·H 2O (am)] is 8.1, and silica is 2.0 (Stumm and Morgan, 1981). 

The presence of competing and, less commonly, complexing ions may significantly alter chromate 

adsorption. Although sulfate is adsorbed less strongly on 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., 

" " 

(aq) and KHCrO4(aq)] 

is suggested as 1 possible mechanism for removal of Cr(VI) by 

CaCrO 4 

Fe 2O 3·H 2O (am) (Leckie et al., 1980). 

Adsorption of Cr(III) to soils has received only a nominal amount of research attention. The 

reason for this may be that sorption of Cr(III) by soil is commonly ascribed to solid phase 

formation. Chromium(III) 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

K d). Soils containing Mn oxides oxidize Cr(III) into Cr(VI) form thus resulting in lower 

K d values. The relation between oxide/hydroxide contents of iron and manganese and 

their effects on K d have not been adequately quantified except for a few soils. 

C 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 K d values or any correlative relationships between K d and the oxide contents. Studies by 

Stollenwerk and Grove (1985) and Sheppard et al. (1987) using soils showed that K d 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 d 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(III) oxidation state. The reduced Cr 

has been shown to coprecipitate with ferric hydroxide. Therefore, observed removal of Cr(VI) 

from solution when contacted with chromium-reductive soils may stem from both adsorption and 

precipitation reactions. Similarly, Rai et al. (1988) also showed that certain soils containing 

manganese oxides may oxidize Cr(III) to Cr(VI). Depending on solution concentrations, the 

oxidized form (+6) of chromium may also precipitate in the form of Ba(S,Cr)O 4. Such complex 

geochemical behavior chromium in soils implies that depending on the properties of a soil, the 

measured K d values may reflect both adsorption and precipitation reactions. 

2- 2- - 2- - 

Adsorption studies have shown that competing anions such as SO4 , CO3 /HCO3, 

HPO4 , H2PO4 - - 

NO3 and Cl , significantly reduce Cr(VI) adsorption on oxide minerals and soils (Leckie et al., 

1980; MacNaughton, 1977; Rai et al., 1986; Rai et al., 1988; Stollenwerk and Grove, 1985). 

The data regarding the effects of soil organic matter on Cr(VI) adsorption are rather sparse. 

In 1 study (Stollenwerk and Grove, 1985) which evaluated the effects of soil organic matter on 

adsorption of Cr(VI), the results indicated that organic matter did not influence Cr(VI) adsorption 

properties (see Appendix E). 

5.4.6.2 K d Look-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

2- 2- - 

competing anions such as SO4 , and CO3 /HCO3 . 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. 


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 K d with pH. The K d values for all 4 soils were observed to decline with increasing pH 

and at pH values beyond about 9, K d values for Cr(VI) are #1 ml/g (see Appendix E). 

5.4.6.2.2 Limits of K d 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(III). Therefore the 

measured K d values for such soils reflect both redox-mediated precipitation and adsorption 

phenomena. The data also show that soils with DCB extractable iron contents of about 

0.25 mmol/g or less do not appear to reduce Cr(VI). Therefore, 3 ranges of DCB extractable iron 

contents were selected which represent the categories of soils that definitely reduce ($0.3 

mmol/g), probably reduce (0.26 - 0.29 mmol/g), and do not reduce (#0.25 mmol/g) Cr(VI) to 

Cr(III) form. 

5.4.6.2.3 Limits of K d 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/l), the chromium K d values are reduced by about an order of 

magnitude as compared to a noncompetitive condition. Therefore, a sulfate concentration of 

about 2 x 10 -3 M (191.5 mg/l) has been used as a limit at which an order of magnitude reduction 

in K d values are expected. Four ranges of soluble sulfate concentrations (0 - 1.9, 2 -18.9, 19 - 

189, and $190 mg/l) 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

Table 5.7. Estimated range of Kd values for chromium (VI) as a function of soil pH, extractable iron content, and soluble 

sulfate. (Data analysis and generation of the table of Kd values are described in Appendix E.) 

pH 

4.1 - 5.0 5.1 - 6.0 6.1 - 7.0 $7.1 

DCB Extractable Fe 

(mmol/g) 


(mmol/g) 


(mmol/g) 


(mmol/g) 

#0.25 0.26 - 0.29 $0.30 #0.25 0.26 - 0.29 $0.30 ##0.25 0.26 - 0.29 $$0.30 ##0.25 0.26 - 0.29 $$0.30 

Kd (ml/g) 

Soluble 

Sulfate 

Conc 

(mg/l) 

Min 25 400 990 20 190 390 8 70 80 0 0 1 

0 - 1.9 

Max 35 700 1770 34 380 920 22 180 350 7 30 60 

5.24 

Min 12 190 460 10 90 180 4 30 40 0 0 1 

2 - 18.9 

Max 15 330 820 15 180 430 10 80 160 3 14 30 

Min 5 90 210 4 40 80 2 15 20 0 0 0 

19 - 189 

Max 8 150 380 7 80 200 5 40 75 2 7 13 

Min 3 40 100 2 20 40 1 7 8 0 0 0 

$190 

Max 4 70 180 3 40 90 2 20 35 1 3 6

5.5 Lead Geochemistry and K d Values 



Lead has 3 known oxidation states, 0, +2, and +4, and the most common redox state encountered 

in the environment is the divalent form. Total dissolved lead concentrations in natural waters are 

very low (~10-8 M). Dissolved lead in natural systems may exist in free ionic form and also as 

hydrolytic and complex species. Speciation calculations show that at pH values exceeding 7, 

aqueous lead exists mainly as carbonate complexes [PbCO 3 

5.25 

" (aq), and Pb(CO3) 2 

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 PbCO 3, Pb 3(OH) 2(CO 3) 2, PbSO 4, 

Pb 5(PO 4) 3(Cl), and Pb 4SO 4(CO 3) 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 (1-x)Pb xSO 4), apatite [Ca (1-x)Pb x(PO 4) 3OH], calcite (Ca (1-x)Pb xCO 3), and 

iron sulfides (Fe (1-x)Pb xS). 

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. 


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 µg/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 ( 204 Pb, 206 Pb, 207 Pb, and 208 Pb). The 

isotopes, 206 Pb, 207 Pb, and 208 Pb are the stable end products of the 238 U, 235 U, and 232 Th thorium 

decay series, respectively (Robbins, 1980). Additionally, heavier isotopes of lead ( 210 Pb, 211 Pb, 

212 Pb, and 214 Pb) are known to occur in nature as intermediate products of uranium and thorium 

decay (Robbins, 1978). The

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). 


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) 3 , is well established. Although several polymeric hydrolytic species such as Pb2OH3+ , 

3+ 4+ 4+ 

Pb3(OH) 3 , Pb4(OH) 4 , and Pb6(OH) 8 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 leadcarbonate 

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 µg/l based on the data for natural waters tabulated by Duram 

et al. (1971) and Hem (1985). A total of 21 aqueous species (uncomplexed Pb 2+ , 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 (

eactions. 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 Pb 2+ and its hydrolysis species, adsorb more 

commonly than anionic lead complexes. 


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 (PbCO 3), hydrocerussite [Pb 3(OH) 2(CO 3) 2], anglesite (PbSO 4), or 

chloropyromorphite [Pb 5(PO 4) 3Cl] may control aqueous lead concentrations. A review paper by 

McLean and Bledsoe (1992) included data which showed that lead concentrations in a calcareous 

soil was controlled by lead-phosphate compounds at lower pH and by mixed mineral phases at pH 

values exceeding 7.5. A study conducted by Mattigod et al. (1986) indicated that the mineral 

leadhillite [Pb 4SO 4(CO 3) 2(OH) 2] may be the solubility controlling solid for lead in a mine-waste 

contaminated soil. 

5.27

Table 5.8. Lead aqueous species included in the 

speciation calculations. 


Pb 2+ 

PbOH + " - 2- 

, Pb(OH) 2(aq), 

Pb(OH)3, 

Pb(OH)4 

3+ 

+ 

Pb2(OH) 3, 

Pb3(OH) 3 

" 2- + 

PbCO3(aq), Pb(CO3) 2 , PbHCO3 

" 2- 

PbSO4(aq), Pb(SO4) 2 

+ 

PbNO3 PbCl + " - 2- 

, PbCl2(aq), PbCl3, 

PbCl4 

PbF + " - 2- 

, PbF2(aq), PbF3, 

PbF4 

5.28


100 

80 

60 

40 

20 

0 

o 

PbSO4 (aq) 

Pb 2+ 

+ 

PbHCO3 

3 4 5 6 7 8 9 10 

5.29 

pH 

o 

PbCO3 (aq) 

o 

Pb(OH)2 (aq) 

2+ 

Pb(CO3)2 

PbOH + 

Figure 5.2. Calculated distribution of lead aqueous species as a function of pH for the 


concentration of 1 µg/l 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)Pb xSO 4], apatite [Ca (1-x)Pb x(PO 4) 3OH], calcite [Ca (1-x)Pb xCO 3], 

and iron sulfides, [Fe (1-x)Pb xS] (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 

PbSO 4, Pb 5(PO 4) 3OH, PbCO 3, and PbS, respectively.

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 

( 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 (

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. 



The review of lead K d 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: 

C Lead may precipitate in soils if soluble concentrations exceed about 4 mg/l at pH 4 and 

about 0.2 mg/l at pH 8. In the presence of phosphate and chloride, these solubility limits 

may be as low as 0.3 mg/l at pH 4 and 0.001 mg/l at pH 8. Therefore, in experiments in 

which concentrations of lead exceed these values, the calculated K d values may reflect 

precipitation reactions rather than adsorption reactions. 

C 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. 

C 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

C Adsorption of lead increases with increasing organic matter content of soils. 

C Increasing equilibrium solution concentrations correlates with decreasing lead adsorption 

(decrease in K d). 

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/l at pH 4 and about 0.2 mg/l at pH 8. In the presence of 

phosphate and chloride, these solubility limits may be as low as 0.3 mg/l at pH 4 and 0.001 mg/l at 

pH 8. Therefore, in experiments in which concentrations of lead exceed these values, the 

calculated K d 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 K d for lead by about 3 orders of magnitude. By comparison quantitative 

data is lacking on the effects of more common inorganic ligands (phosphate, chloride, and 

carbonate) on lead adsorption on soils. 

A number of adsorption studies indicate that within the pH range of soils (4 to 11), lead 

adsorption increases with increasing pH (Braids et al., 1972; Bittel and Miller, 1974; Griffin and 

Shimp, 1976; Haji-Djafari et al., 1981; Hildebrand and Blum, 1974; Overstreet and Krishamurthy, 

1950; Scrudato and Estes, 1975; Zimdahl and Hassett, 1977). Griffin and Shimp (1976) also 

noted that clay minerals adsorbing increasing amounts of lead with increasing pH may also be 

attributed to the formation of lead carbonate precipitates which was observed when the solution 

pH values exceeded 5 or 6. 

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

effect on lead adsorption have been established by Gerritse 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 K d values increased 

with decreasing lead concentrations in solution (from 0.2 mg/l to 0.0062 mg/l). 

5.5.6.2 K d Look-Up Tables 

Among all available data, Gerritse et al (1982) obtained adsorption data at lead concentrations 

(0.0001 - 0.01 mg/l) 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 K d 

values vary log-linearly as a function of equilibrium lead concentrations within the range of 

0.00001 to 0.2 mg/l. 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 K d as a function of soil pH and equilibrium lead 

concentrations. 


The pH ranges in the look-up table (Table 5.9) were selected from the rate of change that we 

noted in the K d data as a function of pH. The K d 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 K d values for lead as a function of soil pH, and 

equilibrium lead concentrations. 

Equilibrium Lead 

Concentration (µg/l) K d (ml/g) 

0.1 - 0.9 

1.0 - 9.9 

10 - 99.9 

100 - 200 

5.33 

Soil pH 

4.0 - 6.3 6.4 - 8.7 8.8 - 11.0 

Minimum 940 4,360 11,520 

Maximum 8,650 23,270 44,580 

Minimum 420 1,950 5,160 

Maximum 4,000 10,760 20,620 

Minimum 190 900 2,380 

Maximum 1,850 4,970 9,530 

Minimum 150 710 1,880 

Maximum 860 2,300 4,410

5.5.6.2.2 Limits of K d Values with Respect to Equilibrium Lead Concentrations 

The limits of equilibrium lead concentrations (0.0001 mg/l 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 K d values. 

5.6 Plutonium Geochemistry and K d Values 



In the ranges of pH and conditions typically encountered in the environment, plutonium can exist 

in all 4 oxidation states, namely +3, 4, +5, and +6. Under oxidizing conditions, Pu(IV), Pu(V), 

and Pu(VI) are common, whereas, under reducing conditions, Pu(III) 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 PuO 2@xH 2O) 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 K d values 

ranging from 11 to 300,000 ml/g. Plutonium in the higher oxidation state adsorbed on iron oxide 

surfaces may be reduced to the tetravalent state by Fe(II) present in the iron oxides. 

Two factors that influence the mobilization of adsorbed plutonium under environmental pH 

conditions (>7) are the concentrations of dissolved carbonate and hydroxyl ions. Both these 

ligands form very strong mixed ligand complexes with plutonium, resulting in desorption and 

increased mobility in the environment. 


Plutonium is produced by fissioning uranium fuel and is used in the construction of nuclear 

weapons. Plutonium has entered the environment either through accidental releases or through 

disposal of wastes generated during fuel processing and the production and detonation of nuclear 

weapons. Plutonium has 15 isotopes, but only 4 of these isotopes namely, 238 Pu [t ½ (half life) = 

5.34

86 y], 239 Pu (t ½ = 24,400 y), 240 Pu (t ½ = 6,580 y), 241 Pu (t ½ = 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 disproportionation 1 

reactions of plutonium are not significant (Cleveland, 1979). Under reducing conditions, Pu(III) 

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, 

1979; Rai et al., 1980b). 

Of the contaminated sites considered in EPA/DOE/NRC (1993), radioactive contamination by 

238 Pu, 239 Pu, and/or 240 Pu 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. 


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(III) > 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: 

3Pu 4+ + 2H2O = 2Pu 3+ 2+ + 

+ PuO2 +4H 

+ + 3+ 2+ 3PuO2 + 4H = Pu + 2PuO2 +2H2O. 

5.35

(Choppin, 1983). Plutonium hydrolytic species may have up to 4 coordinated hydroxyls. 

The tendency of plutonium in various oxidation states to form complexes depends on the ionic 

potential defined as the ratio (z/r) of the formal charge (z) to the ionic radius (r) of an ion. 

Among plutonium redox species, Pu(IV) exhibits the highest ionic potential and therefore forms 

" 

the strongest complexes with various ligands. Based on the equilibrium constants (K r ,298) for the 

plutonium complexation reactions, ligands, such as chloride and nitrate, form weak complexes 

" 

(log K r ,298 of 1 to 2) with plutonium, whereas fluoride, sulfate, phosphate, citrate, and oxalate 

" 

form stronger complexes (log K r ,298 of 6 to 30). Among the strongest complexes of plutonium 

2- 

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 of plutonium may 

be the dominant soluble species in natural environments at lower pH (below 5 to 6) values (Allard 

and Rydberg, 1983). 

Because dissolved plutonium can exist in multiple redox states and form hydrolytic and complex 

species in solution, it is useful to assess the probable dominant plutonium aqueous species that 

may exist in typical ground water. Therefore, the aqueous speciation of dissolved plutonium was 

calculated as a function of pH using the MINTEQA2 code and a concentration of 3.2x10 -10 mg/l 

(1.36x10 -15 M) total dissolved plutonium. This concentration is based on the maximum activity of 

239,240 Pu measured by Simpson et al. (1984) in 33 water samples taken from the highly alkaline 

Mono Lake in California. The species distribution was calculated assuming that multiple 

plutonium valence states might be present based on thermodynamic equilibrium considerations. 

This calculation is dependent on redox conditions as well as the pH and composition of the water. 

Therefore, a set of oxic conditions that might be associated with surface or near-surface disposal 

facilities or contaminated sites were selected for these illustrative calculations. These redox 

conditions are based on an experimentally determined pH/Eh 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 

where pe = negative log of the electron activity. 1 

The pe is related to Eh by the equation 

pe % pH ' 15.23, (5.1) 

Eh ' 2.303RT 

F 

where R = universal gas constant (1.9872 cal/mol·K) 

1 The electron activity is defined as unity for the standard hydrogen electrode. 

5.36 

pe (5.2)

T = temperature in degrees kelvin 

F = Faraday constant (96,487 coulombs/equivalent). 

At 25.0 " C (298 K), 

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. Thermodynamic data 

for these species were taken primarily from Lemire and Tremaine (1980) and other secondary 

sources and database modifications described by Krupka and Serne (1996). 

Results are plotted as a species distribution diagram (Figure 5.3). The data show that, under very 

2+ + 

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 

+ 2- 

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 

4- 

Pu(HPO4) 4 . At pH values exceeding 6.5, the bulk of the dissolved plutonium (~90 percent) 

2- " 

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 

2- 

complex species, Pu(OH) 2(CO3) 2 . 

Polymeric species of plutonium may not occur under environmental conditions because the total 

plutonium concentrations in nature are at least 7 orders of magnitude less than the concentrations 

required for the formation of such species (Choppin, 1983). It is important to note that the 

speciation of plutonium would change significantly with changing redox conditions, pH, the types 

and total concentrations of complexing ligands and major cationic constituents. 


Allard and Rydberg (1983) calculated that the aqueous concentrations of plutonium in nature may 

be controlled by the solubility of the solid phase PuO 2@xH 2O. Many observations show that 

plutonium associated with soils and particulate organic matter is present in tetravalent oxidation 

state (Nelson and Lovett, 1980; Nelson et al., 1987; Silver, 1983). Calculations by Allard and 

Rydberg (1983) based on available thermodynamic data show that, under reducing conditions, the 

solubility of dissolved plutonium would be limited by the solid phase PuO 2 at pH values greater 

than 8, and by the solid phase Pu 2(CO 3) 3 of trivalent plutonium at lower pH values. 

5.37

Table 5.10. Plutonium aqueous species included in the speciation calculations. 

Redox 

State 


Pu(III) Pu3+ , PuOH2+ + " 

, Pu(OH) 2, 

Pu(OH)3(aq) 

+ - 3- 

PuCO3, Pu(CO3) 2, 

Pu(CO3) 3 

+ - 

PuSO4, Pu(SO4) 2 

2+ 2+ 

PuH2PO4 , PuCl 

Pu(IV) Pu 4+ , PuOH 3+ 2+ - " 

, Pu(OH) 2 , Pu(OH)3, 

Pu(OH)4(aq) 

4- 2- 

Pu(OH) 4(CO3) 2 , Pu(OH)2(CO3) 2 

2+ 

PuSO4 , Pu(SO4) " 

2(aq), PuHPO4 2- 

Pu(HPO4) 3 , Pu(HPO4) 4 

5.38 

2+ , Pu(HPO4) " 

PuCl 3+ , PuF 3+ 2+ + " 

, PuF2 , PuF3, 

PuF4(aq) 

+ 

Pu(V) PuO2, PuO2OH " (aq), (PuO2) 2OH + 

2+, 

Pu(VI) PuO2 PuO2OH + " 

, PuO2(OH) 2(aq), 

- 2+ + 

PuO2(OH) 3, 

(PuO2) 2(OH) 2 , (PuO2) 3(OH) 5 

" 2- 4- 

PuO2CO3(aq), PuO2(CO3) 2 , PuO2(CO3) 3 

4- 

2(aq), 

PuO2Cl + , PuO2F + " - 2- 

, PuO2F2(aq), PuO2F3, PuO2F4 " + 

PuO2SO4(aq), PuO2H2PO4


100 

80 

60 

40 

20 

0 

+ 

PuO2 

2+ 

PuF2 

Pu 3+ 

3 4 5 6 7 8 9 10 

5.39 

pH 

2- 

Pu(OH)2(CO3)2 

4- 

Pu(HPO4)4 

o 

Pu(OH)4 (aq) 

Figure 5.3. Calculated distribution of plutonium aqueous species as a function of pH for the 


concentration of 3.2 x 10 -10 mg/l (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 PuO 2@xH 2O 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 PuO 2@xH 2O phase at concentrations about 2 orders of magnitude less than that of 

amorphous PuO 2@xH 2O. Therefore, under oxidizing conditions, amorphous PuO 2@xH 2O, 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, PuO 2@xH 2O may control plutonium 

concentrations at about 10 -10 M (Rai et al., 1980a).


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 (K d = 11 ml/g) to extremely high (K d = 300,000 ml/g) (Baes and Sharp, 1983; 

Coughtrey et al., 1985; Thibault et al., 1990). 

A number of studies indicate that iron hydroxides adsorb and reduce penta- and hexavalent 

plutonium to its tetravalent state at the solid surface. Experimental data showed that tetra- and 

pentavalent plutonium aqueous species oxidize to hexavalent form upon adsorption onto 

manganese dioxide surfaces whereas, pentavalent plutonium adsorbed on goethite 

disproportionate into tetra and hexavalent forms (Keeney-Kennicutt and Morse, 1985). 

Subsequently, the hexavalent form of plutonium was observed to have been reduced to tetravalent 

state. Additionally, these reactions were found to occur faster under light conditions than under 

dark conditions suggesting photochemical catalysis of adsorbed plutonium redox change 

reactions. 

Laboratory studies have indicated that increasing carbonate concentrations decreased adsorption 

of tetra- and pentavalent plutonium on goethite surfaces (Sanchez 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+240 Pu 

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



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 K d 

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: 

C Typically, in many experiments, the oxidation state of plutonium in solution was not 

determined or controlled. Therefore it would be inappropriate to compare the K d data 

obtained from different investigations. 

C 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 K d values would reflect mainly precipitation reactions and 

not adsorption reactions. 

C Adsorption data show that the presence of ligands influence plutonium adsorption onto 

soils. Increasing concentrations of ligands decrease plutonium adsorption. 

C If no complexing ligands are present plutonium adsorption increases with increasing pH 

(between 5.5 and 9.0). 

C Plutonium is known to adsorb onto soil components such as aluminum and iron oxides, 

hydroxides, oxyhydroxides, and clay minerals. However, the relationship between the 

amounts of these components in soils and the measured adsorption of plutonium has not 

been quantified. 

The factors which influence plutonium adsorption were identified from the following sources of 

data. A description and assessment of these data are provided in Appendix G. Because 

plutonium in nature can exist in multiple oxidation states (III, 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 


radionuclides, the studies by Duff et al. (1999) and Fisher et al. (1999) were identified and may be 

of interest to the reader. 

5.41

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 et al., 1975), are known to reduce adsorption of 

plutonium. Studies of suspended particles from natural water systems also showed that increasing 

concentrations of dissolved organic carbon decreased plutonium adsorption (Nelson et al., 1987). 

Experiments using synthetic ligands such as EDTA (1 mmol/l), DTPA (1 mmol/l), and HEDTA 

(100 mmol/l) 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/l 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/l HCO3 , Pu(IV) 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 6.8x10 -7 to 

1x10 -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 et al., 1975), attapulgite (Billon, 1982), and 

oxides, hydroxides, and oxyhydroxides (Evans, 1956; Charyulu et al., 1991; Sanchez et al., 1985; 

Tamura, 1972; Ticknor, 1993; Van Dalen et al., 1975) has been studied, correlative relationships 

between the type and quantities of soil minerals in soils and the overall plutonium adsorption 

behavior of the soils have not been established. 

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 -8 M initial 

concentration to study the adsorption behavior of Pu(IV) on 14 different soil samples from 

5.42

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 K d values. 

5.6.6.2 K d Look-Up Table 

The look-up table for plutonium K d 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 K d 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 K d 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 µm. 

Table 5.11. Estimated range of K d values for plutonium as a function of the 

soluble carbonate and soil clay content values. 

K d (ml/g) 

Clay Content (wt.%) 

0 - 30 31 - 50 51 - 70 

Soluble Carbonate 

(meq/l) 

5.43 


(meq/l) 


(meq/l) 

0.1 - 2 3 - 4 5 - 6 0.1 - 2 3 - 4 5 - 6 0.1 - 2 3 - 4 5 - 6 

Minimum 5 80 130 380 1,440 2,010 620 1,860 2,440 

Maximum 420 470 520 1,560 2,130 2,700 1,980 2,550 3,130

5.6.6.2.2 Limits of K d 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/l (0.1 to 6 mmol/l 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/l of dissolved carbonate (Richards, 1954). 

The regression relationship indicates that within the range of 0.1 to 6 mmol/l of dissolved 

carbonate, the K d 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/l) 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/l ) 

found soil saturation extracts, K d 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 

K d would be increasing pH. 

5.7 Radon Geochemistry and K d Values 



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” K d 

values for modeling contaminant transport and risks in soil /water systems. Radon is soluble in 

water, and the hydrostatic pressure on ground water below the water table is sufficient to keep 

dissolved radon in solution. 

The generation of radon is however affected by the concentrations of its parent elements which, 

along with radon’s decay products, are of regulatory concern. Because aqueous speciation, 

precipitation/dissolution, or adsorption/desorption processes can affect the movement of radon’s 

parents and decay products in soils, these processes should be considered when modeling 

contaminant transport in a total environmental system, including air transport pathways. 

5.44


Radon is a colorless, odorless, essentially inert gas. All radon isotopes are radioactive. The 

longest-lived isotope of radon is 222 Rn which has a half life (t ½) 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 220 Rn and/or 222 Rn 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: 

C 

C 

C 

222 226 238 

Rn by alpha decay from Ra in the U decay series 

220 

Rn (t½=54 sec) by alpha decay from 224 Ra in the 232 Th decay series 

219 

Rn (t½=3.9 sec) by alpha decay from 223 Ra in the 238 U decay series. 

The final, stable daughter products in these 3 decay series are 206 Pb, 208 Pb, and 207 Pb, 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 222 Rn in a water phase to that in a gas phase ranges from 0.52 at 0 " C to 

0.16 at 40 " C. This ratio has been used, for example, for the solubility of radon in water in 

mathematical models designed to calculate radon diffusion coefficients in soils (e.g., Nielson et 

al., 1984). The solubility of radon in organic liquids is greater than that in water. 


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 (t ½=3.8 d) for 222 Rn, 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


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 222 Rn and 220 Rn from solid 

aquifer material. They determined that radon and other decay products from the U-Th series 

were released by alpha recoil 1 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.” 


Because adsorption processes are not important relative to the movement of gaseous radon 

through aqueous environments, a review of K d values for radon was not conducted. 

Compilations, such as Thibault et al. (1990), do not list any K d values for radon. A K d value of 

zero should be considered for radon. 

5.8 Strontium Geochemistry and K d Values 



Strontium in solution is expected to be predominantly present as the uncomplexed Sr 2+ ion. Only 

in highly alkaline soils could strontianite (SrCO 3) 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 222 Rn from 226 Ra). 

5.46

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 CaCO 3 and/or precipitation as SrCO 3 may 

become an increasingly important mechanism controlling strontium removal from solution 

(Lefevre et al., 1993). A direct correlation between solution pH and strontium K d has been 

reported (Prout, 1958; Rhodes, 1957). This trend is likely the result of hydrogen ions competing 

with Sr 2+ for exchange sites and the result of pH increasing the CEC. Strontium K d 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 al., 1980). Calcium is an important competing cation affecting 90 Sr 

K d values (Kokotov and Popova, 1962; Schulz, 1965). The most important ancillary parameters 

affecting strontium K d values are CEC, pH, and concentrations of calcium and stable strontium. 


Strontium exists in nature only in the +2 oxidation state. The ionic radius of Sr 2+ is 1.12 Å, 

very close to that of Ca 2+ at 0.99 Å (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: 84 Sr (0.55 percent), 86 Sr (9.75 percent), 87 Sr (6.96 

percent), and 88 Sr (82.74 percent). The other radioisotopes of strontium are between 80 Sr and 

95 Sr. Only 90 Sr [half life (t½) = 28.1 y], a fission product, is of concern in waste disposal 

operations and environmental contamination. The radionuclide 89 Sr also is obtained in high yield, 

but the half-life is too short (t ½ = 52 d) to create a persistent environmental or disposal problem. 

Because of atmospheric testing of nuclear weapons, 90 Sr is distributed widely in nature. The 

average 90 Sr activity in soils in the United States is approximately 100 mCi/mi 2 . As a calcium 

analog, 90 Sr 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 90 Sr has been identified at 11 of the 45 Superfund National Priorities 

List (NPL). 


There is little tendency for strontium to form complexes with inorganic ligands (Faure and Powell, 

1972). The solubility of the free Sr 2+ 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 Sr 2+ , with no colloidal or 

anionic strontium present in the leachate. Stevenson and Fitch (1986) concluded that strontium 

should not form strong complexes with fulvic or humic acids based on the assumptions that 

strontium would exhibit similar stability with organic ligands as calcium and that strontium could 

not effectively compete with calcium for exchange sites because calcium would be present at 

5.47

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/l total dissolved strontium. Hem (1985, p. 135) lists 

this value as a median concentration of dissolved strontium for larger United States public water 

supplies based on analyses from Skougstad and Horr (1963). The strontium aqueous species 

included in the speciation calculations are listed in Table 5.12. These MINTEQA2 calculations 

support the contention that strontium will exist in groundwaters predominantly as the 

uncomplexed Sr2+ ion. The Sr2+ ion dominates the strontium speciation throughout the pH range 

of 3 to 10. Between pH 3 and 8.5, the Sr2+ species constitutes approximately 98 percent of the 

" 

total dissolved strontium. The remaining 2 percent is composed of the neutral species SrSO4(aq). " 

Between pH 9 and 10, SrCO3(aq) is calculated to be between 2 and 12 percent of the total 

" 

dissolved strontium. As the pH increases above 9, the SrCO3(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 µg/l. 


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 (SrSO 4) and strontianite (SrCO 3), have calcium 

counterparts, anhydrite (CaSO 4), and calcite (CaCO 3). In an acidic environment, most of the 

strontium solids will be highly soluble, and, if the activity of Sr 2+ in solution exceeds 

approximately 10 -4 mol/l, 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 



Sr 2+ , SrOH + 

" " + 

SrCO3(aq), SrSO4(aq), 

SrNO3 

SrCl + , SrF + 

- " + 2- 

SrPO3, SrHPO4(aq), 

SrH2PO4, SrP2O7 5.48

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),Sr x)SO 4 in more-alkaline environments (Ainsworth 

and Rai, 1987; Felmy et al., 1993). 


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 90 Sr, 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 exchange 1 (Ames and Rai, 1978; Lefevre et al., 1993; 

McHenry, 1958). 

Among the most important environmental parameters affecting the magnitude of a strontium K d 

value is the soil CEC (Ames and Rai, 1978; Lefevre et al., 1993; McHenry, 1958). This finding is 

consistent with cation exchange proposed as the mechanism generally controlling strontium 

adsorption. The results of Serne and LeGore (1996) also indicate that strontium adsorption is 

largely controlled by cation exchange. They reported that 90 Sr adsorption was reversible; that is, 

strontium could be easily desorbed (exchanged) from the surfaces of soils. Natural soils that had 

been in contact with 90 Sr for approximate 27 y could be leached of adsorbed 90 Sr as readily as 

similar soils containing recently adsorbed strontium, indicating that 90 Sr does not become more 

recalcitrant to leaching with time. Furthermore, these studies suggested that cation exchange, and 

not (co)precipitation, was responsible for 90 Sr sorption because the latter would leach at a much 

slower rate. 

Some studies indicate that a fraction of some 90 Sr 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 90 Sr into nonexchangeable forms by three 

soils. They observed that less than 10% of the total applied carrier-free 90 Sr 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) + 90 Sr 2% (aq) = 90 SrX(s) + Ca 2% (aq) 

where X designates an exchange surface site. Adsorption phenomena are discussed in more detail 

in Volume I of this report. 

5.49

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 90 Sr added to the soil 

could be displaced by repeated acidic ammonium acetate (pH 4.6) extractions. Wiklander 

proposed that the retention of 90 Sr 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 90 Sr 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 90 Sr 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 90 Sr present in these sediments was extracted by warm 1N NaCl or NH 4OAC 

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 1B 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 $ NH 4 > 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 90 Sr (Table 5.1). Consequently, mass action would improve the likelihood 

of calcium out competing 90 Sr 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 HCl and the ionic strength was adjusted by adding 4 M NaNO 3. For a dilute 

solution, the strontium K d 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 NaNO 3 

(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 K d 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 

(K d = 700 to 800 ml/g) than that reported by Rhodes (1957). Prout (1958) also reported only a 

5.50

slight pH effect on strontium K d values in high ionic strength solutions. Rhodes (1957) and Prout 

(1958) reported that increases in ionic strength resulted in lower strontium K d values. 



Two simplifying assumptions underlying the selection of strontium K d values included in the lookup 

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/l, ionic strength levels greater than about 0.1 M, and pH levels greater than about 12. 

Based on these assumptions and limitation, strontium K d 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 K d 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 (

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 

K d 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 K d 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 K d value. As the 

pH or the CEC of a system increases, so does the strontium K d 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 K d Values with Respect to Dissolved Calcium Concentrations 

Of the 63 experiments reporting strontium K d values, 32 also reported dissolved calcium 

concentrations (Appendix H). The mean calcium concentration in this data set was 56 mg/l, with 

a minimum of 0 mg/l and a maximum of 400 mg/l. Calcium concentration had a correlation with 

strontium K d 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

Table 5.13. Look-up table for estimated range of K d 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 (

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, K d, data indicates significant 

retention of thorium by most soil types. 


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: 

C 

C 

C 

238 234 

U decay series: Th [t½ (half life) = 24.1 d) and 230 Th (t½ = 8.0 x 10 4 y) 

232 232 

Th decay series: Th (t½ = 1.41 x 10 10 y) and 228 Th (t½ = 1.913 y) 

235 231 

U decay series: Th (t½ = 25.5 h) and 227 Th (t½ = 18.5 d). 

Natural thorium consists of essentially 1 isotope, 232 Th, with trace quantities of the other isotopes. 

Thorium is fertile nuclear material in that the principal isotope 232 Th can be converted by capture 

of a thermal neutron and 2 beta decays to fissionable 233 U which does not exist in nature. The 

application of thorium as a reactor fuel in the ThO 2 ceramic form is described in detail by Belle 

and Berman (1984). 

Thorium occurs only in the +4 oxidation state in nature. The Th 4+ ion is the largest tetravalent 

cation known with a radius of approximately 1.0 Å. Although the Th 4+ 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 µg/l (0.1 pCi/l 232 Th), 

5.54

although mg/l concentrations of 232 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 thoriumcontaining 

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,etc.)SiO 4] and thorianite (crystalline ThO 2). 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)PO 4] and 

zircon (ZrSiO 4). 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 al., 1967; Frondel, 1958). 

Concentrations of thorium can be several weight percent in these deposits. 

Because of their long half lives, 228 Th (t ½ = 1.913 y), 230 Th (t ½ = 8.0 x 10 4 y), and 232 Th (t ½ = 

1.41 x 10 10 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 228 Th, 230 Th, 

and/or 232 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. 


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

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 µg/l 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 µg/l 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 

2+ + " 

, Th(OH)3, 

and Th(OH)4(aq). 

The latter 2 hydrolytic complexes have the widest range 

Th(OH) 2 

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 

6+ 8+ 9+ 

for the thorium polynuclear hydrolyzed species Th2(OH) 2 , Th4(OH) 8 , and Th6(OH) 15 

based on 

the review of Bases and Mesmer (1976). 

Table 5.14. Thorium aqueous species included in the 



Th 4+ , ThOH 3+ 2+ + 

, Th(OH) 2 , Th(OH)3, 

Th(OH)4°(aq), 

6+ 8+ 9+ 

Th2(OH) 2 , Th4(OH) 8 , Th6(OH) 15 

- 6- 

Th(OH) 3CO3 and Th(CO3) 5 

ThF 3+ 2+ + 

, ThF2 , ThF3, 

ThF4°(aq) 

ThCl 3+ 2+ + 

, ThCl2 , ThCl3, 

ThCl4°(aq) 

2+ 2- 4- 

ThSO4 , Th(SO4) 2°(aq), Th(SO4) 3 , Th(SO4) 4 

4+ 3+ 2+ 

ThH3PO4 , ThH2PO4 , Th(H2PO4) 2 , 

2- 

Th(HPO4) 2°(aq), Th(HPO4) 3 

5.56

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). Östhols et al. (1994) have recently published 

- and Th(CO3) 5 

thermodynamic constants for the thorium carbonate complexes Th(OH) 3CO 3 


100 

80 

60 

40 

20 

0 

Th 4+ 

ThOH + 

2+ 

Th(OH) 2 

+ 

Th(OH) 3 

3 4 5 6 7 8 9 10 

5.57 

pH 

o 

Th(OH) 4 (aq) 

6- that 

are based on their solubility studies of microcrystalline ThO 2 at different partial pressures of CO 2 

in aqueous media. 

Figure 5.4. Calculated distribution of thorium hydrolytic species as a function of pH. 

[The species distribution is based on a concentration of 1 µg/l total 

dissolved thorium in pure water (i.e., absence of complexing ligands other 

than OH - ) and thermodynamic data from Langmuir and Herman (1980).]

The distribution of thorium aqueous species (Figure 5.5) was also calculated as a function of pH 

using the MINTEQA2 for a concentration of 1 µg/l total dissolved thorium and the water 

composition in Table 5.1. The thermodynamic data were principally from Langmuir and Herman 

- 6- 

(1980). The thermodynamic constants for the aqueous species Th(OH) 3CO3 and Th(CO3) 5 from 

Östhols 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 µg/l. At 

- 

the higher concentration of dissolved thorium, the stability of Th(OH) 3CO3 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 

3- 2- 

soil/water systems. Langmuir and Herman (1980) used citrate (C6H5O7 ), oxalate (C2O4 ), and 

4- 

ethylenediamine tetra-acetic acid (EDTA) (C10H12O8N2 ) to show the possible role of organic 

complexes in the mobility of thorium in natural waters. Based on the stability constants available 

for thorium citrate, oxalate, and ethylenediamine complexes, calculations by Langmuir and 

Herman (1980) indicate that thorium organic complexes likely predominate over inorganic 

complexes in organic-rich waters and soils. For the concentrations considered by Langmuir and 

Herman (1980), the ThEDTA " (aq) complex dominates all other thorium aqueous species over the 

pH range from 2 to 8. This would in turn have an important effect on the solubility and 

adsorption of thorium in such waters. 


The main thorium-containing minerals, thorite [(Th,U,Ce,Fe,etc.)SiO 4], thorianite (crystalline 

ThO 2), monazite [(Ce,La,Th)PO 4) and zircon (ZrSiO 4), 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 lowtemperature 

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


100 

80 

60 

40 

20 

0 

+ 

ThF3 2+ 

ThF2 o 

ThF4 (aq) 

ThF 3+ 

2- 

Th(HPO4) 3 

o 

ThHPO4 (aq) 

3 4 5 6 7 8 9 10 

5.59 

pH 

- 

Th(OH) 3CO3 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 µg/l total dissolved thorium and thermodynamic data 

from Langmuir and Herman (1980) and Östhols et al. (1994, for 

- 6- 

Th(OH) 3CO3 and Th(CO3) 5 ). The thermodynamic database used for these 

speciation calculations did not include the constants for thorium humic acid 

complexes.] 

The maximum concentration of dissolved thorium that may occur in a low-temperature aqueous 

system can be predicted with the solubility of hydrous thorium oxide. This solid is known to 

precipitate in laboratory experiments conducted at low temperature, oversaturated conditions 

over several weeks. If this solid precipitates in a natural environment, it will likely alter with time 

to a more crystalline solid that has a lower solubility. The solubility of hydrous thorium oxide has 

been studied experimentally by Rai and coworkers (Felmy et al., 1991; Rai et al., 1995; Ryan and 

Rai, 1987). In 0.1 M NaClO 4 solutions, the measured solubility of hydrous thorium oxide ranges

from about 10 -8.5 mol/l (0.0007 mg/l) to less than 10 -9 mol/l (0.0002 mg/l) in the pH range from 5 

to 10 (Ryan and Rai, 1987). The concentration of dissolved thorium increases to approximately 

10 -2.6 mol/l (600 mg/l) 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 NaClO 4 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 NaClO 4 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. 


Thorium concentrations in surface- and groundwaters may also be controlled to very low levels 

(# few µg/l) 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 K d data by soil type needed to model radionuclide migration 

from a nuclear waste geological disposal vault to the biosphere. Thibault et al. list K d values for 

thorium that range from 207 to 13,000,000 ml/g. The range of thorium K d values listed for 

organic soil was 1,579 to 1.3 x 10 7 ml/g. Based on our experience, the very high K d 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 K d 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 1B, precipitation of solids containing the contaminant of interest results in K d 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 Östhols (1995) studied the adsorption of thorium 

on amorphous colloidal particles of silica (SiO 2). 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 ("-FeOOH) and nsutite 

((-MnO 2) 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

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 

(-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 adsorption goes from 0 percent to 

approximately 90-100 percent) was measured to be in the pH range from 2 to 5. For conditions 

considered in their study, ionic strength was found to have no effect on the adsorption of thorium 

on goethite. LaFlamme and Murray did however observe a strong influence of carbonate 

alkalinity on thorium adsorption. In their experiments at pH 9.0±0.6, they observed a decrease of 

thorium adsorption with the addition of 100 meq/l carbonate alkalinity, and no measurable 

adsorption of thorium at carbonate alkalinity greater than 300 meq/l. At the low particle 

concentrations used in their experiments, LaFlamme and Murray attributed this reduction to the 

2- - 

competition for surface sites by CO3 and HCO3 and the formation of soluble thorium-carbonate 

complexes with a net negative charge. 



Two generalized, simplifying assumptions were established for the selection of thorium K d 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: 

C 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 K d value will be used in the look-up table to 

approximate thorium behavior in systems with high thorium concentrations. 

C Thorium adsorption occurs at concentrations less than 10 -9 M. The extent of thorium 

adsorption can be estimated by soil pH. 


However, these simplifying assumptions are clearly compromised in systems containing high 

alkaline (LaFlamme and Murray, 1987), carbonate (LaFlamme 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 in 

more detail in the following sections. 

5.61

Based on the assumptions and limitations described above, thorium K d 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 K d values (not percent 


C Low ionic strength (< 0.1 M) 

C pH values between 4 and 10.5 

C Dissolved thorium concentrations less than 10 -9 M 

C Low humic material concentrations (

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 K d values in Table 5.15. 

5.9.6.2.1 Limits of K d Values with Respect to Organic Matter and Aluminum/Iron-Oxide 

Concentrations 

Of the 17 entries in the thorium K d data set (Appendix I), none of them had accompanying 

organic matter or aluminum- and iron-oxide mineral concentration data. It was anticipated that 

the presence of organic matter would decrease thorium K d 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 aluminumand/or 

iron-oxides would increase thorium K d values by increasing the number of adsorption 

(surface complexation) sites. 

5.9.6.2.2 Limits of K d Values with Respect to Dissolved Carbonate Concentrations 

Of the 17 entries in the thorium K d data set (Appendix I), none of them had accompanying 

carbonate concentration data. However, 5 entries had calcite (CaCO 3) 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 K d 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 K d values due to formation of the weaker forming 

carbonate-thorium complexes. 

Table 5.15. Look-up table for thorium K d 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 (

5.10 Tritium Geochemistry And K d Values 



Tritium, a radioactive isotope of hydrogen with a half life (t ½) 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. 


Tritium ( 3 H) is a radioactive isotope of hydrogen. Three isotopes of hydrogen are known. These 

include the 2 stable isotopes 1 H (protium or H) and 2 H (deuterium or D), and the radioactive 

isotope 3 H (tritium or T). Tritium has a half life (t ½) of 12.3 y, and disintegrates into helium-3 

( 3 He) by emission of a weak beta ($ - ) 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 1 H, 2 H, and 3 H 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 6 Li. 

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

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. 


Because tritium oxidizes rapidly to form isotopic water, aqueous speciation reactions do not 

affect the mobility of tritium in soil/water systems. 


Neither precipitation or coprecipitation processes affect the mobility of tritium in soil/water 

systems. 


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, K d, 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 K d 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. 


A review of the literature pertaining to K d values for tritium was not conducted given the limited 

availability of K d values for tritium (see section above) and limited importance of sorption 

processes relative to the mobility of tritium in aqueous environments. 

5.11 Uranium Geochemistry and K d Values 



In essentially all geologic environments, +4 and +6 are the most important oxidation states of 

5.65

uranium. Uranium(VI) species dominate in oxidizing environments. Uranium(VI) retention by 

soils and rocks in alkaline conditions is poor because of the predominance of neutral or negatively 

charged species. An increase in CO 2 pressure in soil solutions reduces U(VI) adsorption by 

promoting the formation of poorly sorbing carbonate complexes. Uranium(IV) species dominate 

in reducing environments. Uranium(IV) tends to hydrolyze and form strong hydrolytic 

complexes. Uranium(IV) also tends to form sparingly soluble precipitates that commonly control 

U(IV) 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(IV) solubility. There are several 

ancillary environmental parameters affecting uranium migration. The most important of these 

parameters include redox status, pH, ligand (carbonate, fluoride, sulfate, phosphate, and dissolved 

carbon) concentrations, aluminum- and iron-oxide mineral concentrations, and uranium 



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 

238 U, 0.711 percent 235 U, and 0.0054 percent 234 U by weight. The half-lives of these isotopes are 

4.51 x 10 9 y, 7.1 x 10 8 y, and 2.47 x 10 5 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(IV) minerals include uraninite (UO 2 through UO 2.25) and coffinite [USiO 4] (Frondel, 1958; 

Langmuir, 1978). Aqueous U(IV) 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 [(K 2(UO 2) 2(VO 4) 2], schoepite 

(UO 3·2H 2O), rutherfordine (UO 2CO 3), tyuyamunite [Ca(UO 2) 2(VO 4) 2], autunite 

[Ca(UO 2) 2(PO 4) 2], potassium autunite [K 2(UO 2) 2(PO 4) 2], and uranophane [Ca(UO 2) 2(SiO 3OH) 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 lignite 1 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 234 U, 235 U, and/or 238 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

National Priorities List (NPL) sites and 26 of the 38 NRC Site Decommissioning Site Plan 

(SDMP) sites. 


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(III) easily oxidizes to U(IV) under most reducing conditions 

+ 1 

) readily disproportionates to U(IV) and U(VI). 

found in nature. The U(V) aqueous species (UO 2 

Consequently, U(IV) and U(VI) are the most common oxidation states of uranium in nature. 

Uranium will exist in the +6 and +4 oxidation states, respectively, in oxidizing and more reducing 

environments. 

2+ 4+ 4+ 

Both uranium species, UO2 and U , hydrolyze readily. The U ion is more readily hydrolyzed 

2+ 

than UO2 , 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/l), 

2+ 

fluoride (0.2 mg/l), phosphate (0.1 mg/l), and sulfate (100 mg/l). Below pH 3, UF2 was the 

dominant uranium species. The speciation of dissolved U(IV) at pH values greater than 3 is 

+ N 

dominated by hydrolytic species such as U(OH) 3 and U(OH)4(aq). 

Complexes with chloride, 

fluoride, phosphate, and sulfate were not important above pH 3. The total U(IV) concentration in 

solution is generally quite low, between 3 and 30 µg/l, because of the low solubility of U(IV) solid 

phases (Bruno et al., 1988; Bruno et al., 1991). Precipitation is discussed further in the next 

section. 

Dissolved U(VI) hydrolyses to form a number of aqueous complexes. The distribution of U(VI) 

species is presented in Figures 5.6a-b and 5.7. The distribution of uranyl hydrolytic species 

(Figures 5.6a-b) was calculated as a function of pH using the MINTEQA2 code. The U(VI) 

aqueous species included in the speciation calculations are listed in Table 5.16. The 

thermodynamic data for these aqueous species were taken primarily from Wanner and Forest 

(1992). Because dissolved uranyl ions can be present as polynuclear 2 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 µg/l, 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 + 2+ 4+ 

= UO2 + U . 

2 - 

A polynuclear species contains more than 1 central cation moiety, e.g., (UO2) 2CO3(OH) 3 and 

4+ Pb4(OH) 4 . 

5.67

gives 0.1 to 10 µg/l 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 µg/l. 

2+ 

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 µg/l (Figure 5.6a) and 1,000 µg/l dissolved 

+ 

U(VI) (Figure 5.6b). At 1,000 µg/l dissolved uranium, some polynuclear species, (UO2) 3(OH) 5 

2+ 

and (UO2) 2(OH) 2 , 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 µg/l dissolved U(VI). 

At pH values less than 5, the UO2F + species dominates the system, whereas at pH values greater 

" 2- 4- 

than 5, carbonate complexes [UO2CO3(aq), UO2(CO3) 2 , UO2(CO3) 3 ] dominate the system. 

These calculations clearly show the importance of carbonate chemistry on U(VI) speciation. For 

this water composition, complexes with chloride, sulfate, and phosphate were relatively less 

important. Consistent with the results in Figure 5.7, Langmuir (1978) concluded that the uranyl 

complexes with chloride, phosphate, and sulfate were not important in a typical groundwater. 

The species distribution illustrated in Figure 5.7 changes slightly at pH values greater than 6 if the 

concentration of total dissolved uranium is decreased from 1,000 to 1 µg/l. At the lower 

- 

concentration of dissolved uranium, the species (UO2) 2CO3(OH) 3 is no longer present as a 

dominant aqueous species. 

2+ " - 

Sandino and Bruno (1992) showed that UO2 -phosphate complexes [UO2HPO4(aq) and UO2PO4] 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 

2+ 

be about +3.3 for U(VI) in UO2 . Kim (1986) concluded that, in general, +6 actinides, including 

U(VI), would have approximately the same tendency to form humic- or fulvic-acid complexes as 

to hydrolyze or form carbonate complexes. This suggests that the dominant reaction with the 

uranyl ion that will take place in a groundwater will depend largely on the relative concentrations 

of hydroxide, carbonate, and organic material concentrations. He also concluded, based on 

comparison of stability constants, that the tendency for U 4+ to form humic- or fulvic-acid 

complexes is less than its tendency to hydrolyze or form carbonate complexes. Importantly, 

5.68

U(IV) and U(VI) can form stable organic complexes, thereby increasing their solubility and 

mobility. 

Table 5.16. Uranium(VI) aqueous species included in the 



2+ 

UO2 , UO2OH + N - 2- 

, UO2(OH) 2(aq), 

UO2(OH) 3, 

, UO2(OH) 4 , 

(UO2) 2OH3+ 2+ 2+ + 

, (UO2) 2(OH) 2 , (UO2) 3(OH) 4 , (UO2) 3(OH) 5, 

+ 9+ 

, U6(OH) 15 

- 

(UO2) 3(OH) 7, 

(UO2) 4(OH) 7 

N 2- 4- 5- 

UO2CO3(aq), UO2(CO3) 2 , UO2(CO3) 3 , UO2(CO3) 3 , 

6- 2 

, (UO2) 11(CO3) 6(OH) 12 

(UO 2) 3(CO 3) 6 

5.69 

- - 

, (UO2) 2CO3(OH) 3 

- N + 2+ 

UO2PO4, UO2HPO4(aq), UO2H2PO4, UO2H3PO4 , 

N 

UO2(H2PO4) 2(aq), 

UO2(H2PO4)(H3PO4) + , 

N 2- 

UO2SO4(aq), UO2(SO4) 2 

+ 

UO2NO3 UO2Cl + N 

, UO2Cl2(aq), UO2F + N - 2- 

, UO2F2(aq), UO2F3, UO2F4 + 

UO2SiO(OH) 3 


Dissolution, precipitation, and coprecipitation have a much greater effect on the concentrations of 

U(IV) than on the concentration of U(VI) in groundwaters. In most cases, these processes will 

likely not control the concentration of U(VI) in oxygenated groundwaters far from a uranium 

source. Near a uranium source, or in reduced environments, these processes tend to become 

increasingly important and several (co)precipitates may form depending on the environmental 

conditions (Falck, 1991; Frondel, 1958). Reducing conditions may exist in deep aquifers, marsh 

areas, or engineered barriers that may cause U(IV) to precipitate. Important U(IV) minerals 

include uraninite (compositions ranging from UO 2 to UO 2.25), coffinite (USiO 4), and ningyoite 

[CaU(PO 4) 2·2H 2O] (Frondel, 1958; Langmuir, 1978). Important U(VI) minerals include carnotite 

[(K 2(UO 2) 2(VO 4) 2], schoepite (UO 3·2H 2O), rutherfordine (UO 2CO 3), tyuyamunite 

[Ca(UO 2) 2(VO 4) 2], autunite [Ca(UO 2) 2(PO 4) 2], potassium autunite [K 2(UO 2) 2(PO 4) 2], and 

uranophane [Ca(UO 2) 2(SiO 3OH) 2] (Frondel, 1958; Langmuir, 1978). Carnotite, a U(VI) mineral, 

is found in the oxidized zones of uranium ore deposits and uraninite, a U(IV) mineral, is a primary

mineral in reducing ore zones (Frondel, 1958). The best way to model the concentration of 

precipitated uranium is not with the K d construct, but through the use of solubility constants. 


100 

80 

60 

40 

20 

0 

2+ 

UO2 

UO2OH + 

3 4 5 6 7 8 9 10 

5.70 

o 

UO2(OH)2 (aq) 

pH 

- 

UO2(OH)3 

Figure 5.6a. Calculated distribution of U(VI) hydrolytic species as a function of pH 

at 0.1 µg/l total dissolved U(VI). [The species distribution is based on 

U(VI) dissolved in pure water (i.e., absence of complexing ligands other 

than OH - ) and thermodynamic data from Wanner and Forest (1992).]


100 

80 

60 

40 

20 

0 

2+ 

UO2 

2+ 

(UO2)2(OH)2 

UO2OH + 

+ 

(UO2)3(OH)5 

3 4 5 6 7 8 9 10 

5.71 

pH 

o 

UO2(OH)2 (aq) 

- 

UO2(OH)3 

Figure 5.6b. Calculated distribution of U(VI) hydrolytic species as a function of pH at 

1,000 µg/l total dissolved U(VI). [The species distribution is based on 

U(VI) dissolved in pure water and thermodynamic data from Wanner and 

Forest (1992).]


100 

80 

60 

40 

20 

0 

UO2F + 

2+ 

UO2 

Other Species 

o 

UO2HPO4 (aq) 

o 

UO2CO3 (aq) 

3 4 5 6 7 8 9 10 

5.72 

- 

(UO2)2CO3(OH)3 

o 

UO2(OH)2 (aq) 

pH 

4- 

UO2(CO3)3 

2- 

UO2(CO3) 2 

- 

UO2(OH)3 

Figure 5.7. Calculated distribution of U(VI) aqueous species as a function of pH for the 


concentration of 1,000 µg/l total dissolved U(VI) and thermodynamic data 

from Wanner and Forest (1992).] 


In low ionic strength solutions with low U(VI) concentrations, dissolved uranyl concentrations 

will likely be controlled by cation exchange and adsorption processes. The uranyl ion and its 

complexes adsorb onto clays (Ames 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 Ca 2+ , Mg 2+ , and K + , will displace the uranyl ion from soil exchange sites, forcing 

it into solution. For this reason, the uranyl ion is particularly mobile in high ionic-strength

solutions. Not only will other cations dominate over the uranyl ion in competition for exchange 

sites, but carbonate ions will form strong soluble complexes with the uranyl ion, further lowering 

the activity of this ion while increasing the total amount of uranium in solution (Yeh and Tripathi, 

1991). 

Some of the sorption processes to which uranyl ion is subjected are not completely reversible. 

Sorption onto iron and manganese oxides can be a major process for extraction of uranium from 

solution (Hsi and Langmuir, 1985; Waite 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(VI) in soils and sediments. The 

mechanisms by which uranium is sequestered by organic matter have not been worked out in 

detail. One possible process involves adsorption of uranium to humic substances through rapid 

ion-exchange and complexation processes with carboxylic and other acidic functional groups 

(Boggs et al., 1985; Borovec et al., 1979; Idiz et al., 1986; Shanbhag and Choppin, 1981; Szalay, 

1964). These groups can coordinate with the uranyl ion, displacing waters of hydration, to form 

stable complexes. A process such as this probably accounts for a significant fraction of the 

organically bound uranium in surface and subsurface soils. Alternatively, sedimentary organics 

may act to reduce dissolved U(VI) species to U(IV) (Nash et al., 1981). 

Uranium sorption to iron oxide minerals and smectite clay has been shown to be extensive in the 

absence of dissolved carbonate (Ames et al., 1982; Hsi and Langmuir, 1985; Kent et al., 1988). 

However, in the presence of carbonate and organic complexants, sorption has been shown to be 

substantially reduced or severely inhibited (Hsi and Langmuir, 1985; Kent 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


5.11.6.1 General Availability of K d Values 

More than 20 references (Appendix J) that reported K d values for the sorption of uranium onto 

soils, crushed rock material, and single mineral phases were identified during this review. 1 These 

studies were typically conducted to support uranium migration investigations and safety 

assessments associated with the genesis of uranium ore deposits, remediation of uranium mill 

tailings, agriculture practices, and the near-surface and deep geologic disposal of low-level and 

high-level radioactive wastes (including spent nuclear fuel). These studies indicated that pH and 

dissolved carbonate concentrations are the 2 most important factors influencing the adsorption 

behavior of U(VI). 

The uranium K d 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 K d values that are less than 10 ml/g. At pH values less 

than 3.5 and greater than 8, this variability includes K d values of less than 1 ml/g. 

Uranium K d 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 K d 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 et al., 1992, 1994), clays (McKinley et 

al., 1995; Turner et al., 1996; Waite et al., 1992), and quartz (Waite et al., 1992). This pHdependent 

behavior is related to the pH-dependent surface charge properties of the soil minerals 

and complex aqueous speciation of dissolved U(VI), especially near and above neutral pH 

conditions where dissolved U(VI) forms strong anionic uranyl-carbonato complexes with 

dissolved carbonate. 

5.11.6.2 Look-Up Table 

Solution pH was used as the basis for generating a look-up table for the range of estimated 

minimum and maximum K d values for uranium. Given the orders of magnitude variability 

observed for reported uranium K d values, a subjective approach was used to estimate the 

1 


radionuclides, the studies by Pabalan et al. (1998), Payne et al. (1998), Redden et al. (1998), 

Rosentreter et al. (1998), and Thompson et al. (1998) were identified and may be of interest to 

the reader. 

5.74

minimum and maximum K d values for uranium as a function of pH. These values are listed in 

Table 5.17. For K d 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-K d values listed in Table 5.17. 

Table 5.17. Look-up table for estimated range of K d values for uranium based on pH. 

K d 

(ml/g) 

5.75 

pH 

3 4 5 6 7 8 9 10 

Minimum

No attempt was made to statistically fit the K d 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 CO 2 or carbonate phases present in 

the soil samples. Given the complexity of these reaction processes, it is recommended that the 

reader consider the application of geochemical reaction codes, and surface complexation models 

in particular, as the best approach to predicting the role of dissolved carbonate in the adsorption 

behavior of uranium and derivation of U(VI) K d values when site-specific K d values are not 

available. 

5.11.6.2.2 Limits of K d Values with Respect to Clay Content and CEC 

No attempt was made to statistically fit the K d 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 K d values and the clay content or CEC of their soils. Their systems covered the pH 

conditions from 3 to 7. 

However, clays have an important role in the adsorption of uranium in soils. Attempts have been 

made (e.g., Borovec, 1981) to represent this functionality with a mathematical expression, but 

such studies are typically for limited geochemical conditions. Based on studies by 

Chisholm-Brause (1994), Morris et al. (1994), McKinley et al. (1995), Turner et al. (1996), and 

others, uranium adsorption onto clay minerals is complicated and involves multiple binding sites, 

including exchange and edge-coordination sites. The reader is referred to these references for a 

detailed treatment of the uranium adsorption on smectite clays and application of surface 

complexation modeling techniques for such minerals. 

5.11.6.2.3 Use of Surface Complexation Models to Predict Uranium K d 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 K d 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 d values for the geochemical conditions of interest, the reader is encouraged to 

5.76

apply this technology to predict bounding uranium K d 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(VI). 

Important solution speciation, (co)precipitation/dissolution, and adsorption reactions were 

discussed for each contaminant. The species distributions for each contaminant were calculated 

using the chemical equilibria code MINTEQA2 (Version 3.11, Allison 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 0 , Sr 2+ , and HTO, 

respectively (Ames and Rai, 1978; Rai and Zachara, 1984). Chromium (as chromate, CrO 4 

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(III), Pu(IV), Pu(V), and Pu(VI)] and it forms stable 

complexes with a large number of ligands. Because of redox sensitivity, the speciation of uranium 

exhibits a large number of stable complexes. Uranium(VI) also forms polynuclear complex 

species [complexes containing more than 1 mole of uranyl [e.g., (UO 2) 2CO 3OH - ]. 

One general conclusion that can be made from the results in Table 5.19 is that, as the pH 

increases, the aqueous complexes tend to become increasingly more negatively charged. For 

example, lead, plutonium, thorium, and uranium are cationic at pH 3. At pH values greater 

than 7, they exist predominantly as either neutral or anionic species. Negatively charged 

complexes tend to adsorb less to soils than their respective cationic species. This rule-of-thumb 

stems from the fact that most minerals in soils have a net negative charge. Conversely, the 

solubility of several of these contaminants decreases dramatically as pH increases. Therefore, the 

net contaminant concentration in solution does not necessarily increase as the dominant aqueous 

species becomes more negatively charged. 

5.77 

2- ),

Table 5.18. Selected chemical and transport properties of the contaminants. 

Element Radionuclide 

1 

Primary Species at pH 7 

and Oxidizing Conditions Redox 

5.78 

Sensitive 2 

Cationic Anionic Neutral Not 

Retarded 3 

Transport Through 

Soils at pH 7 

Retarded 3 

Cd x x x 

Cs x x x 

Cr x x x x 

Pb x x x x 

Pu x x x x x 

Rn x x x 

Sr x x x 

Th x x x 

3 H x x x 

U x x x x x 

1 Contaminants that are primarily a health concern as a result of their radioactivity are identified in 

this column. Some of these contaminants also exist as stable isotopes (e.g., cesium and strontium). 

2 The redox status column identifies contaminants (Cr, Pu, and U) that have variable oxidation 

states within the pH and Eh limits commonly found in the environment and contaminants (Cd and 

Pb) whose transport is affected by aqueous complexes or precipitates involving other redoxsensitive 

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.

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 

pH 3 pH 7 pH 10 

Species % Species % Species % 

Cd Cd 2+ 97 Cd 2+ 

CdHCO 3 + 

CdCO 3 " (aq) 

5.79 

84 

6 

6 

CdCO 3 " (aq) 96 

Cs Cs + 100 Cs + 100 Cs + 100 

Cr HCrO 4 - 99 CrO 4 2- 

Pb Pb 2+ 

PbSO 4 " (aq) 

Pu PuF 2 2+ 

PuO 2 + 

Pu 3+ 

96 

4 

69 

24 

5 

HCrO 4 - 

PbCO 3 " (aq) 

Pb 2+ 

PbHCO 3 + 

PbOH + 

Pu(OH) 2(CO 3) 2 2- 

Pu(OH) 4 " (aq) 

78 

22 

75 

15 

7 

3 

94 

5 

CrO 4 2- 99 

PbCO 3 " (aq) 

Pb(CO 3) 2 2- 

Pb(OH) 2 " (aq) 

Pb(OH) + 

Pu(OH) 2(CO 3) 2 2- 

Pu(OH) 4 " (aq) 

Rn Rn 0 100 Rn 0 100 Rn 0 100 

Sr Sr 2+ 99 Sr 2+ 99 Sr 2+ 

Th ThF 2 2+ 

ThF 3 + 

54 

42 

Th(HPO 4) 3 2- 

Th(OH) 3CO 3 - 

76 

22 

SrCO 3 " (aq) 

50 

38 

9 

3 

90 

10 

86 

12 

Th(OH) 3CO 3 - 99 

3 H HTO 100 HTO 100 HTO 100 

U 

0.1 µg/l 

U 

1,000 µg/l 

UO 2F + 

UO 2 2+ 

UO 2F 2 " (aq) 

UO 2F + 

UO 2 2+ 

UO 2F 2 " (aq) 

62 

31 

4 

61 

33 

4 

UO 2(CO 3) 2 2- 

UO 2(OH) 2 " (aq) 

UO 2CO 3 " (aq) 

UO 2PO 4 - 

UO 2(CO 3) 2 2- 

(UO 2) 2CO 3(OH) 3 - 

UO 2(OH) 2 " (aq) 

UO 2CO 3 " (aq) 

58 

19 

17 

3 

41 

30 

13 

12 

UO 2(CO 3) 3 4- 

UO 2(OH) 3 - 

UO 2(CO 3) 2 2- 

UO 2(CO 3) 3 4- 

UO 2(OH) 3 - 

UO 2(CO 3) 2 2- 

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. 

63 

31 

4 

62 

32 

4

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 K d 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 variablecharged 

minerals (aluminum- and iron-oxide minerals), partitioning of contaminants to organic 

matter, CEC, formation of polynuclear complexes, oxidation state of contaminants and 

complexing/precipitating ligands, and H + -competition for adsorption sites. 

The redox status of a system also influences the sorption of several contaminants included in this 

study (Table 5.20). Like pH, redox has direct and indirect effects on contaminant 

(co)precipitation. The direct effect occurs with contaminants like uranium and chromium where 

the oxidized species form more soluble solid phases than the reduced species. Redox conditions 

also have a direct effect on the sorption of plutonium, but the effects are quite complicated. The 

indirect effects occur when the contaminants adsorb to redox sensitive solid phases or precipitate 

with redox sensitive ligands. An example of the former involves the reductive dissolution of ferric 

oxide minerals, which can adsorb (complex) metals strongly. As the ferric oxide minerals 

dissolve, the adsorption potential of the soil is decreased. Another indirect effect of redox on 

contaminant sorption involves sulfur-ligand chemistry. Under reducing conditions, S(VI) (SO 4 

sulfate) will convert into S(II) (S 2- , 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 K d values by 

several orders of magnitude (Ames and Rai, 1978; Rai and Zachara, 1984). 

5.80 

2- ,

Table 5.20. Some of the more important aqueous- and solid-phase parameters 

affecting contaminant sorption. 1 

Element Important Aqueous- and Solid-Phase Parameters Influencing 

Contaminant Sorption 2 

Cd [Aluminum/Iron-Oxide Minerals], [Calcium], Cation Exchange 

Capacity, [Clay Mineral], [Magnesium], [Organic Matter], pH, Redox, 

[Sulfide] 

Cs [Aluminum/Iron-Oxide Minerals], [Ammonium], Cation Exchange 

Capacity, [Clay Mineral], [Mica-Like Clays], pH, [Potassium] 

Cr [Aluminum/Iron-Oxide Minerals], [Organic Matter], pH, Redox 

Pb [Aluminum/Iron-Oxide Minerals], [Carbonate, Fluoride, Sulfate, 

Phosphate], [Clay Mineral], [Organic Matter], pH, Redox 

Pu [Aluminum/Iron-Oxide Minerals], [Carbonate, Fluoride, Sulfate, 

Phosphate], [Clay Mineral], [Organic Matter], pH, Redox 

Rn None 

Sr Cation Exchange Capacity, [Calcium], [Carbonate], pH, [Stable 

Strontium] 

Th [Aluminum/Iron-Oxide Minerals], [Carbonate], [Organic Matter], pH 

3 H None 

U [Aluminum/Iron-Oxide Minerals], [Carbonate, Fluoride, Sulfate, 

Phosphate], [Clay Mineral], [Organic Matter], pH, Redox, [U] 

1 For groundwaters with low ionic strength and low concentrations of contaminant, 

chelating agents (e.g., EDTA), and natural organic matter. 

2 Parameters listed in alphabetical order. Square brackets represent concentration. 

5.81

6.0 REFERENCES 

Adriano, D. C. 1992. Biogeochemistry of Trace Metals. Lewis Publishers, Boca Raton, Florida. 

Ainsworth, C. C., and D. Rai. 1987. Selected Chemical Characterization of Fossil Fuel Wastes. 

EPRI EA-5321, Electric Power Research Institute, Palo Alto, California. 

Allard, B., and J. Rydberg. 1983. “Behavior of Plutonium in Natural Waters.” In Plutonium 

Chemistry, W. T. Carnall and G. R. Choppin (eds.), pp 275-295, ACS Symposium Series 216, 

American Chemical Society, Washington, D.C. 

Allison, J. D., D. S. Brown, and K. J. Novo-Gradac. 1991. MINTEQA2/PRODEFA2, A 

Geochemical Assessment Model for Environmental Systems: Version 3.0 User's Manual. 

EPA/600/3-91/021, U.S. Environmental Protection Agency, Athens, Georgia. 

Alloway, B. J. 1990. “Cadmium.” In Heavy Metals in Soils, B. J. Alloway (ed.), pp. 100-121, 

Blackie & Son, Glasgow, Scotland. 

Ames, L. L., J. E. McGarrah, B. A. Walker, and P. F. Salter. 1982. “Sorption of Uranium and 

Cesium by Hanford Basalts and Associated Secondary Smectites.” Chemical Geology, 

35:205-225. 

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 520/6-78-007A, prepared for the 

U.S. Environmental Protection Agency by the Pacific Northwest Laboratory, Richland, 

Washington. 

Artiola, J., and W. H. Fuller. 1979. “Effect of Crushed Limestone Barriers on Chromium 

Attenuation in Soils.” Journal of Environmental Quality, 8:503-510. 

Aston, S. R. 1980. “Evaluation of Chemical Forms of Plutonium in Seawater.” Marine 

Chemistry, 8:317-326. 

Ault, M. R. 1989. “Gamma Emitting Isotopes of Medical Origin Detected in Sanitary Waste 

Samples.” Radiation Protection Management, 6:48-52. 

Azizian, M. F., and P. O. Nelson. 1998. “Lead Sorption, Chemically Enhanced Desorption, and 

Equilibrium Modeling in an Iron-Oxide-Coated Sand and Synthetic Groundwater System.” In 

Adsorption of Metals by Geomedia. Variables, Mechanisms, and Model Applications, 

E. A. Jenne (ed.), pp. 166-180, Academic Press, San Diego, California. 

6.1

Baes, C. F., Jr., and R. E. Mesmer. 1976. The Hydrolysis of Cations. John Wiley and Sons, 

New York, New York. 

Baes, C. F., III, and R. D. Sharp. 1981. “Predicting Radionuclide Leaching from Root Zone Soil 

for Assessment Applications.” Transactions of the American Nuclear Society, 38:111 112. 

Baes, C. F., and R. D. Sharp. 1983. “A Proposal for Estimation of Soil Leaching Constants for 

Use in Assessment Models.” Journal of Environmental Quality, 12:17-28. 

Balistrieri, L. S., and J. W. Murray. 1982. “The Adsorption of Cu, Pb, Zn, and Cd on Goethite 

from Major Ion Seawater.” Geochimica et Cosmochimica Acta, 46:1253-1265. 

Ball, J. W., and D. K. Nordstrom. 1998. “Critical Evaluation and Selection of Standard State 

Thermodynamic Properties for Chromium Metal and Its Aqueous Ions, Hydrolysis Species, 

Oxides, and Hydroxides.” Journal of Chemical and Engineering Data, 43:895-918. 

Barney, G. S. 1984. “Radionuclide Sorption and Desorption Reactions with Interbed Materials 

from the Columbia River Basalt Formation.” In Geochemical Behavior of Radioactive Waste, 

G. S. Barney, J. D. Navratil, and W. W. Schulz (eds.), pp. 1-23. American Chemical Society, 

Washington, D.C. 

Bartlett, R. J., and B. James. 1979. “Behavior of Chromium in Soils. III. Oxidation.” Journal 

of Environmental Quality, 8:31-35. 

Bartlett, R. J., and J. M. Kimble. 1976. “Behavior of Chromium in Soils. I. Trivalent Forms.” 

Journal of Environmental Quality, 5:379-383. 

Bates, R. L., and J. A. Jackson (eds.). 1980. Glossary of Geology. American Geological 

Institute, Second Edition, Falls Church, Virginia. 

Belle, J., and R. M. Berman. 1984. “Application of Thorium Dioxide in Nuclear Power 

Reactors.” In Thorium Dioxide: Properties and Nuclear Applications. J. Belle and 

R. M. Berman (eds.), pp. 1-22, DOE/NE-0060, U.S. Department of Energy, 


Benjamin, M. M., and J. O. Leckie. 1980. “Adsorption of Metals at Oxide Interfaces: Effects on 

the Concentration of Adsorbate and Competing Metals.” In Contaminants and Sediments, 

Volume 2, R. A. Baker (ed.), pp. 305-332, Ann Arbor Science, Ann Arbor, Michigan. 

Benjamin, M. M., and J. O. Leckie. 1981. “Multiple-Site Adsorption of Cd, Cu, Zn, and Pb on 

Amorphous Iron Oxyhydroxide.” Journal of Colloid and Interface Science, 79:209-221. 

6.2

Bensen, D. W. 1960. Review of Soil Chemistry Research at Hanford. HW-67201. General 

Electric Company, Richland, Washington. 

Billon, A. 1982. “Fixation D’elements Transuraniens a Differents Degres D’oxydation Sur Les 

Argiles.” In Migration in the Terrestrial Environment of Long-lived Radionuclides from the 

Nuclear Fuel Cycle, pp. 167-176, IAEA-SM-257/32. International Atomic Energy Agency, 

Vienna, Austria. 

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. 

Blowes, D. W., and C. J. Ptacek. 1992. “Geochemical Remediation of Groundwater by 

Permeable Reactive Walls: Removal of Chromate by Reaction with Iron-Bearing Solids.” 

In Proceeding of the Subsurface Restoration Conference, June 21-24, 1992, Dallas, Texas, 

pp. 214-216, Rice University Press, Houston, Texas. 

Boggess, W. R., and B. G. Wixson. 1977. Lead in the Environment. NSF/RA-770214, National 

Science Foundation, Washington, D.C. 

Boggs, S., Jr., D. Livermore, and M. G. Seitz. 1985. Humic Substances in Natural Waters and 

Their Complexation with Trace Metals and Radionuclides: A Review. ANL-84-78, Argonne 

National Laboratory, Argonne, Illinois. 

Bondietti, E. A., S. A. Reynolds and M. H. Shanks. 1975. “Interaction of Plutonium with 

Complexing Substances in Soils and Natural Waters.” In Transuranium Nuclides in the 

Environment, pp. 273-287, IAEA-SM-199/51. International Atomic Energy Agency. Vienna, 

Austria. 

Bondietti, E. A., and J. R. Trabalka. 1980. “Evidence for Plutonium(V) in an Alkaline, 

Freshwater Pond.” Radioanalytical Letters, 43:169-176. 

Borovec, Z. 1981. “The Adsorption of Uranyl Species by Fine Clay.” Chemical Geology, 

32:45-58. 

Borovec, Z., B. Kribek, and V. Tolar. 1979. “Sorption of Uranyl by Humic Acids.” Chemical 

Geology, 27:39-46. 

Bovard, P., A. Grauby, and A. Saas. 1970. “Chelating Effect of Organic Matter and its Influence 

on the Migration of Fission Products.” In Proceedings of Symposium: Isotopes and 

Radiation in Soil Organic Matter Studies, pp. 471-495, STI/PUB-190, NSA 24:5659-5668, 

International Atomic Energy Agency (IAEA), 1968, CONF 680725, Vienna, Austria. 

6.3

Bowen, H. J. M. 1979. Environmental Chemistry of the Elements. Academic Press, London, 

England. 

Brady, P. V., R. T. Cygan, and K. L. Nagy. 1998. “Surface Charge and Metal Sorption to 

Kaolinite.” In Adsorption of Metals by Geomedia. Variables, Mechanisms, and Model 

Applications, E. A. Jenne (ed.), pp. 371-382, Academic Press, San Diego, California. 

Brady, P. V., B. P. Spalding, K. M. Krupka, R. D. Waters, P. Zhang, D. J. Borns, and 

W. D. Brady. 1999. Site Screening and Technical Guidance for Monitored Natural 

Attenuation at DOE Sites. SAND99-0464, Sandia National Laboratories, Albuquerque, 

New Mexico. 

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. 

Bruggenwert, M. G. M., and A. Kamphorst. 1979. “Survey of Experimental Information on 

Cation Exchange in Soil Systems.” In Soil Chemistry: B. Physico-Chemical Models, G. H. 

Bolt (ed.), Elsevier Scientific Publishing Company, New York, New York. 

Bruno, J., I. Casas, and I. Puigdomenech. 1988. “The Kinetics of Dissolution of UO 2 (s) Under 

Reducing Conditions.” Radiochimica Acta, 11:44-45. 

Bruno, J., I. Casas, and I. Puigdomenech. 1991. “The Kinetics of Dissolution of UO 2 Under 

Reducing Conditions and the Influence of an Oxidized Surface Layer (UO 2+x): Application of 

a Continuous Flow-through Reactor.” Geochimica et Cosmochimica Acta, 55:647-658. 

Bunzl, K., H. Flessa, W. Kracke, and W. Schimmack. 1995. “Association of Fallout 239+240 Pu and 

241 Am with Various Soil Components in Successive Layers of a Grassland Soil.” 

Environmental Science and Technology, 29:2513-2518. 

Cavallaro, N., and M. B. McBride. 1978. “Copper and Cadmium Adsorption Characteristics of 

Selected Acid and Calcareous Soils.” Soil Science Society of America Journal, 42:550-556. 

Cerling, T. E., and B. P. Spalding. 1982. “Distribution and Relationship of Radionuclides to 

Streambed Gravels in a Small Watershed.” Environmental Geology, 4:99-116. 

Charyulu, M. M., I. C. Pius, A. Kadam, M. Ray, C. K. Sivaramakrishnan, and S. K. Patil. 1991. 

“The Behavior of Plutonium in Aqueous Basic Media.” Journal of Radioanalytical and 

Nuclear Chemistry, 152: 479-485. 

Chen, C-C., C. Papelis, and K. F. Hayes. 1998. “Extended X-ray Absorption Fine Structure 

(EXAFS) Analysis of Aqueous Sr II Ion Sorption at Clay-Water Interfaces.” In Adsorption of 

6.4

Metals by Geomedia. Variables, Mechanisms, and Model Applications, E. A. Jenne (ed.), 

pp. 333-348, Academic Press, San Diego, California. 

Chisholm-Brause, C., S. D. Conradson, C. T. Buscher, P. G. Eller, and D. E. Morris. 1994. 

“Speciation of Uranyl Sorbed at Multiple Binding Sites on Montmorillonite.” Geochimica et 

Cosmochimica Acta, 58:3625-2631. 

Choppin, G. R. 1983. “Aspects of Plutonium Solution Chemistry.” In Plutonium Chemistry, W. 

T. Carnall and G. R. Choppin (eds.), pp. 213-230, ACS Symposium Series 216, American 

Chemical Society, Washington, D.C. 

Choppin, G. R., and J. W. Morse. 1987. “Laboratory Studies of Actinides in Marine Systems.” 

In Environmental Research on Actinide Elements, J. E. Pinder, J. J. Alberts, K. W. McLeod, 

and R. Gene Schreckhise (eds.), pp. 49-72, CONF-841142, Office of Scientific and Technical 

Information, U.S. Department of Energy, Washington, D.C. 

Chow, T. J. 1978. “Lead in Natural Waters.” In The Biogeochemistry of Lead in the 

Environment. Part A. Ecological Cycles., J. O. Nriagu (ed.), pp. 185-218, Elsevier/North 

Holland, New York, New York. 

Cleveland, J. M. 1979. The Chemistry of Plutonium. American Nuclear Society, LaGrange 

Park, Illinois. 

Coleman, N. T., R. J. Lewis, and D. Craig. 1963. “Sorption of Cesium by Soils and its 

Displacement by Salt Solutions.” Soil Science Society of America Proceedings, 22:390-294. 

Cotton, F. A., and G. Wilkinson. 1980. Advanced Inorganic Chemistry. A Comprehensive Text. 

John Wiley and Sons, New York, New York. 

Coughtrey, P. J., D. Jackson and M. C. Thorne. 1985. Radionuclide Distribution and Transport 

in Terrestrial and Aquatic Ecosystems. A Compendium of Data. A. A. Balkema, 

Netherlands. 

Cygan, R. T., K. L. Nagy, and P. V. Brady. 1998. “Molecular Models of Cesium Sorption on 

Kaolinite.” In Adsorption of Metals by Geomedia. Variables, Mechanisms, and Model 

Applications, E. A. Jenne (ed.), pp. 383-399, Academic Press, San Diego, California. 

Davis, J. A., and J. O. Leckie. 1978 “Surface Ionization and Complexation at the Oxide/Water 

Interface. II. Surface Properties of Amorphous Iron Oxyhydroxide and Adsorption of Metal 

Ions.” Journal of Colloid and Interface Science, 67:90-107. 

Davis, J. A., and J. O. Leckie. 1980. “Surface Ionization and Complexation at the Oxide/Water 

Interface. 3. Adsorption of Anions.” Journal of Colloid Interface Science, 74:32-43. 

6.5

Deer, W. A., R. A. Howie, and J. Zussman. 1967. Rock-Forming Minerals. Volume 1. Orthoand 

Ring Silicates. Longmans, London, England. 

Delegard, C. H. 1987. “Solubility of PuO 2·xH 2O in Alkaline Hanford High-Level Waste 

Solution.” Radiochimica Acta, 41:11-21. 

Delegard, C. H., G. S. Barney, and S. A. Gallagher. 1984. “Effects of Hanford High-Level 

Waste Components on the Solubility and Sorption of Cobalt, Strontium, Neptunium, 

Plutonium, and Americium.” In Geochemical Behavior of Disposed Radioactive Waste, G. 

S. Barney, J. D. Navratil, and W. W. Schulz (eds.), ACS Symposium Series 246, pp. 95-112. 

American Chemical Society, Washington, D.C. 

Douglas, L. A. 1989. “Vermiculites.” In Minerals in Soil Environments, J. B. Dixon and 

S. B. Week (eds.), Second Edition, pp. 635-674, Soil Science Society of America, Madison, 

Wisconsin. 

Driesens, F. C. M. 1986. “Ionic Solid Solutions in Contact with Aqueous Solutions.” In 

Geochemical Processes at Mineral Surfaces, J. A. Davis and K. F. Hayes (eds.), pp. 524-560, 

ACS Symposium Series 323. American Chemical Society, Washington, D.C. 

Duff, M. C., and C. Amrhein. 1996. “Uranium(VI) Adsorption on Goethite and Soil in 

Carbonate Solutions.” Soil Science Society of America Journal, 60(5):1393-1400. 

Duff, M. C., D. B. Hunter, I. R. Triay, P. M. Bertsch, D. T. Reed, S. R. Sutton, 

G. Shea-McCarthy, J. Kitten, P. Eng, S. J. Chipera, and D. T. Vaniman. 1999. “Mineral 

Associations and Average Oxidation States of Sorbed Pu on Tuff.” Environmental Science 

and Technology, 32:2163-2169. 

Duram, W. H., J. D. Hem, and S. G. Heidel. 1971. Reconnaissance of Selected Minor Elements 

in Surface Waters of the United States, October 1970. U.S. Geological Survey Circular 643, 

U.S. Geological Survey, Alexandria, Virginia. 

Eary, L. E., and D. Rai. 1987. “Kinetics of Chromium(III) Oxidation to Chromium(VI) by 

Reaction with Manganese Dioxide.” Environmental Science and Technology, 21:1187-1193. 

Eary, L. E., and D. Rai. 1989. “Kinetics of Chromate Reduction by Ferrous Ions Derived from 

Hematite and Biotite at 25ºC.” American Journal of Science, 289:180-213. 

EPA (U.S. Environmental Protection Agency). 1992. Background Document for Finite Source 

Methodology for Wastes Containing Metal. HWEP-S0040, U.S. Environmental Protection 

Agency, Office of Solid Waste, Washington, D.C. 

6.6

EPA (U.S. Environmental Protection Agency). 1996. Soil Screening Guidance: Technical 

Background Document. EPA/540/R-96/018, U.S. Environmental Protection Agency, 


EPA/DOE/NRC (Cooperative Effort by the U.S. Environmental Protection Agency, U.S. 

Department of Energy, and U.S. Nuclear Regulatory Commission). 1993. Environmental 

Characteristics of EPA, NRC, and DOE Sites Contaminated with Radioactive Substances. 

EPA 402-R-93-011, U.S. Environmental Protection Agency, Washington, D.C. 

Evans, E. J. 1956. Plutonium Retention in Chalk River Soil. CRHP-660, Chalk River 

Laboratory, Chalk River, Canada. 

Falck, W. E. 1991. CHEMVAL Project. Critical Evaluation of the CHEMVAL Thermodynamic 

Database with Respect to its Contents and Relevance to Radioactive Waste Disposal at 

Sellafield and Dounreay. DoE/HMIP/RR/92.064, Department of Environment, Her 

Majesty’s Stationary Office, London, England. 

Faure, G., and J. L. Powell. 1972. Strontium Isotope Geology. Springer-Verlag, Berlin, 

Germany. 

Felmy, A. R., D. Rai, and M. J. Mason. 1991. “The Solubility of Hydrous Thorium(IV) Oxide in 

Chloride Media: Development of an Aqueous Ion-Interaction Model.” Radiochimica Acta, 

55:177-185. 

Felmy, A. R., D. Rai, and D. A. Moore. 1993. “The Solubility of (Ba,Sr)SO 4 Precipitates: 

Thermodynamic Equilibrium and Reaction Path Analysis.” Geochimica et Cosmochimica 

Acta, 57:4345-4363. 

Fisher, N. S., S. W. Fowler, F. Boisson, J. Carroll, K. Rissanen, B. Salbu, T. G. Sazykina, and 

K-L.-Sjoeblom. 1999. “Radionuclide Bioconentrations of Factors and Sediment Partition 

Coefficients in Arctic Sea Subject to Contamination from Dumped Nuclear Wastes.” 

Environmental Science and Technology, 32:1979-1982. 

Forbes, E. A., A. M. Posner, and J. P. Quirk. 1976. “The Specific Adsorption of Divalent Cd, 

Co, Cu, Pb, and Zn on Goethite.” Journal of Soil Science, 27:154-166. 

Freeze, R. A., and J. A. Cherry. 1979. Groundwater. Prentice-Hall, Inc., Englewood Cliffs, 

New Jersey. 

Frondel, C. 1958. Systematic Mineralogy of Uranium and Thorium. Geological Survey 

Bulletin 1064, U.S. Geological Survey, Washington, D.C. 

6.7

Gadde, R. R., and H. A. Laitinen. 1974. “Study of the Sorption of Lead by Hydrous Ferric 

Oxide.” Environmental Letters, 5:223-235. 

Gambrell, R. P., R. A. Khalid, M. B. Verloo, and W. H Patrick, Jr. 1977. Transformations of 

Heavy Metals and Plant Nutrients in Dredged Sediments as Affected by Oxidation-Reduction 

Potential and pH. Volume II: Materials and Methods, Results and Discussion, Contract Rep. 

D-77-4, CE, U.S. Army Engineer Waterways Experiment Station, Vicksburg, Mississippi. 

Gascoyne, M. 1982. “Geochemistry of the Actinides and Their Daughters.” In Uranium Series 

Disequilibrium: Applications to Environmental Problems, M. Ivanovich and R. S. Harmon 

(eds.), pp. 33-55, Clarendon Press, Oxford, England. 

Gerritse, R. G., R. Vriesema, J. W. Dalenberg, and H. P. De Roos. 1982. “Effect of Sewage 

Sludge on Trace Element Mobility in Soils.” Journal of Environmental Quality. 11:359-364. 

Gesell, T. F., and W. M. Lowder (eds.). 1980. Natural Radiation Environment III. Volumes 1 

and 2. Proceedings of Symposium Held at Houston, Texas, April 23-28. 1978. U.S. 

Department of Energy CONF 780422, National Technical Information Service, Springfield, 

Virginia. 

Giesy, J. G., Jr. 1980. “Cadmium Interactions with Naturally Occurring Organic Ligands.” In 

Cadmium in the Environment - Part 1 Ecological Cycling, J. O. Nriagu (ed.), pp. 237-256, 

John Wiley and Sons, New York, New York. 

Giesy, J. P., G. J. Leversee, and D. R. Williams. 1977. “Effects of Natural Occurring Aquatic 

Organic Fractions on Cadmium Toxicity to Simocephalus Serrulatus (Daphnidae) and 

Gambusia Affinis (Poeciliidae).” Water Research, 12:1013-1020. 

Glover, P. A., F. J. Miner and W. O. Polzer. 1976. “Plutonium and Americium Behavior in the 

Soil/Water Environment. I. Sorption of Plutonium and Americium by Soils.” In Proceedings 

of Actinide-Sediment Reactions Working Meeting, Seattle, Washington, pp. 225-254, BNWL- 

2117, Battelle Pacific Northwest Laboratories, Richland, Washington. 

Goldschmidt, V. M. 1954. Geochemistry. Clarendon Press, Oxford, England. 

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. 

Graves, B. (ed.). 1987. Radon, Radium, and Other Radioactivity in Ground Water. 

Hydrogeologic Impact and Application to Indoor Airborne Contamination. Proceedings of 

the NWWA Association Conference, April 7-9, 1987, Somerset, New Jersey. Lewis 

Publishers, Chelsea, Michigan. 

6.8

Griffin, R. A., A. K. Au, and R. R. Frost. 1977. “Effect of pH on Adsorption of Chromium from 

Landfill-Leachate by Clay Minerals.” Journal of Environmental Science Health, 12:431-449. 

Griffin, R. A., and N. F. Shimp. 1976. “Effect of pH on Exchange-Adsorption or Precipitation of 

Lead from Landfill Leachates by Clay Minerals.” Environmental Science and Technology, 

10:1256-1261. 

Haji-Djafari, S., P. E. Antommaria, and H. L. Crouse. 1981. “Attenuation of Radionuclides and 

Toxic Elements by In Situ Soils at a Uranium Tailings Pond in central Wyoming.” In 

Permeability and Groundwater Contaminant Transport, T. F. Zimmie and C. O. Riggs (eds.), 

pp 221-242. ASTM STP 746, American Society of Testing Materials, Washington, D.C. 

Hammond, P. B. 1977. “Human Health Implications.” In Lead in the Environment, W. R. 

Bogges and B. G. Wixson (eds.), pp. 195-198, NSF/RA-770214, National Science 

Foundation, Washington, D.C. 

Hassett, J. J. 1974. “Capacity of Selected Illinois Soils to Remove Lead from Aqueous 

Solution.” Communications in Soil Science and Plant Analysis, 5:499-505. 

Hem, J. C. 1977. “Reactions of Metal Ions at Surfaces of Hydrous Iron Oxide.” Geochimica et 


Hem, J. D. 1972. “Chemistry and Occurrence of Cadmium and Zinc in Surface Water and 

Ground Water.” Water Resources and Research, 8:661-679. 

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. 

Hildebrand, E. E., and W. E. Blum. 1974. “Lead Fixation by Clay Minerals.” 

Naturewissenschaften, 61:169-170. 

Hsi, C-K. D., and D. Langmuir. 1985. “Adsorption of Uranyl onto Ferric Oxyhydroxides: 

Application of the Surface Complexation Site-binding Model.” Geochimica et Cosmochimica 

Acta, 49:1931-1941. 

Huang, C. P., H. Z. Elliott, and R. M. Ashmead. 1977. “Interfacial Reactions and the Fate of 

Heavy Metals in Soil-Water Systems.” Journal of Water Pollution Control Federation, 

49:745-755. 

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. 

6.9

IAEA (International Atomic Energy Agency). 1979. Behavior of Tritium in the Environment 

(Proceedings of the International Symposium on the Behavior of Tritium in the Environment 

Jointly Organized by the International Atomic Energy Agency and the OECD Nuclear 

Energy Agency and Held in San Francisco, 16-20 October 1978. Proceedings Series, 

International Atomic Energy Agency, Vienna, Austria. 

Idiz, E. F., D. Carlisle, and I. R. Kaplan. 1986. “Interaction Between Organic Matter and Trace 

Metals in a Uranium Rich Bog, Kern County, California, U.S.A.” Applied Geochemistry, 

1:573-590. 

Izrael, Y. A., and F. Y. Rovinskii. 1970. Hydrological Uses of Isotopes Produced in 

Underground Nuclear Explosions for Peaceful Purposes. UCRL-Trans-10458, International 

Atomic Energy Agency (IAEA), Vienna. 

Jackson, A. P., and B. J. Alloway. 1992. “Transfer of Cadmium from Soils to the Human Food 

Chain.” In Biogeochemistry of Trace Metals, D. C. Adriano (ed.), pp.-109-147, Lewis 

Publishers, Boca Raton, Florida. 

Jenne, E. A., J. W. Ball, J. M. Burchard, D. V. Vivit, and J. H. Barks. 1980. “Geochemical 

Modeling; Apparent Solubility Controls on Ba, Zn, Cd, Pb, and F in Waters of the Missouri 

Tri-State Mining Area.” Trace Substances in Environmental Health, 14:353-361. 

John, M. K. 1971. “Influence of Soil Characteristics of Adsorption and Desorption of 

Cadmium.” Environmental Letters, 2:173-179. 

Juo, A. S. R., and S. A Barber. 1970. “The Retention of Strontium by Soils as Influenced by pH, 

Organic Matter and Saturated Cations.” Soil Science, 109:143-148. 

Juste, C., and M. Mench. 1992. “Long-term Application of Sewage Sludge and Its Effect on 

Metal Uptake by Crops.” In Biogeochemistry of Trace Metals, D. C. Adriano (ed.), 

pp. 159-194, Lewis Publishers, Boca Raton, Florida. 

Kaplan, D. I., T. L. Gervais, and K. M. Krupka. 1998. “Uranium(VI) Sorption to Sediments 

Under High pH and Ionic Strength Conditions.” Radiochimica Acta, 80:201-211. 

Kaplan, D. I., R. J. Serne, A. T. Owen, J. Conca, T. W. Wietsma, and T. L. Gervais. 1996. 

Radionuclide Adsorption Distribution Coefficient Measured in Hanford Sediments for the 

Low Level Waste Performance Assessment Project. PNNL-11385, Pacific Northwest 

National Laboratory, Richland, Washington. 

Kargbo, D. M., D. S. Fanning, H. I. Inyang, and R. W. Duell. 1993. “Environmental Significance 

of Acid Sulfate “Clays” as Waste Covers.” Environmental Geology, 22:218-226. 

6.10

+ 

Keeney-Kennicutt, W. L., and J. W. Morse. 1985. “The Redox Chemistry of Pu(V)O2 Interaction with Common Mineral Surfaces in Dilute Solutions and Seawater.” Geochimica 


Kent, D. B., V. S. Tripathi, N. B. Ball, J. O. Leckie, and M. D. Siegel. 1988. Surface- 

Complexation Modeling of Radionuclide Adsorption in Subsurface Environments. 

NUREG/CR-4807, U.S. Nuclear Regulatory Commission, Washington, D.C. 

Khalid, R. A. 1980. “Chemical Mobility of Cadmium in Sediment-Water Systems.” In Cadmium 

in the Environment - Part 1. Ecological Cycling. J. O. Nriagu (ed.), pp. 257-298, John Wiley 

and Sons, New York, New York. 

Kharkar, D. P., K. K. Turekian, and K. K. Bertine. 1968. “Stream Supply of Dissolved Silver, 

Molybdenum, Antimony, Selenium, Chromium, Cobalt, Rubidium, and Cesium to the 

Oceans.” Geochimica et Cosmochimica Acta, 32:285-298. 

Kim, J. J. 1986. “Chemical Behavior of Transuranic Elements in Aquatic Systems.” In 

Handbook on the Physics and Chemistry of the Actinides, A. J. Freeman and C. Keller (eds.), 

pp. 413-455, Elsevier Science Publishers, Amsterdam, Holland. 

Kokotov, Y. A., and R. F. Popova. 1962. “Sorption of Long-Lived Fission Products by Soils 

and Argillaceous Minerals III: Selectivity of Soils and Clays Toward 90 Sr Under Various 

Conditions.” Soviet Radiochemistry, 4:292-297. 

Korte, N. E., J. Skopp, W. H. Fuller, E. E. Niebla, and B. A. Alesii. 1976. “Trace Element 

Movement in Soils: Influence of Soil Physical and Chemical Properties.” Soil Science 

Journal, 122:350-359. 

Krupka, K. M. and R. J. Serne. 1998. Effects on Radionuclide Concentrations by 

Cement/Ground-Water Interactions to Support Performance Assessment of Low-Level 

Radioactive Waste Disposal Facilities. NUREG/CR-6377, Pacific Northwest National 

Laboratory, Richland, Washington. 

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. 

Lagerwerff, J. V., and D. L. Brower. 1973. “Exchange Adsorption or Precipitation of Lead in 

Soils Treated With Chlorides of Aluminum, Calcium and Sodium.” Soil Science Society of 

America Proceedings, 27:1951-1954. 

6.11

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. 

Langmuir, D. 1978. “Uranium Solution-mineral Equilibria at Low Temperatures with 

Applications to Sedimentary Ore Deposits.” Geochimica et Cosmochimica Acta, 42:547-569. 

Langmuir, D., and J. S. Herman. 1980. “The Mobility of Thorium in Natural Waters at Low 

Temperatures.” Geochimica et Cosmochimica Acta, 44:1753-1766. 

Leckie, J. O., M. M. Benjamin, K. Hayes, G. Kaufman, and S. Altman. 1980. 

Adsorption/Coprecipitation of Trace Elements from Water With Iron Oxyhydroxides. 

EPRI-RP-910, Electric Power Research Institute, Palo Alto, California. 

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. 

Lemire, R. J., and P. R. Tremaine. 1980. “Uranium and Plutonium Equilibria in Aqueous 

Solutions to 200EC.” Journal of Chemical Engineering Data, 25:361-370. 

Levi-Minzi, R., G. F. Soldatini, and R. Riffaldi. 1976. “Cadmium Adsorption by Soils.” Journal 

of Soil Science, 27:10-15. 

Levy, R., and C. W. Francis. 1976. “Adsorption and Desorption of Cadmium by Synthetic and 

Natural Organo-Clay Complexes.” Geoderma, 18:193-205. 

Lindenmeier, C. W., R. J. Serne, J. L. Conca, A. T. Owen, and M. I. Wood. 1995. Solid Waste 

Leach Characteristics and Contaminant-Sediment Interactions Volume 2: Contaminant 

Transport Under Unsaturated Moisture Contents. PNL-10722, Pacific Northwest 


Lindsay, W. L. 1979. Chemical Equilibria in Soils. J. Wiley and Sons, New York, New York. 

MacNaughton, M. G. 1977. “Adsorption of Chromium(VI) at the Oxide-Water Interface.” In 

Biological Implications of Metals in the Environment, H. Drucker and R. F. Wildung (eds.), 

pp. 244-253, CONF-750929, National Technical Information Service, Springfield, Virginia. 

Mattigod, S. V., and A. L. Page. 1983. “Assessment of Metal Pollution in Soils.” In Applied 

Environmental Geochemistry, I. Thornton (ed.), pp. 355-394, Academic Press, New York, 

New York. 

6.12

Mattigod, S. V., A. L. Page, and I. Thornton. 1986. “Identification of Some Trace Metal 

Minerals in a Mine-Waste Contaminated Soil.” Soil Science Society of America Journal, 

50:254-258. 

McBride, M. B. 1980. “Chemisorption of Cd 2+ on Calcite Surfaces.” Soil Science Society of 

America Journal, 44:26-28. 

McBride, M. B., L. D. Tyler, and D. A. Hovde. 1981. “Cadmium Adsorption by Soils and 

Uptake by Plants as Affected by Soil Chemical Properties.” Soil Science Society of America 

Journal, 45:739-744. 

McHenry, J. R. 1954. Adsorption and Retention of Cesium by Soils of the Hanford Project. 

HW-S1011, Westinghouse Hanford Company, Richland, Washington. 

McHenry, J. R. 1958. “ Ion Exchange Properties of Strontium in a Calcareous Soil.” Soil 

Science Society of America Proceedings, 22:514-518. 

McKinley, J. P., J. M. Zachara, S. C. Smith, and G. D. Turner. 1995. “The Influence of Uranyl 

Hydrolysis and Multiple Site-Binding Reactions on Adsorption of U(VI) to Montmorillonite.” 

Clays and Clay Minerals, 43(5):586-598. 

McLean, J. E., and B. E. Bledsoe. 1992. Behavior of Metals in Soils. EPA/540/S-92/018, U.S. 

Environmental Protection Agency, Ada, Oklahoma. 

Meybeck, M. 1982. “Carbon, Nitrogen, and Phosphorous Transport by World Rivers.” 

American Journal of Science, 282:401-450. 

Morris, D. E., C. J. Chisholm-Brause, M. E. Barr, S. D. Conradson, and P. G. Eller. 1994. 

2+ 

“Optical Spectroscopic Studies of the Sorption of UO2 Species on a Reference Smectite.” 

Geochimica et Cosmochimica Acta, 58:3613-3623. 

Nakashima, S., J. R. Disnar, A. Perruchot, and J. Trichet. 1984. “Experimental Study of 

Mechanisms of Fixation and Reduction of Uranium by Sedimentary Organic Matter Under 

Diagenetic or Hydrothermal Conditions.” Geochimica et Cosmochimica Acta, 48:2321-2329. 

Nakayama, E., T. Kumamoto, S. Tsurubo, and T. Fujinaga. 1981. “Chemical Speciation of 

Chromium in Sea Water. Part 2. Effects of Manganese Oxides and Reducible Organic 

Materials on the Redox Processes of Chromium.” Analytica Chimica Acta, 130:401-404. 

Nash, K., S. Fried, A. M. Freidman, and J. C. Sullivan. 1981. “Redox Behavior, Complexing, 

and Adsorption of Hexavalent Actinides by Humic Acid and Selected Clays.” Environmental 

Science and Technology, 15:834-837. 

6.13

Navrot, J., A. Singer, and A. Banin. 1978. “Adsorption of Cadmium and Its Exchange 

Characteristics in Some Israeli Soils.” Journal of Soil Science, 29:505-511. 

Nelson, D. M., R. P. Larson, and W. R. Penrose. 1987. “Chemical Speciation of Plutonium in 

Natural Waters.” In Environmental Research on Actinide Elements, J. E. Pinder, J. J. 

Alberts, K. W. McLeod, and R. Gene Schreckhise (eds.), pp. 27-48, CONF-841142 

(DE86008713), Office of Scientific and Technical Information, U.S. Department of Energy, 


Nelson, D. M., and M. B. Lovett. 1980. “Measurements of the Oxidation State and 

Concentration of Plutonium in Interstitial Waters in the Irish Sea.” In Impacts of 

Radionuclide Releases into the Marine Environment, IAEA Staff (ed.), pp. 105-118, 

International Atomic Energy Agency (IAEA), Vienna, Austria. 

Nelson D. M., and K. A. Orlandini. 1979. Identification of Pu(V) in Natural Waters. 

ANL-79-65, Argonne National Laboratory, Argonne, Illinois. 

Nielson, K. K., V. C. Rogers, and G. W. Gee. 1984. “Diffusion of Radon Through Soils: A 

Pore Distribution Model.” Soil Science Society of America Journal, 48:482-487. 

Nishita, H. 1978. “Extractability of Plutonium-238 and Curium-242 from a Contaminated Soil as 

a Function of pH and Certain Soil Components. CH 3COOH-NH 4OH System.” In 

Environmental Chemistry and Cycling Process, pp. 403-416, CONF-760429. Technical 

Information Center, U.S. Department of Energy, Washington, D.C. 

NRC (U.S. Nuclear Regulatory Commission). 1993. Site Decommissioning Management Plan. 

NUREG-1444, U.S. Nuclear Regulatory Commission, Washington, D.C. 

Nriagu, J. O. 1978. The Biogeochemistry of Lead in the Environment. Part A. Ecological 

Cycles. Elsevier/North-Holland, New York, New York. 

Nriagu, J. O. 1980a. Cadmium in the Environment - Part 1. Ecological Cycling. John Wiley 


Nriagu, J. O. 1980b. “Production, Uses, and Properties of Cadmium.” In Cadmium in the 

Environment - Part 1 Ecological Cycling, J. O. Nriagu (ed.), pp. 35-70, John Wiley 


Nriagu, J. O., and P. B. Moore. 1984. Phosphate Minerals. Springer-Verlag, New York, New 

York. 

Nriagu, J. O., and E. Nieboer (eds.). 1988. Chromium in the Natural and Human Environments, 

Volume 20. John Wiley & Sons, New York, New York. 

6.14

Oscarson, D. W., and H. B. Hume. 1998. “Effect of Solid:Liquid Ratio on the Sorption of Sr 2+ 

and Cs + on Bentonite.” In Adsorption of Metals by Geomedia. Variables, Mechanisms, and 

Model Applications, E. A. Jenne (ed.), pp. 277-289, Academic Press, San Diego, California. 

Östhols, E. 1995. “Thorium Sorption on Amorphous Silica.” Geochimica et Cosmochimica 

Acta, 59:1235-1249. 

Östhols, E., J. Bruno, and I. Grenthe. 1994. “On the Influence of Carbonate on Mineral 

Dissolution: III. The Solubility of Microcrystalline ThO 2 in CO 2-H 2O Media.” Geochimica et 


Overstreet, R., and C. Krishnamurthy. 1950. “An Experimental Evaluation of Ion-Exchange 

Relationships.” Soil Science, 69:41-50. 

Pabalan, R. T., D. R. Turner, F. P. Bertetti, and J. D. Prikryl. 1998. “Uranium VI Sorption onto 

Selected Mineral Surfaces: Key Geochemical Parameters.” In Adsorption of Metals by 

Geomedia. Variables, Mechanisms, and Model Applications, E. A. Jenne (ed.), pp. 99-130, 

Academic Press, San Diego, California. 

Palmer, C. D., and R. W. Puls. 1994. Natural Attenuation of Hexavalent Chromium in 

Groundwater and Soils. EPA/540/S-94/505, U.S. Environmental Protection Agency, Ada, 

Oklahoma. 

Palmer, C. D. and P. R. Wittbrodt. 1991. “Processes Affecting the Remediation of Chromium- 

Contaminated Sites.” Environmental Health Perspectives, 92:25-40. 

Payne, T. E., G. R. Lumpkin, and T. D. Waite. 1998. “Uranium VI Adsorption on Model 

Minerals: Controlling Factors and Surface Complexation Modeling.” In Adsorption of 



Peters, R. W., and L. Shem. 1992. “Adsorption/Desorption Characteristics of Lead on Various 

Types of Soil.” Environmental Progress, 11:234-240. 

Petruzelli, G., G. Guidi, and L. Lubrano. 1978. “Organic Matter as an Influencing Factor on 

Copper and Cadmium Adsorption by Soils.” Water Air and Soil Pollution, 9:263-268. 

Pittwell, L. R. 1974. “Metals Coordinated by Ligands Normally Found in Natural Waters.” 

Journal of Hydrology, 21:301-304. 

Prout, W. E. 1958. “Adsorption of Radioactive Wastes by Savannah River Plant Soil.” Soil 

Science, 84:13-17. 

6.15

Rai, D., A. R. Felmy, D. A. Moore, and M. J. Mason. 1995. “The Solubility of Th(IV) and 

U(IV) Hydrous Oxides in Concentrated NaHCO 3 and Na 2CO 3 Solutions.” In Scientific Basis 

for Nuclear Waste Management XVIII, Part 2, T. Murakami and R. C. Ewing (eds.), 

pp. 1143-1150, Materials Research Society Symposium Proceedings, Volume 353, Materials 

Research Society, Pittsburgh, Pennsylvania 

Rai, D., B. M. Sass, and D. A. Moore. 1987. “Chromium(III) Hydrolysis Constants and 

Solubility of Chromium(III) Hydroxide.” Inorganic Chemistry, 26:345-349. 

Rai, D., R. J. Serne, and D. A. Moore. 1980a. “Solubility of Plutonium Compounds and Their 

Behavior in Soils.” Soil Science Society of America Journal, 44:490-495. 

Rai D., R. J. Serne, and J. L. Swanson. 1980b. “Solution Species of Plutonium in the 

Environment.” Journal of Environmental Quality, 9:417-420. 

Rai, D., and J. M. Zachara. 1984. Chemical Attenuation Rates, Coefficients, and Constants in 

Leachate Migration. Volume 1: A Critical Review. EA-3356, Electric Power Research 

Institute, Palo Alto, California. 

Rai, D., J. M. Zachara, L. E. Eary, C. C. Ainsworth, J. D. Amonette, C. E. Cowan, R. W. 

Szelmeczka, C. T. Resch, R. L. Schmidt, S. C. Smith, and D. C. Girvin. 1988. Chromium 

Reactions in Geologic Materials. EPRI-EA-5741, Electric Power Research Institute, Palo 

Alto, California. 

Rai, D., J. M. Zachara, L. E. Eary, D. C. Girvin, D. A. Moore, C. T. Resch, B. M. Sass, and R. 

L. Schmidt. 1986. Geochemical Behavior of Chromium Species. EPRI-EA-4544. Electric 

Power Research Institute, Palo Alto, California. 

Rama, and W. S. Moore. 1984. “Mechanism of Transport of U-Th Series Radioisotopes from 

Solids into Ground Water.” Geochimica et Cosmochimica Acta, 48:395-399. 

Read, D., T. A. Lawless, R. J. Sims, and K. R. Butter. 1993. “Uranium Migration Through 

Intact Sandstone Cores.” Journal of Contaminant Hydrology, 13:277-289. 

Redden, G. D., J. Li, and J. Leckie. 1998. “Adsorption of U VI and Citric Acid on Goethite, 

Gibbsite, and Kaolinite: Comparing Results for Binary and Ternary Systems.” In Adsorption 

of Metals by Geomedia. Variables, Mechanisms, and Model Applications, E. A. Jenne (ed.), 


Reid, J. C., and B. McDuffie. 1981. “Sorption of Trace Cadmium in Clay Minerals and River 

Sediments: Effects of pH and Cd(II) Concentrations in a Synthetic River Water Medium.” 

Water Air and Soil Pollution, 15:375-386. 

6.16

Relyea, J. F. and D. A. Brown. 1978. “Adsorption and Diffusion of Plutonium in Soil.” In 

Environmental Chemistry and Cycling Process, CONF-760429, Technical Information 

Center, U.S. Department of Energy, Washington, D.C. 

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. 

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. 

Rhodehamel, E. C., V. B. Kron, and V. M. Dougherty. 1971. Bibliography of Tritium Studies 

Related to Hydrology Through 1966. Geological Survey Water Supply Paper 1900, U.S. 

Geological Survey, Washington, D.C. 

Rhodes, D. W. 1957. “The Effect of pH on the Uptake of Radioactive Isotopes from Solution by 

a Soil.” Soil Science Society of America Proceedings, 21:389-392. 

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. 

Richard, F. C., and A. C. M. Bourg. 1991. “Aqueous Geochemistry of Chromium: A Review.” 

Water Research, 25:807-816. 

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. 219-284, Elsevier/North-Holland, New York, New York. 

Richards, L. A. 1954. Diagnosis and Improvement of Saline and Alkali Soils. Agricultural 

Handbook 60, U.S. Department of Agriculture, Washington, D.C. 

Robbins, J. A. 1978. “Geochemical and Geophysical Applications of Radioactive Lead.” In The 

Biogeochemistry of Lead in the Environment. Part A. Ecological Cycles, J. O. Nriagu (ed.), 

pp. 285-394, Elsevier/North-Holland, New York, New York. 

Roberts H., and R. G. Menzel. 1961. “Availability of Exchange and Nonexchangeable Sr-90 to 

Plants.” Agriculture and Food Chemistry, 9:95-98. 

Rosentreter, J. J., H. S. Quarder, R. W. Smith, and T. McLing. 1998. “Uranium Sorption onto 

Natural Sands as a Function of Sediment Characteristics and Solution pH.” In Adsorption of 

6.17



Routson, R. C. 1973. A Review of Studies on Soil-Waste Relationships on the Hanford 

Reservation from 1944 to 1967. BNWL-1464, Pacific Northwest Laboratory, Richland, 

Washington. 

Routson, R. C., G. S. Barney, and R. M. Smith. 1980. Hanford Site Sorption Studies for the 

Control of Radioactive Wastes: A Review. WHO-SA-155, Rev. 1, Rockwell Hanford 

Operations, Richland, Washington. 

RTI (Research Triangle Institute). 1994. Chemical Properties for Soil Screening Levels. 

Prepared for the Environmental Protection Agency, Office of Emergency and Remedial 

Response, Washington, D.C. 

Ruby, M. V., A. Davis, and A. Nicholson. 1994. “In Situ Formation of Lead Phosphates in Soils 

as a Method to Immobilize Lead.” Environmental Science and Technology, 28:646-654. 

Ryan, J. L., and D. Rai. 1987. “Thorium(IV) Hydrous Oxide Solubility.” Inorganic Chemistry, 

26:4140-4142. 

Sanchez, A. L., J. W. Murray, and T. H. Sibley. 1985. “The Adsorption of Pu(IV) and (V) of 

Goethite.” Geochimica et Cosmochimica Acta, 49:2297-2307. 

Sandino, A., and J. Bruno. 1992. “The Solubility of (UO 2) 3(PO 4) 2·4H 2O(s) and the Formation of 

U(VI) Phosphate Complexes: Their Influence in Uranium Speciation in Natural Waters.” 


Santillan-Medrano, J., and J. J. Juriank. 1975. “The Chemistry of Lead and Cadmium in Soil: 

Solid Phase Formation.” Soil Science Society of America Proceedings, 39:851-856. 

Sass, B. M., and D. Rai. 1987. “Solubility of Amorphous Chromium(III)-Iron(III) Hydroxide 

Solid Solutions.” Inorganic Chemistry, 26:2228-2232. 

Sax, N. I., and R. J. Lewis, Sr. 1987. Hawley’s Condensed Chemical Dictionary. Eleventh 

Edition, Van Nostrand Reinhold Company, New York, New York. 

Scheider, K. J., and A. M. Platt (eds.). 1974. High-Level Waste Management Alternatives. 

Volume I. BNWL-1900, pp. 2.D.7-2.D.8, Pacific Northwest National Laboratory, Richland, 

Washington. 

Schulz, R. K. 1965. “Soil Chemistry of Radionuclides.” Health Physics, 11:1317-1324. 

6.18

Schultz, R. K., R. Overstreet, and I. Barshad. 1960. “On the Chemistry of Cesium 137.” Soil 

Science, 89:16-27. 

Schulz, R. K., and H. H. Riedel. 1961. “Effect of Aging on Fixation of Strontium-90 by Soils.” 

Soil Science, 91:262-264. 

Schwertmann, U., and R. M. Taylor. 1989. “Iron Oxides.” In Minerals in Soil Environments, 

Second Edition. J. B. Dixon and S. B. Week (eds.), pp. 379-438, Soil Science Society of 

America, Madison, Wisconsin. 

Scrudato, R. J., and E. L. Estes. 1975. “Clay-Lead Sorption Studies.” Environmental Geology., 

1:167-170. 

Serkiz, S. M. And W. H. Johnson. 1994. Uranium Geochemistry in Soil and Groundwater at 

the F and H Seepage Basins (U). EPD-SGS-94-307, Westinghouse Savannah River 

Company, Savannah River Site, Aiken, South Carolina. 

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, Volume 1, Pacific Northwest Laboratory, Richland, 

Washington. 

Serne, R. J. 1977. “Geochemical Distribution of Selected Trace Metals in San Francisco Bay 

Sediments.” In Biological Implications of Metals in the Environment. Proceedings of the 

Fifteenth Annual Hanford Life Sciences Symposium at Richland, Washington, 

September 29-October 1, 1975. H. Drucker and R. E. Wildung (eds.), pp. 280-296, 

CONF-75029, Energy Research and Development Administration, Washington, D.C. 

Serne, R. J., and V. L. LeGore. 1996. Strontium-90 Adsorption-Desorption Properties and 

Sediment Characterization at the 100 N-Area. PNL-10899, Pacific Northwest National 


Shanbhag, P. M., and G. R. Choppin. 1981. “Binding of Uranyl by Humic Acid.” Journal of 

Inorganic Nuclear Chemistry, 43:3369-3372. 

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. 

Silver, G. L. 1983. “Comment on the Behavior of the Chemical Forms of Plutonium in Seawater 

and Other Aqueous Solutions.” Marine Chemistry, 12:91-96. 

6.19

Simpson. H. J., R. M. Trier, Y. H. Li, R. F. Anderson, and A. L. Herczeg. 1984. Field 

Experiment Determinations of Distribution Coefficient of Actinide Elements in Alkaline Lake 

Environments. NUREG/CR-3940, prepared for the U.S. Nuclear Regulatory Commission by 

Columbia University, Palisades, New York. 

Singh, S. S. 1979. “Sorption and Release of Cadmium in Some Canadian Soils.” Canadian 

Journal of Soil Science, 59:119-130. 

Singh, B., and G. S. Sekjon. 1977. “Adsorption, Desorption, and Solubility Relationships of 

Lead and Cadmium in Some Alkaline Soils.” Journal of Soil Science, 28:271-275. 

Skougstad, M. W., and C. A. Horr. 1963. “Occurrence and Distribution of Strontium in Natural 

Waters.” U.S. Geological Survey Water Supply Paper 1496-D, pp. D55-D97, U.S. 

Geological Survey, Alexandria, Virginia. 

Smith, J. T., and R. N. J. Comans. 1996. “Modelling the Diffusive Transport and Remobilisation 

of 137 Cs in Sediments: The Effects of Sorption Kinetics and Reversibility.” Geochimica et 


Soldatini, G. F., R. Riffaldi, and R. Levi-Minzi. 1976. “Lead adsorption by Soils.” Water, Air 

and Soil Pollution. 6:111-128. 

Sposito, G. 1984. The Surface Chemistry of Soils. Oxford University Press, New York, 

New York. 

Stevenson, F. J., and A. Fitch. 1986. “Chemistry of Complexation Metal Ions with Soil Solution 

Organics.” In Interactions of Soil Minerals with Natural Organics and Microbes, 

P. M. Huang and M. Schnitzer (eds.), SSSA Special Publication No. 17, Soil Science Society 

of America, Inc., Madison, Wisconsin. 

Stollenwerk, K. G., and D. B. Grove. 1985. “Adsorption and Desorption of Hexavalent 

Chromium in an Alluvial Aquifer Near Telluride, Colorado.” Journal of Environmental 

Quality, 14:150-155. 

Strenge, D. L., and S. R. Peterson. 1989. Chemical Databases for the Multimedia 

Environmental Pollutant Assessment System. PNL-7145, Pacific Northwest Laboratory, 

Richland, Washington. 

Stumm, W., and J. J. Morgan. 1981. Aquatic Chemistry. An Introduction Emphasizing 

Chemical Equilibria in Natural Waters. John Wiley and Sons, New York, New York. 

6.20

Szalay, A. 1964. “Cation Exchange Properties of Humic Acids and their Importance in the 

2+ 

Geochemical Enrichment of UO2 and Other Cations.” Geochimica et Cosmochimica Acta, 

28:1605-1614. 

Tait, C. D., S. A. Ekberg, P. D. Palmer, and D. E. Morris. 1995. Plutonium Carbonate 

Speciation Changes as Measured in Dilute Solutions with Photoacoustic Spectroscopy. 

LA-12886-MS, Los Alamos National Laboratory, Los Alamos, New Mexico. 

Tamura, T. 1972. “Sorption Phenomena Significant in Radioactive Waste Disposal.” In 

Underground Waste Management and Environmental Implications, American Association of 

Petroleum Geology Memoirs, 18:318-330. 

Tanner, A. B., 1980. “Radon Migration in the Ground: A Supplementary Review.” In Natural 

Radiation Environment III. Volumes 1 and 2. Proceedings of Symposium Held at Houston, 

Texas, April 23-28. 1978, T. F. Gesell and W. M. Lowder (eds.), pp. 5-56, U.S. Department 

of Energy CONF 780422, National Technical Information Service, Springfield, Virginia. 

Taylor A. W. 1968. “Strontium Retention in Acid Soils of the North Carolina Coastal Plain.” 

Soil Science, 106:440-447. 

Ter Haar G. L., R. B. Holtzman, and H. F. Lucas. 1967. “Lead and Lead-210 in Rainwater.” 

Nature, 216:353-354. 

Thibault, D. H., M. I. Sheppard, and P. A. Smith. 1990. A Critical Compilation and Review of 

Default Soil Solid/Liquid Partition Coefficients, K d, for Use in Environmental Assessments. 

AECL-10125, Whiteshell Nuclear Research Establishment, Pinawa, Manitoba, Canada. 

Thompson, H. A., G. A. Parks, and G. E. Brown, Jr. 1998. “Structure and Composition of 

Uranium VI Sorption Complexes at the Kaolinite-Water Interface.” In Adsorption of Metals by 

Geomedia. Variables, Mechanisms, and Model Applications, E. A. Jenne (ed.), pp. 349-370, 

Academic Press, San Diego, California. 

Ticknor, K. V. 1993. “Actinide Sorption by Fracture-Filling Minerals.” Radiochimica Acta, 

60:33-42. 

Tripathi, V. S. 1984. Uranium(VI) Transport Modeling: Geochemical Data and Submodels. 

Ph.D. Dissertation, Stanford University, Stanford, California. 

Turner, G. D., J. M. Zachara, J. P. McKinley, and S. C. Smith. 1996. “Surface-Charge 

2+ 

Properties and UO2 Adsorption of a Subsurface Smectite.” Geochimica et Cosmochimica 

Acta, 60(18):3399-3414. 

6.21

UNSCEAR (United Nations Scientific Committee on the Effects of Atomic Radiation). 1982. 

Ionizing Radiation: Sources and Biological Effects. UNIPUB No. E.82.IX.8, 06300P, 

UNIPUB, New York, New York. 

Van Dalen, A., F. DeWitte, and J. Wikstra. 1975. Distribution Coefficients for Some 

Radionuclides Between Saline Water and Clays, Sandstones and Other Samples from Dutch 

Subsoil. Report 75-109, Reactor Centrum, Nederland. 

Waite, T. D., J. A. Davis, T. E. Payne, G. A. Waychunas, and N. Xu. 1994. “Uranium(VI) 

Adsorption to Ferrihydrite: Application of a Surface Complexation Model.” Geochimica et 


Waite, T. D., T. E. Payne, J. A. Davis, and K. Sekine. 1992. Alligators Rivers Analogue Project. 

Final Report Volume 13. Uranium Sorption. ISBN 0-642-599394 (DOE/HMIP/RR/92/0823, 

SKI TR 92:20-13. 

Wanner, H., and I. Forest (eds.). 1992. Chemical Thermodynamics Series, Volume 1: Chemical 

Thermodynamics of Uranium. North-Holland, Elsevier Science Publishing Company, Inc., 

New York, New York. 

Wang, W-Z., M. L. Brusseau, and J. F. Artiola. 1998. “Nonequilibrium and Nonlinear Sorption 

during Transport of Cadmium, Nickel, and Strontium Through Subsurface Soils.” In 

Adsorption of Metals by Geomedia. Variables, Mechanisms, and Model Applications, 

E. A. Jenne (ed.), pp. 427-443, Academic Press, San Diego, California. 

Weast, R. C., and M. J. Astle (eds.). 1980. CRC Handbook of Chemistry and Physics. CRC 

Press, Inc., Boca Raton, Florida. 

Weber, W. J., and H. S. Poselt. 1974. “Equilibrium Models and Precipitation Reactions for 

Cadmium(II).” In Chemical Oceanography 2nd Edition. J. P. Riley and G. Skirrow (eds.), 

pp. 311-356, Academic Press, New York, New York. 

White, A. F., and M. F. Hochella, Jr. 1989. “Electron Transfer Mechanism Associated with the 

Surface Dissolution and Oxidation of Magnetite and Ilmenite.” In Water-Rock Interaction 

WRI-6, D. L. Miles (ed.), p.765-768. A. A. Balkema, Rotterdam, Holland. 

Wiklander, L. 1964. “Uptake, Adsorption and Leaching of Radiostrontium in a Lysimeter 

Experiment.” Soil Science, 97:168-172. 

Yamaguchi, T., Y. Sakamoto, and T. Ohnuki. 1994. “Effect of the Complexation on Solubility 

of Pu(IV) in Aqueous Carbonate System.” Radiochimica Acta, 66/67:9-14. 

6.22

Yariv, S., and H. Cross. 1979. Geochemistry of Colloid Systems. Springer-Verlag, New York, 

New York. 

Yeh, G., and V. S. Tripathi. 1991. “A Model for Simulating Transport of Reactive Multispecies 

Components: Model Development and Demonstration.” Water Resources Research, 

27:3075-3094. 

Yong, R. N., and E. M. MacDonald. 1998. “Influence of pH, Metal Concentration, and Soil 

Component Removal on Retention of Pb and Cu by an Illitic Soil.” In Adsorption of Metals 

by Geomedia. Variables, Mechanisms, and Model Applications, E. A. Jenne (ed.), 


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 


6.23

APPENDICES

APPENDIX A 

Acronyms, Abbreviations, Symbols, and Notation

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 Element Symbol Element Symbol Element 

Ac Actinium 

Ag Silver 

Al Aluminum 

Am Americium 

Ar Argon 

As Arsenic 

At Astatine 

Au Gold 

B Boron 

Ba Barium 

Be Beryllium 

Bi Bismuth 

Bk Berkelium 

Br Bromine 

C Carbon 

Ca Calcium 

Cb Columbium 

Cd Cadmium 

Ce Cerium 

Cf Californium 

Cl Chlorine 

Cm Curium 

Co Cobalt 

Cr Chromium 

Cs Cesium 

Cu Copper 

Dy Dysprosium 

Er Erbium 

Es Einsteinium 

Eu Europium 

F Fluorine 

Fe Iron 

Fm Fermium 

Fr Francium 

Ga Gallium 

Gd Gadolinium 

Ge Germanium 

H Hydrogen 

He Helium 

Hf Hafnium 

Hg Mercury 

Ho Holmium 

I Iodine 

In Indium 

Ir Iridium 

K Potassium 

Kr Krypton 

La Lanthanum 

Li Lithium 

Lu Lutetium 

Lw Lawrencium 

Md Mendelevium 

Mg Magnesium 

Mn Manganese 

Mo Molybdenum 

N Nitrogen 

Na Sodium 

Nb Niobium 

Nd Neodymium 

Ne Neon 

Ni Nickel 

No Nobelium 

Np Neptunium 

O Oxygen 

Os Osmium 

P Phosphorus 

Pa Protactinium 

Pb Lead 

Pd Palladium 

Pm Promethium 

A.4 

Po Polonium 

Pr Praseodymium 

Pt Platinum 

Pu Plutonium 

Ra Radium 

Rb Rubidium 

Re Rhenium 

Rh Rhodium 

Rn Radon 

Ru Ruthenium 

S Sulfur 

Sb Antimony 

Sc Scandium 

Se Selenium 

Si Silicon 

Sm Samarium 

Sn Tin 

Sr Strontium 

Ta Tantalum 

Tb Terbium 

Tc Technetium 

Te Tellurium 

Th Thorium 

Ti Titanium 

Tl Thallium 

Tm Thulium 

U Uranium 

V Vanadium 

W Tungsten 

W Wolfram 

Xe Xenon 

Y Yttrium 

Yb Ytterbium 

Zn Zinc 

Zr Zirconium

A.3.0 List of Symbols and Notation 

Db Porous media bulk density (mass/length3 ) 

Å Angstrom, 10-10 meters 

ads Adsorption or adsorbed 

Ai Concentration of adsorbate (or species) I on the solid phase at equilibrium 

am Amorphous 

aq Aqueous 

CEC Cation exchange capacity 

Ci Curie 

d Day 

dpm Disintegrations per minute 

e - 

Free electron 

Eh Redox potential of an aqueous system relative to the standard hydrogen electrode 

F Faraday constant, 23,060.9 cal/V·mol 

g Gram 

3 

H Tritium 

h Hour 

I Ionic strength 

IAP Ion activity product 

IEP Isoelectric point 

Kd Concentration-based partition (or distribution) coefficient 

Kr,298 Equilibrium constant at 298 K 

Kr,T Equilibrium constant at temperature T 

l Liter 

M Molar 

m Meter 

mCi Millicurie, 10 -3 Curies 

meq Milliequivalent 

mi Mile 

ml Milliliter 

mol Mole 

mV Millivolt 

N Constant in the Freundlich isotherm model 

n Total porosity 

ne Effective porosity 

pCi Picocurie, 10 -12 Curies 

pE Negative common logarithm of the free-electron activity 

pH Negative logarithm of the hydrogen ion activity 

pHzpc pH for zero point of charge 

ppm Parts per million 

A.5

R Ideal gas constant, 1.9872 cal/mol·K 

Rf Retardation factor 

s Solid phase species 

sec Second 

SI Saturation index, as defined by log (IAP/Kr,T) T Absolute temperature, usually in Kelvin unless otherwise specified 

t Time 

t½ Half life 

TDS Total dissolved solids 

TU Tritium unit which is equivalent to 1 atom of 3H (tritium) per 1018 atoms 

of 1H (protium) 

vc Velocity of contaminant through a control volume 

vp Velocity of the water through a control volume 

y Year 

Z Valence state 

z Charge of ion 

{ } Activity 

[ ] Concentration 

A.6

APPENDIX B 

Definitions

Appendix B 

Definitions 

Adsorption - partitioning of a dissolved species onto a solid surface. 

Adsorption Edge - the pH range where solute adsorption sharply changes from ~10% to ~90%. 

Actinon - name occasionally used, especially in older documents, to refer to 219 Rn 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, (. 

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 4 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 222 Rn from 226 Ra). 

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 

where X designates an exchange surface site. 

CaX(s) + 90 Sr 2% (aq) = 90 SrX(s) + Ca 2% (aq) 

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 µm (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

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 U 4+ can be written as 

U 4+ + H 2O = UOH 3+ + H + . 

Hydrolytic Species - an aqueous species formed from a hydrolysis reaction. 

Ionic Potential - ratio (z/r) of the formal charge (z) to the ionic radius (r) of an ion. 

Isoelectric Point (iep) - pH at which a mineral’s surface has a net surface charge of zero. More 

precisely, it is the pH at which the particle is electrokinetically uncharged. 

Lignite - a coal that is intermediate in coalification between peat and subbituminous coal. 

Marl - an earthy substance containing 35-65% clay and 65-35% carbonate formed under marine 

or freshwater conditions 

Mass Transfer - transfer of mass between 2 or more phases that includes an aqueous solution, 

such as the mass change resulting from the precipitation of a mineral or adsorption of a metal 

on a mineral surface. 

Mass Transport - time-dependent movement of 1 or more solutes during fluid flow. 

Mire - a small piece of marshy, swampy, or boggy ground. 

Model Validation - integrated test of the accuracy with which a geochemical model and its 

thermodynamic database simulate actual chemical processes. 

Monomeric Species - an aqueous species containing only 1 center cation (as compared to a 

polymeric species). 

Near Field - the portion of a contaminant plume that is near the point source and whose chemical 

composition is significantly different from that of the uncontaminated portion of the aquifer. 

Peat - an unconsolidated deposit of semicarbonized plant remains in a water saturated 

environment. 

B.4

Polynuclear Species - an aqueous species containing more than 1 central cation moiety, e.g., 

- 4+ 

(UO2) 2CO3(OH) 3 and Pb4(OH) 4 . 

Protium (H) - stable isotope 1 H of hydrogen. 

Retrograde Solubility - solubility that decreases with increasing temperature, such as those of 

calcite (CaCO 3) 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 220 Rn which forms 

from the decay of thorium. 

Tritium (T) - radioactive isotope 3 H of hydrogen. 

Tritium Units - units sometimes used to report tritium concentrations. A tritium unit (TU) is 

equivalent to 1 atom of 3 H (tritium) per 10 18 atoms of 1 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

C.1.0 Background 

Appendix C 

Partition Coefficients For Cadmium 

Cadmium K d 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 K d values and were conducted in systems consisting of 

C Natural soils (as opposed to pure mineral phases) 

C Low ionic strength solutions (

Table C.1. Descriptive statistics of the cadmium K d data set for soils. 

Cadmium 

K d 

(ml/g) 

Clay 

Content 

(wt.%) 

pH CEC 

(meq/100g) 

C.3 

TOC 

(mg/l) 

Cd Conc. 

(mg/l) 

Fe Oxides 

(wt.%) 

Mean 226.7 14.2 5.88 21 5.5 3.67 1.32 

Standard 

Error 

44.5 1.7 0.09 3 0.85 0.48 0.53 

Median 121.8 10.24 5.83 23 2.0 0.01 0.38 

Mode 80.0 6 6.8 2 0.4 0.01 0.19 

Std. Dev 586.6 13.5 1.16 15 6.8 6.27 2.12 

Sample 

Variance 

344086 182 1.34 245 45.9 39.4 4.51 

Range 4359 86.2 6.20 58 32.4 34.9 8.28 

Minimum 0.50 .9 3 2 0.2 0.01 0.01 

Maximum 4360 87.1 9.2 60 32.6 35 8.29 

No. 

Samples 

174 62 174 22 63 172 16 

C.2.0 Approach and Regression Models 

C.2.1 Correlations with Cadmium K d Values 

Linear regression analyses were conducted between the ancillary parameters and cadmium K d 

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(K d). 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 K d 

values: total organic carbon and pH, clay content and pH, total organic carbon and iron-oxides, 

and pH and CEC.

Cadmium 

K d 

Table C.2. Correlation coefficients (r) of the cadmium K d data set for soils. 

Cadmium 

K d 

log (K d) 0.69 1 

1 

log (K d) Clay 

Content 

Clay Conc. -0.04 0.03 1 

pH 0.50 0.75 0.06 1 

CEC 0.40 0.41 0.62 0.35 1 

C.4 

pH CEC TOC Cd Conc. 

TOC 0.20 0.06 0.13 -0.39 0.27 1 

Cd Conc. -0.02 -0.10 -0.39 0.22 -0.03 -0.09 1 

Fe Oxide 

Conc. 

0.18 0.11 -0.06 0.16 0.19 0.18 0.01 

C.2.2 Cadmium K d Values as a Function of pH 

The cadmium K d values plotted as a function of pH are presented in Figure C.1. A large amount 

of scatter exists in these data. At any given pH, the range of K d 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 K d values was not fruitful. 

The look-up table (Table C.3) for cadmium K d values was categorized by pH. The regression 

equation for the line presented in Figure C.1 is: 

Cd K d = -0.54 + 0.45(pH). (C.1) 

The minimum and maximum values were estimated based on the scatter of data points observed in 

Figure C.1.

Figure C.1. Relation between cadmium K d values and pH in soils. 

Table C.3. Look-up table for estimated range of K d 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 (

C.3.0 Data Set for Soils 

Table C.4 lists the available K d values for cadmium identified for experiments conducted with only 

soils. The K d values are listed with ancillary parameters that included clay content, pH, CEC, 

TOC, solution cadmium concentrations, and iron-oxide concentrations 

Cd K d 

(ml/g) 

Clay 

Cont. 

(wt%) 

pH CEC 

(meq/ 

100 g) 

Table C.4. Cadmium K d data set for soils. 

TOC 

(wt%) 

[Cd] 

(mg/l) 

Fe 

Oxides 

(wt.%) 

52.5 54.7 4.8 30.2 1.54 1 0.33 0.005 M 

CaNO 3 

288.4 8.3 5.7 2 0.61 1 0.1 0.005 M 

CaNO 3 

13.9 51.2 5.4 2.4 0.26 1 0.08 0.005 M 

CaNO 3 

186.6 0.9 5.9 22.54 6.62 1 1.68 0.005 M 

CaNO 3 

52.7 17.6 3.9 26.9 11.6 1 1.19 0.005 M 

CaNO 3 

91.2 28.2 6 11 1.67 1 0.19 0.005 M 

CaNO 3 

28.8 2.8 6.9 4.1 0.21 1 0.06 0.005 M 

CaNO 3 

97.9 6.2 6.6 8.6 0.83 1 0.3 0.005 M 

CaNO 3 

5.5 3.8 4.3 2.7 1.98 1 0 0.005 M 

CaNO 3 

755.1 23.9 7.6 48.1 4.39 1 0.19 0.005 M 

CaNO 3 

C.6 

Solution Soil 

Identification 

Comments Ref. a 

Alligator Ap Converted 

Freund. to K d 

Using 1ppm 

Cecil Ap Converted 


Using 1ppm 

Cecil B Converted 


Using 1ppm 

Kula Ap1 Converted 


Using 1ppm 

Lafitte Ap Converted 


Using 1ppm 

Molokai Ap Converted 


Using 1ppm 

Norwood Ap Converted 


Using 1ppm 

Olivier Ap Converted 


Using 1ppm 

Spodisol Converted 


Using 1ppm 

Webster Ap Converted 


Using 1ppm 

1 

1 

1 

1 

1 

1 

1 

1 

1 

1

Cd K d 

(ml/g) 

Clay 

Cont. 

(wt%) 

pH CEC 

(meq/ 

100 g) 

TOC 

(wt%) 

[Cd] 

(mg/l) 

Fe 

Oxides 

(wt.%) 

14.4 2.8 5.3 2 2.03 1 0.42 0.005 M 

CaNO 3 

C.7 

Solution Soil 


87.1 8.4 60 1.44 1 1.07 Water Vertic 

Torrifluvent 


Windsor Ap Converted 


Using 1ppm 

Converted 


Using 1ppm 

33.88 5.2 33.8 32.6 1 Water Organic Converted 


Using 1ppm 

20.42 5.8 23.8 3 1 8.29 Water Boomer, Ultic 

Haploxeralf 

10.47 6 25 3.2 1 1.07 Water UlticPalexeral 

f 

80 8.2 8.2 0.21 35 0.01 M 

NaCl 

200 7.8 15.4 0.83 25 0.01 M 

NaCl 

133.3 8.3 18.9 0.23 30 0.01 M 

NaCl 

181.8 7.6 31.8 0.79 25 0.01 M 

NaCl 

266.7 7.9 37 0.86 15 0.01 M 

NaCl 

8 8 3.7 1.6 11.2 0.01 M 

NaNO 3 

17 8 4.8 1.6 11.2 0.01 M 

NaNO 3 

32 8 5.3 1.6 11.2 0.01 M 

NaNO 3 

64 8 6 1.6 11.2 0.01 M 

NaNO 3 

92 8 6.2 1.6 11.2 0.01 M 

NaNO 3 

110 8 6.8 1.6 11.2 0.01 M 

NaNO 3 

250 8 7.3 1.6 11.2 0.01 M 

NaNO 3 

Converted 


Using 1ppm 

Converted 


Using 1ppm 

Gevulot Calc. Fig 1. 3 

Bet Yizhaq Calc. Fig 1. 3 

Gilat Calc. Fig 1. 3 

Maaban 

Michael 

1 

2 

2 

2 

2 

Calc. Fig 1. 3 

Hahoterim Calc. Fig 1. 3 

Downer 

Loamy Sand 

Downer 

Loamy Sand 

Downer 

Loamy Sand 

Downer 

Loamy Sand 

Downer 

Loamy Sand 

Downer 

Loamy Sand 

Downer 

Loamy Sand 

4 

4 

4 

4 

4 

4 

4

Cd K d 

(ml/g) 

Clay 

Cont. 

(wt%) 

pH CEC 

(meq/ 

100 g) 

TOC 

(wt%) 

[Cd] 

(mg/l) 

Fe 

Oxides 

(wt.%) 

580 8 8.5 1.6 11.2 0.01 M 

NaNO 3 

0.5 6 3.1 0.4 11.2 0.01 M 

NaNO 3 

3.3 6 3.8 0.4 11.2 0.01 M 

NaNO 3 

7.5 6 4.5 0.4 11.2 0.01 M 

NaNO 3 

10 6 5.5 0.4 11.2 0.01 M 

NaNO 3 

34 6 6.1 0.4 11.2 0.01 M 

NaNO 3 

45 6 6.8 0.4 11.2 0.01 M 

NaNO 3 

80 6 7.5 0.4 11.2 0.01 M 

NaNO 3 

150 6 8 0.4 11.2 0.01 M 

NaNO 3 

420 6 8.4 0.4 11.2 0.01 M 

NaNO 3 

900 6 9.1 0.4 11.2 0.01 M 

NaNO 3 

2.1 13 3 16.8 11.2 0.01 M 

NaNO 3 

10 13 3.7 16.8 11.2 0.01 M 

NaNO 3 

30 13 4.2 16.8 11.2 0.01 M 

NaNO 3 

57 13 4.6 16.8 11.2 0.01 M 

NaNO 3 

C.8 

Solution Soil 


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 4 




4 

4 

4 

4 

4 

4 

4 

4 

4 

4 

4

Cd K d 

(ml/g) 

Clay 

Cont. 

(wt%) 

pH CEC 

(meq/ 

100 g) 

TOC 

(wt%) 

[Cd] 

(mg/l) 

Fe 

Oxides 

(wt.%) 

101 13 5 16.8 11.2 0.01 M 

NaNO 3 

195 13 5.2 16.8 11.2 0.01 M 

NaNO 3 

420 13 5.8 16.8 11.2 0.01 M 

NaNO 3 

1,200 13 6.2 16.8 11.2 0.01 M 

NaNO 3 

4,000 13 6.8 16.8 11.2 0.01 M 

NaNO 3 

1.2 16 3.3 9.8 11.2 0.01 M 

NaNO 3 

7.1 16 4.1 9.8 11.2 0.01 M 

NaNO 3 

27 16 4.8 9.8 11.2 0.01 M 

NaNO 3 

53 16 5.1 9.8 11.2 0.01 M 

NaNO 3 

170 16 5.6 9.8 11.2 0.01 M 

NaNO 3 

300 16 6.1 9.8 11.2 0.01 M 

NaNO 3 

390 16 6.2 9.8 11.2 0.01 M 

NaNO 3 

910 16 6.5 9.8 11.2 0.01 M 

NaNO 3 

1,070 16 6.8 9.8 11.2 0.01 M 

NaNO 3 

43 10 4.8 2.4 11.2 0.01 M 

NaNO 3 

67 10 5.7 2.4 11.2 0.01 M 

NaNO 3 

130 10 6.3 2.4 11.2 0.01 M 

NaNO 3 

C.9 

Solution Soil 








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 


Delaware 

River 


Delaware 

River 

4 

4 

4 

4 

4 

4 

4 

4 

4 

4 

4 

4

Cd K d 

(ml/g) 

Clay 

Cont. 

(wt%) 

pH CEC 

(meq/ 

100 g) 

TOC 

(wt%) 

[Cd] 

(mg/l) 

Fe 

Oxides 

(wt.%) 

150 10 6.7 2.4 11.2 0.01 M 

NaNO 3 

370 10 7.3 2.4 11.2 0.01 M 

NaNO 3 

880 10 8 2.4 11.2 0.01 M 

NaNO 3 

1,950 10 9.2 2.4 11.2 0.01 M 

NaNO 3 

1,000 12 8 1 3.7 Carbonate 

Groundwate 

r 

4,360 12.4 8 1 2.5 Carbonate 

Groundwate 

r 

536.8 25.2 6.8 27.5 0.01 M 

NaCl 

440 25.2 6.8 27.5 0.01 M 

NaCl 

9 4.3 0.01 0.001M 

CaCl 2 

23.4 4.3 0.01 0.001M 

CaCl 2 

15.8 4.4 0.01 0.001M 

CaCl 2 

11.3 4.5 0.01 0.001M 

CaCl 2 

31.2 4.5 0.01 0.001M 

CaCl 2 

32.5 4.5 0.01 0.001M 

CaCl 2 

23 4.5 0.01 0.001M 

CaCl 2 

17.1 4.7 0.01 0.001M 

CaCl 2 

C.10 

Solution Soil 



Delaware 

River 


Delaware 

River 


Delaware 

River 


Delaware 

River 


Interbed pH of 

Groundwater 

Alluvium pH of 

Groundwater 

Soil A Desorption 6 

Soil A Desorption 6 

Agricultural 

Danish Soil 

Agricultural 

Danish Soil 

Agricultural 

Danish Soil 

Agricultural 

Danish Soil 

Agricultural 

Danish Soil 

Agricultural 

Danish Soil 

Agricultural 

Danish Soil 

Agricultural 

Danish Soil 

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 

4 

4 

4 

4 

5 

5 

7 

7 

7 

7 

7 

7 

7 

7

Cd K d 

(ml/g) 

Clay 

Cont. 

(wt%) 

pH CEC 

(meq/ 

100 g) 

TOC 

(wt%) 

[Cd] 

(mg/l) 

Fe 

Oxides 

(wt.%) 

13.1 4.8 0.01 0.001M 

CaCl 2 

24.9 4.6 0.01 0.001M 

CaCl 2 

26.8 4.7 0.01 0.001M 

CaCl 2 

36.2 4.7 0.01 0.001M 

CaCl 2 

32.9 4.7 0.01 0.001M 

CaCl 2 

37.2 4.7 0.01 0.001M 

CaCl 2 

29.2 4.8 0.01 0.001M 

CaCl 2 

28.3 4.8 0.01 0.001M 

CaCl 2 

22.6 4.9 0.01 0.001M 

CaCl 2 

37.4 4.9 0.01 0.001M 

CaCl 2 

40.9 4.9 0.01 0.001M 

CaCl 2 

63.5 4.7 0.01 0.001M 

CaCl 2 

25.2 5.4 0.01 0.001M 

CaCl 2 

29.9 5.3 0.01 0.001M 

CaCl 2 

33.7 5.2 0.01 0.001M 

CaCl 2 

44.3 5.1 0.01 0.001M 

CaCl 2 

42.8 5.1 0.01 0.001M 

CaCl 2 

53.5 5 0.01 0.001M 

CaCl 2 

56.2 4.9 0.01 0.001M 

CaCl 2 

C.11 

Solution 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 

Agricultural 

Danish Soil 


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 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7

Cd K d 

(ml/g) 

Clay 

Cont. 

(wt%) 

pH CEC 

(meq/ 

100 g) 

TOC 

(wt%) 

[Cd] 

(mg/l) 

Fe 

Oxides 

(wt.%) 

68.7 5 0.01 0.001M 

CaCl 2 

82.3 5.1 0.01 0.001M 

CaCl 2 

75.7 5 0.01 0.001M 

CaCl 2 

95.2 4.8 0.01 0.001M 

CaCl 2 

103 4.8 0.01 0.001M 

CaCl 2 

160 4.8 0.01 0.001M 

CaCl 2 

43.3 5.4 0.01 0.001M 

CaCl 2 

55.2 5.4 0.01 0.001M 

CaCl 2 

52.2 5.3 0.01 0.001M 

CaCl 2 

40.3 5.6 0.01 0.001M 

CaCl 2 

56.1 5.5 0.01 0.001M 

CaCl 2 

67.5 5.5 0.01 0.001M 

CaCl 2 

102.9 5.4 0.01 0.001M 

CaCl 2 

164.4 5.5 0.01 0.001M 

CaCl 2 

163.8 5.3 0.01 0.001M 

CaCl 2 

202.1 5.2 0.01 0.001M 

CaCl 2 

172.4 5.2 0.01 0.001M 

CaCl 2 

149 5.2 0.01 0.001M 

CaCl 2 

72.8 5.6 0.01 0.001M 

CaCl 2 

C.12 

Solution 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 

Agricultural 

Danish Soil 


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 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7

Cd K d 

(ml/g) 

Clay 

Cont. 

(wt%) 

pH CEC 

(meq/ 

100 g) 

TOC 

(wt%) 

[Cd] 

(mg/l) 

Fe 

Oxides 

(wt.%) 

81.6 5.7 0.01 0.001M 

CaCl 2 

90 5.7 0.01 0.001M 

CaCl 2 

94.3 5.6 0.01 0.001M 

CaCl 2 

48.1 6.2 0.01 0.001M 

CaCl 2 

56.5 6.4 0.01 0.001M 

CaCl 2 

81 6.5 0.01 0.001M 

CaCl 2 

122.3 6.4 0.01 0.001M 

CaCl 2 

121.4 6.2 0.01 0.001M 

CaCl 2 

101.5 6 0.01 0.001M 

CaCl 2 

99.3 6 0.01 0.001M 

CaCl 2 

107.8 6 0.01 0.001M 

CaCl 2 

219.5 6.2 0.01 0.001M 

CaCl 2 

179.2 6.2 0.01 0.001M 

CaCl 2 

177 6.1 0.01 0.001M 

CaCl 2 

360.4 6 0.01 0.001M 

CaCl 2 

305.2 6 0.01 0.001M 

CaCl 2 

236.8 5.9 0.01 0.001M 

CaCl 2 

186.3 5.9 0.01 0.001M 

CaCl 2 

174.8 5.8 0.01 0.001M 

CaCl 2 

C.13 

Solution 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 

Agricultural 

Danish Soil 


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 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7

Cd K d 

(ml/g) 

Clay 

Cont. 

(wt%) 

pH CEC 

(meq/ 

100 g) 

TOC 

(wt%) 

[Cd] 

(mg/l) 

Fe 

Oxides 

(wt.%) 

138.7 5.8 0.01 0.001M 

CaCl 2 

132.5 5.7 0.01 0.001M 

CaCl 2 

375.6 5.9 0.01 0.001M 

CaCl 2 

403.3 5.8 0.01 0.001M 

CaCl 2 

510.8 5.8 0.01 0.001M 

CaCl 2 

225.9 5.7 0.01 0.001M 

CaCl 2 

227.3 5.7 0.01 0.001M 

CaCl 2 

248 5.7 0.01 0.001M 

CaCl 2 

253.1 5.6 0.01 0.001M 

CaCl 2 

277.2 5.6 0.01 0.001M 

CaCl 2 

240.7 6.4 0.01 0.001M 

CaCl 2 

227.8 6.5 0.01 0.001M 

CaCl 2 

281.1 6.6 0.01 0.001M 

CaCl 2 

551.2 6.2 0.01 0.001M 

CaCl 2 

519.8 6.2 0.01 0.001M 

CaCl 2 

418.7 6.2 0.01 0.001M 

CaCl 2 

353.7 6.2 0.01 0.001M 

CaCl 2 

400.8 6.4 0.01 0.001M 

CaCl 2 

609.2 6.3 0.01 0.001M 

CaCl 2 

C.14 

Solution 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 

Agricultural 

Danish Soil 


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 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7

Cd K d 

(ml/g) 

Clay 

Cont. 

(wt%) 

pH CEC 

(meq/ 

100 g) 

TOC 

(wt%) 

[Cd] 

(mg/l) 

Fe 

Oxides 

(wt.%) 

545.7 6.3 0.01 0.001M 

CaCl 2 

515.9 6.4 0.01 0.001M 

CaCl 2 

545.7 6.4 0.01 0.001M 

CaCl 2 

760.9 6.4 0.01 0.001M 

CaCl 2 

665.7 6.5 0.01 0.001M 

CaCl 2 

503.2 6.5 0.01 0.001M 

CaCl 2 

515.2 7 0.01 0.001M 

CaCl 2 

488.9 6.9 0.01 0.001M 

CaCl 2 

481 6.9 0.01 0.001M 

CaCl 2 

461.6 6.9 0.01 0.001M 

CaCl 2 

1,151 6.5 0.01 0.001M 

CaCl 2 

868.7 6.6 0.01 0.001M 

CaCl 2 

637.2 6.7 0.01 0.001M 

CaCl 2 

970.9 6.7 0.01 0.001M 

CaCl 2 

950.5 6.8 0.01 0.001M 

CaCl 2 

886.2 6.9 0.01 0.001M 

CaCl 2 

1,106 6.9 0.01 0.001M 

CaCl 2 

970.9 7 0.01 0.001M 

CaCl 2 

2,248 7.1 0.01 0.001M 

CaCl 2 

C.15 

Solution 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 

Agricultural 

Danish Soil 


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 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7 

7

Cd K d 

(ml/g) 

Clay 

Cont. 

(wt%) 

pH CEC 

(meq/ 

100 g) 

TOC 

(wt%) 

[Cd] 

(mg/l) 

Fe 

Oxides 

(wt.%) 

1,909 7.2 0.01 0.001M 

CaCl 2 

1,411 7.3 0.01 0.001M 

CaCl 2 

1,383 7.4 0.01 0.001M 

CaCl 2 

2,337 7.5 0.01 0.001M 

CaCl 2 

C.16 

Solution Soil 


Agricultural 

Danish Soil 

Agricultural 

Danish Soil 

Agricultural 

Danish Soil 

Agricultural 

Danish Soil 


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 

a 1 = Buchter et al., 1989; 2 = Garcia-Miragaya, 1980; 3 = Navrot et al., 1978; 4 = Allen et al., 1995; 5 = Del Debbio, 

1991; 6 = Madrid et al., 1992; 7 = Anderson and Christensen , 1988 

C.4.0 References 

Allen, G. E., Y. Chen, Y, Li, and C. P. Huang. 1995. “Soil Partition Coefficients for Cd by 

Column Desorption and Comparison to Batch Adsorption Measurements.” Environmental 

Science and Technology, 29:1887-1891. 

Anderson, P. R., and T. H. Christensen. 1988. “Distribution Coefficients of Cd, Co, Ni, and Zn 

in Soils.” Journal of Soil Science, 39:15-22. 

Buchter, B., B. Davidoff, M. C. Amacher, C. Hinz, I. K. Iskandar, and H. M. Selim. 1989. 

“Correlation of Freundlich K d 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. 

7 

7 

7 

7

APPENDIX D 

Partition Coefficients For Cesium

D.1.0 Background 

Appendix D 

Partition Coefficients For Cesium 

Three generalized, simplifying assumptions were established for the selection of cesium K d 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: 

C 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. 

C Cesium adsorption onto mica-like minerals occurs much more readily than desorption. 

Thus, K d values, which are essentially always derived from adsorption studies, will greatly 

overestimate the degree to which cesium will desorb from these surfaces. 

C Cesium concentrations in groundwater plumes are low enough, less than approximately 

10 -7 M, such that cesium adsorption follows a linear isotherm. 


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 K d 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 K d values (not percent 




C Dissolved cesium concentrations less than 10 -7 M 


Two separate data sets were compiled. The first one (see Section D.3) included both soils and 

pure mineral phases. The lowest cesium K d 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 al., 

1986) (Table D.1). The value is unexplainably much less than most other cesium K d values 

present in the data set. The largest cesium K d values was 52,000 ml/g for a measurement made on 

a pure vermiculite solid phase (Tamura, 1972). The average cesium K d value was 2635 ± 

530 ml/g. 

Table D.1. Descriptive statistics of cesium K d data set including soil and pure mineral 

phases. [Data set is presented in Section D.3.] 

K d (ml/g) Clay 

(%) 

Mica 

(%) 

D.3 

pH CEC 

(meq/100 g) 

Surface Area 

(m 2 /g) 

Mean 2,635 30 5.5 7.4 30.4 141.3 

Standard Error 530 3.8 0.7 0.1 3.7 29.7 

Median 247 42 4 8.2 4.8 31.2 

Mode 40 42 4 8.2 1.8 17.7 

Standard Deviation 7055 15 4.4 1.7 37.4 230.4 

Sample Variance 49,781,885 226 20.0 2.8 1,396.9 53,106 

Range 51,999 38 13 7.8 129.9 638 

Minimum 0.6 4 2 2.4 0.00098 8 

Maximum 52,000 42 15 10.2 130 646 

No. Observations 177 15 41 139 103 60 

Confidence Level 

(95.0%) 

1,046.6 8.3 1.4 0.3 7.3 59.5

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 median 1 of the 57 K d 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 K d value for this data set. The range of K d 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. [Data set is presented 

in Section D.4.] 

1 The median is that value for which 50 percent of the observations, when arranged in order of 

magnitude, lie on each side. 

Cesium 

K d 

(ml/g) 

Clay 

(%) 

D.4 

Mica 

(%) 

pH CEC 

(meq/100g) 

Surface Area 

(m 2 /g) 

Mean 651 5 5.6 6.9 34 57.5 

Standard Error 188 0.6 0.6 0.3 8.9 13.4 

Median 89 5.0 4 6.7 20 60 

Mode 22 NA 4 4.0 60 70 

Standard Deviation 1423 1.0 4.3 1.9 29.5 44.6 

Sample Variance 2026182 1.0 18.4 3.6 870 1986 

Range 7602 2.0 13 7.8 57.4 123.4 

Minimum 7.1 7.1 2 2.4 2.6 6.6 

Maximum 7610 6.0 15 10.2 70.0 130 

No. Observations 57 3 45 55 11 11 

Confidence Level (95%) 378 2.5 1.29 0.5 19.8 30

The soil-only data set was frequently incomplete with regard to supporting data describing the 

experimental conditions under which the cesium K d values were measured (Table D.2). Quite 

often the properties of the solid phase or the dissolved cesium concentration used in the K d 

experiments were not reported. For instance, there were only 3 cesium K d values that had 

accompanying clay content data, 11 cesium K d values that had accompanying cation exchange 

data, and 11 cesium K d values that had accompanying surface area data (Table D.2). 

Consequently, it was not possible to evaluate adequately the relationship between cesium K d 

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 K d Values 

A matrix of the correlation coefficients for the parameters included in the data set containing K d 

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 K d was CEC (r = 0.52). Also significant was the correlation coefficient 

between cesium K d values and surface area (r = 0.42) and CEC and clay content (r = 0.64). The 

poor correlation between cesium aqueous concentration ([Cs] aq) and cesium K d 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 K d values as well as the concentrations after contact 

with soil (the denominator of the K d term). 

A matrix of the correlation coefficients for the parameters included in the data set containing K d 

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 K d 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 K d 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 K d value data set that 

included soils and pure mineral phases. [Data set is presented in 

Section D.3.] 

Cesium K d 

Cesium 

K d 

1.00 

Clay 

Content 

Clay Content 0.05 1.00 

Mica 0.29 0.00 1.00 

pH 0.10 -0.11 0.08 1.00 

Mica pH CEC Surface Area 

CEC 0.52 a 0.64 a NA 0.37 1.00 

Surface Area 0.42 a 

[Cs] aq -0.07 0.85 a 

0.35 NA -0.11 0.47 a 

D.6 

1.00 

0.29 0.13 -0.17 -0.15 

a Correlation coefficient is significant at the 5% level of significance (P # 0.05). 

Table D.4. Correlation coefficients (r) of the soil-only data set. [Data set is 

presented in Section D.4.] 

Cesium K d 

Cesium 

K d 

1.00 

Clay 

Content 

Clay Content -0.21 1.00 

Mica 0.27 0 1.00 

pH 0.11 0.4 0.07 1.00 

CEC -0.83 NA 0.99 1 

Surface Area -0.31 NA 0.99 1 

Mica pH CEC Surface Area 

0.05 1.00 

-0.03 0.37 1.00 

[Cs] aq 0.18 NA 0.09 -0.04 0.00 0 

1 Correlation coefficient is significant at >5% level of significance (P # 0.05).

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 and pH 

Akiba and Hashimoto (1990) showed a strong correlation between cesium K d values and the CEC 

of a large number of soils, minerals, and rock materials. The regression equation generated from 

their study was: 

log (Cs K d) = 1.2 + 1.0 log (CEC) (D.1) 

A similar regression analysis using the entire data set (mineral, rocks, and soils) is presented in 

Figure D.1. 

Figure D.1. Relation between cesium K d values and CEC. 

D.7

By transposing the CEC and cesium K d data into logarithms, the regression correlation slightly 

increases from 0.52 (Table D.3) to 0.60 (Figure D.1). However, a great amount of scatter in the 

data can still be seen in the logarithmic transposed data. For instance, at log(CEC) of 0.25, the 

cesium K d values range over 4 orders of magnitude. It is important to note that the entire cesium 

K d data set only varies 5 orders of magnitude. Thus, the correlation with CEC, although the 

strongest of all the independent variables examined, did not reduce greatly the variability of 

possible cesium K d values. 

D.2.3 CEC as a Function of Clay Content and pH 

Because CEC values are not always available to contaminant transport modelers, an attempt was 

made to use independent variables more commonly available in the regression analysis. Multiple 

regression analysis was conducted using clay content and pH as independent variables to predict 

CEC values (Figure D.2). Clay content was highly correlated to CEC (r = 0.64). Soil pH was 

not significantly correlated to either CEC or cesium K d values. 

Figure D.2. Relation between CEC and clay content. 

D.8

D.2.4 Cesium Adsorption onto Mica-Like Minerals 

Cesium adsorption onto mica-like minerals has long been recognized as a non-reversible reaction 

(Bruggenwert and Kamphorst, 1979; Comans et al., 1989; Cremers et al., 1988; Douglas, 1989; 

Evans et al., 1983; Francis and Brinkley, 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 K d model to describe adsorption is that the rate at which adsorption 

occurs is equal to the rate at which desorption occurs. This phenomena is referred to as an 

adsorption hysteresis. Cesium adsorption onto mica-like minerals is appreciably faster than its 

desorption. The reason for this is that the cesium ion fits perfectly into the hexagonal ring formed 

on the tetrahedral sheet in the crystallographic structure of mica-like clays. This perfect fit does 

not permit other cations that exist at much greater concentrations in nature to exchange the 

cesium from these sites. This can be demonstrated using the data of Tamura (1972) (Table D.5). 

He measured cesium K d values for mica, vermiculite, and kaolinite using a water and 0.1 M NaCl 

background solution. For mica, the K d value remained about the same for both solutions. For the 

vermiculite and kaolinite, the cesium K d 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 

K d 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 K d values of a mica (Fithian illite) by desorption and 

adsorption experiments. Portions of their data are presented in Table D.6. Cesium K d 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 

1.67x10 -12 M.] 

Mineral 

Phases 

CEC 

(meq/100 g) 

D.9 

K d in Water 

(ml/g) 

K d in 0.1 M NaCl 

(ml/g) 

Mica 20 26,000 28,600 

Vermiculite 127 52,000 2,700 

Kaolinite 11.2 2,500 94

Table D.6. Cesium K d values measured on mica (Fithian illite) via adsorption and 

desorption experiments. [Data are from Comans et al. (1989).] 

Experimental Conditions Adsorption 

Cesium K d 

D.10 

Desorption 

Cesium K d 

K-saturated Mica, 7x10 -6 M Cs 2,890 5,200 

K-saturated Mica, 2x10 -7 M Cs 9,000 11,300 

Ca-saturated Mica, 7x10 -6 M Cs 1,060 4,600 

Ca-saturated Mica, 2x10 -7 M Cs 600,000 1,050,000 

Essentially all K d values reported in the literature are measured using adsorption experiments. 

Thus, in the case of soils containing mica-like soils, using adsorption K d 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 K d 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 137 Cs concentration in the 

groundwaters beneath the Hanford Site in 1994 was 1.94 x 10 -13 M (or 2,310 pCi/l) 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 K d value describes a linear isotherm.

Table D.7. Approximate upper limits of linear range of adsorption isotherms on various 

solid phases. 

Upper Limit of 

Linear Range (M) 

Solid Phase Background 

Aqueous Phase 

D.11 

Reference 

1 x 10 -7 Itado Tuff Deionized Water Akida et al., 1989 

1 x10 -10 Sandstone Deionized Water Akida et al., 1989 

5 x 10 -5 Limestone Deionized Water Akida et al., 1989 

1 x 10 -10 Augite Andesite Deionized Water Akida et al., 1989 

5 x 10 -9 Olivine Basalt Deionized Water Akida et al., 1989 

1 x 10 -8 

5 x 10 -8 

5 x 10 -7 

1 x 10 -6 

1 x 10 -1 

eported 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. 

Figure D.3. K d values calculated from an overall literature 

Freundlich equation for cesium (Equation D.2). 

D.12

Table D.8. Freundlich equations identified in literature for cesium. 

a 1 b 1 Range of Solution Cs 

Concentration (M) 

D.13 

Experimental Ref. 2 

1.7 0.677 Water/Batcombe Sediment 1 

3,300 0.909 Water/Denchworth Sediment 1 

260 0.841 Water/Tedburn Sediment 1 

16 0.749 Water/Teigngrace Sediment 1 

12.2 0.745 1x10 -8 to 1x10 -12 Water/Batcombe Sediment 1 

6,070 0.899 1x10 -8 to 1x10 -12 Water/Denchworth Sediment 1 

1,290 0.849 1x10 -8 to 1x10 -12 

163 0.815 1x10 -8 to 1x10 -12 

1.23 0.657 1x10 -8 to 1x10 -12 

Water/Tedburn Sediment 1 

Water/Teigngrace Sediment 1 

CaCl 2/Batcombe Sediment 1 

0.63 0.659 CaCl 2/Batcombe Sediment 1 

427 0.814 1x10 -8 to 1x10 -12 

CaCl 2/Denchworth Sediment 1 

1.5 0.599 CaCl 2/Denchworth Sediment 1 

48.1 0.754 1x10 -8 to 1x10 -12 

CaCl 2/Tedburn Sediment 1 

17 0.739 CaCl 2/Tedburn Sediment 1 

5.22 0.702 1x10 -8 to 1x10 -12 

CaCl 2/Teigngrace Sediment 1 

4.4 0.716 CaCl 2/Teigngrace Sediment 1 

0.22 1.1 1x10 -9 to 1.5x10 -2 

0.017 0.53 1x10 -9 to 1.5x10 -2 

0.13 1 1x10 -9 to 1.5x10 -2 

0.048 0.67 1x10 -9 to 1.5x10 -2 

5.10x10 -4 

3.00x10 -3 

1.30x10 -5 

0.21 1x10 -9 to 1.5x10 -2 

0.48 1x10 -9 to 1.5x10 -2 

0.013 1x10 -9 to 1.5x10 -2 

Bentonite/Water 2 

Bentonite/Water 2 

Bentonite/Groundwater 2 

Bentonite/Groundwater 2 

Takadata Loam/Water 2 

Takadata Loam/Groundwater 2 

Hachinohe Loam/Water 2 

2.30x10 -5 0.38 1x10 -9 to 1.5x10 -2 Hachinohe Loam/Groundwater 2

a 1 b 1 Range of Solution Cs 

Concentration (M) 

D.14 

Experimental Ref. 2 

2.70x10 -4 0.546 1x10 -8 to 1x10 -2 Unwashed/Kaolinite/pH 2 3 



2.27x10 -3 0.586 1x10 -8 to 1x10 -2 Sodium/Kaolinite/pH 2 3 

5.04x10 -2 0.723 1x10 -8 to 1x10 -2 Sodium/Kaolinite/pH 4 3 

3.49x10 -2 0.703 1x10 -8 to 1x10 -2 Na/Kaolinite/pH 7 3 

0.235 0.821 1x10 -8 to 1x10 -2 Na/Kaolinite/pH 10 3 

3.03x10 -2 

0.804 1x10 -8 to 1x10 -2 

0.135 0.845 1x10 -8 to 1x10 -2 

0.247 0.881 1x10 -8 to 1x10 -2 

8.71x10 -3 

1.02x10 -4 

1.05x10 -2 

3.17x10 -2 

0.694 1x10 -8 to 1x10 -2 

0.503 1x10 -8 to 1x10 -2 

0.709 1x10 -8 to 1x10 -2 

0.755 1x10 -8 to 1x10 -2 

0.224 0.815 1x10 -8 to 1x10 -2 

0.241 0.839 1x10 -8 to 1x10 -2 

0.481 0.897 1x10 -8 to 1x10 -2 

1.84 0.938 1x10 -8 to 1x10 -2 

0.274 0.82 1x10 -8 to 1x10 -2 

3.40x10 -2 

4.90x10 -2 

4.00x10 -2 

0.51 1x10 -7 to 1x10 -3 

0.5 1x10 -7 to 1x10 -3 

Ca/Kaolinite/pH 2 3 




Na/Montmorillonite/pH 2 3 


Na/Montmorillonite./pH 7 3 


Ca/Montmorillonite/pH 2 3 




Granite/pH 8.2 4 

Granite/pH 8.2 4 

0.5 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 = Serne et 

al., 1993; 5 = Shiao et al., 1979.

Table D.9. Descriptive statistics of the cesium Freundlich equations (Table D.8) 

reported in the literature. 

Statistic a b 

Mean 252 0.696 

Standard Error 150.2 0.029 

Median 0.222 0.720 

Mode NA 0.815 

Standard Deviation 1019 0.198 

Sample Variance 1038711 0.039 

Range 6070 1.087 

Minimum 0.000013 0.013 

Maximum 6070 1.1 

95% Confidence Level 302 0.059 

Using the medians of the a and b parameters from the literature, we come up with the overall 

equation: 

Cs adsorbed = 0.222(Cs solution) 0.720 

This equation is plotted in Figure D.4. Using Cs adsorbed and Cs solution from equation D.3, a K d value 

can be calculated according to equations D.4, 

K d = Cs adsorbed/Cs solution. 

Cesium K d values calculated from Equations D.3 and D.4 are presented in Figure D.5. 

D.15 

(D.3) 

(D.4)

Figure D.4. Generalized cesium Freundlich equation 

(Equation D.3) derived from the literature. 

Figure D.5. Cesium K d values calculated from generalized 

Freundlich equation (Equations D.3 and D.4) 

derived from the literature. 

D.16

D.2.6 Approach to Selecting K d Values for Look-up Table 





correlation between pH and CEC, and pH and cesium K d 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 K d values and the solid phase parameters. Thus, the K d 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 K d 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 micalike 

minerals (Table D.11). For both tables, the user will be able to reduce the range of possible 

cesium K d 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.11) were assigned by multiplying the central estimates from 

Table D.10 by a factor of 2.5. The 2.5 scaler was selected based on relationships existing in the 

values in the data set and in Table D.6. Finally, the lower and upper limits for these central 

estimates were estimated based on the assumption that there was 2.5 orders of magnitude 

variability associated with the central estimates. The variability was based on visual inspection of 

a number of figures containing the cesium K d values, including Figure D.1. 

The calculations and equations used to estimate the central, minimum, and maximum estimates 

used in the look-up tables are presented in Table D.12. 

D.17

Table D.10. Estimated range of K d values (ml/g) for cesium based on CEC or clay content for 

systems containing

Mica 

Concentration 

in Clay Fraction 

(%) 

Table D.12. Calculations for values used in look-up table. 

Clay 

Content 

(wt.%) 

CE 1 

(ml/g) 

Logarithm Scale Base-10 Scale 

Log CE 

D.19 

Lower Limit 

(Log CE)/2 

Lower Limit 

10 (log CE)/2 (ml/g) 

Upper Limit 

10 log CE + (log CE)/2 (ml/g) 

D.3.0 K d Data Set for Soils and Pure Mineral Phases 

Table D.13 lists the available cesium K d values identified for experiments conducted with soils and 

pure mineral phases. 

Cesium 

Kd 

(ml/g) 

Clay 

(wt.% 

) 

Table D.13. Cesium K d data base for soils and pure mineral phases 

Mica 

(%) 

pH CEC a 

(meq/100 g) 

SA 1 

(m 2 /g) 

Aqueous Cs 

(µM) 

D.20 

Background 

Aqueous 

247 6.2 1.90x10 -2 Gorleben 

Groundwater 

Soil and Mineral 

Phase ID and 

Information 

Ref 2 

Gorleben Sediment 1 

62 6.2 1.42x10 -1 Gorleben Sediment 1 


16 6.2 1.05 Gorleben Sediment 1 

12 6.2 1.53 Gorleben Sediment 1 

167 8.1 189 5.20x10 -3 Groundwater-1 S1: Quartz, 

Kaolinite, 

Plagioclase 

1 7.8 113 5.20x10 -3 Groundwater-2 S2:Quartz, 

Kaolinite, Dolomite 

1500 9.3 60 70 1.00x10 -1 Water pH 9.3 Bentonite 3 

160 2.4 60 70 1.00x10 -1 Groundwater 

pH 2.4 

1100 9.3 60 70 1.00x10 -1 Groundwater 

pH 9.3 

Bentonite 3 

Bentonite 3 

4100 6.1 20 130 1.00x10 -1 Water pH 6.1 Takadate loam 3 

1400 7.7 20 130 1.00x10 -1 Groundwater 

pH 7.7 

Takadate loam 3 

1100 6.6 70 60 1.00x10 -1 Water pH 6.6 Hachinohe loam 3 

280 8.3 70 60 1.00x10 -1 Groundwater 

pH 8.3 

Hachinohe loam 3 

237 8.2 2 22 1.00x10 -3 ym-22 4 

8220 8.2 109 103 1.00x10 -3 ym-38 4 

325 8.2 6 43 1.00x10 -3 ym-45 4 

22100 8.2 51 19 1.00x10 -3 ym-48 4 

35800 8.2 107 1.00x10 -3 ym-49 4 

42600 8.2 107 1.00x10 -3 ym-49 4 

205 8.2 4 1.00x10 -3 ym-54 4 

2 

2

Cesium 


(ml/g) 

Clay 

(wt.% 

) 

Mica 

(%) 

pH CEC a 

(meq/100 g) 

SA 1 

(m 2 /g) 

Aqueous Cs 

(µM) 

Background 

Aqueous 


Phase ID and 

Information 

15200 8.4 31 1.00x10 -3 low salts JA-18 4 

8440 8.3 31 1.00x10 -3 hi salts JA-18 4 

143 8.2 8 1.00x10 -3 low salts JA-32 4 

73 8.5 8 1.00x10 -3 hi salts JA-32 4 

1390 8.4 100 1.00x10 -3 low salts JA-37 4 

757 8.5 100 1.00x10 -3 hi salts JA-37 4 

95 15 4 4.20x10 -4 0.005 M Na Savannah River 5 

120 15 5.5 4.20x10 -4 0.005 M Na Savannah River 5 


130 15 7 4.20x10 -4 0.005 M Na Savannah River 5 



72 3 4 4.20x10 -4 0.005 M Na 4-Mile Creek 5 

79 3 5.5 4.20x10 -4 0.005 M Na 4-Mile Creek 5 


98 3 7 4.20x10 -4 0.005 M Na 4-Mile Creek 5 


33 4 4 4.20x10 -4 0.005 M Na Par Pond Soil 5 

37 4 5.5 4.20x10 -4 0.005 M Na Par Pond Soil 5 




27 2 4 4.20x10 -4 0.005 M Na Steel Creek Soil 5 

25 2 5.5 4.20x10 -4 0.005 M Na Steel Creek Soil 5 





88 4 4 4.20x10 -4 0.005 M Na Lower 3 Runs Soil 5 

92 4 5.5 4.20x10 -4 0.005 M Na Lower 3 Runs Soil 5 


D.21 

Ref 2

Cesium 


(ml/g) 

Clay 

(wt.% 

) 

Mica 

(%) 

pH CEC a 

(meq/100 g) 

SA 1 

(m 2 /g) 

Aqueous Cs 

(µM) 

Background 

Aqueous 


Phase ID and 

Information 




88 5 4 4.20x10 -4 0.005 M Na Pen Branch Soil 5 

89 5 5.5 4.20x10 -4 0.005 M Na Pen Branch Soil 5 




22 2 4 4.20x10 -4 0.005 M Na Upper 3 Runs Soil 5 

31 2 5.5 4.20x10 -4 0.005 M Na Upper 3 Runs Soil 5 




27 8.25 1.83 17.7 2.72x10 2 0.002 M 

Groundwater 

329 8.25 1.83 17.7 2.90x10 -1 0.002 M 

Groundwater 

960 8.25 1.83 17.7 1.03x10 -3 0.002 M 

Groundwater 

1088 8.25 1.83 17.7 9.11x10 -6 0.002 M 

Groundwater 

1084 8.25 1.83 17.7 1.87x10 -6 0.002 M 

Groundwater 

28 8.6 1.83 17.7 2.63x10 2 0.013 M 

Groundwater 

289 8.6 1.83 17.7 3.31x10 -1 0.013 M 

Groundwater 

951 8.6 1.83 17.7 1.05x10 -3 0.013 M 

Groundwater 

1022 8.6 1.83 17.7 9.77x10 -6 0.013 M 

Groundwater 

1025 8.6 1.83 17.7 1.95x10 -6 0.013 M 

Groundwater 

18 8.2 1.5 10.3 3.61x10 2 0.002 M 

Groundwater 

189 8.2 1.5 10.3 5.00x10 -1 0.002 M 

Groundwater 

D.22 

Ref 2 

Umtanum Basalt 6 










Flow E Basalt 6 

Flow E Basalt 6

Cesium 


(ml/g) 

Clay 

(wt.% 

) 

Mica 

(%) 

pH CEC a 

(meq/100 g) 

SA 1 

(m 2 /g) 

Aqueous Cs 

(µM) 

Background 

Aqueous 

418 8.2 1.5 10.3 2.34x10 -3 0.002 M 

Groundwater 

450 8.2 1.5 10.3 2.17x10 -5 0.002 M 

Groundwater 

487 8.2 1.5 10.3 3.98x10 -6 0.002 M 

Groundwater 

20 8.7 1.5 10.3 3.39x10 2 0.013 M 

Groundwater 

214 8.7 1.5 10.3 4.47x10 -1 0.013 M 

Groundwater 

488 8.7 1.5 10.3 2.00x10 -3 0.013 M 

Groundwater 

549 8.7 1.5 10.3 1.78x10 -5 0.013 M 

Groundwater 

617 8.7 1.5 10.3 3.24x10 -6 0.013 M 

Groundwater 

48 8.3 4.84 31.2 1.71x10 2 0.002 M 

Groundwater 

460 8.3 4.84 31.2 2.13x10 -1 0.002 M 

Groundwater 

1111 8.3 4.84 31.2 8.30x10 -4 0.002 M 

Groundwater 

1466 8.3 4.84 31.2 6.37x10 -6 0.002 M 

Groundwater 

1281 8.3 4.84 31.2 1.39x10 -6 0.002 M 

Groundwater 

56 8.55 4.84 31.2 1.51x10 2 0.013 M 

Groundwater 

389 8.55 4.84 31.2 2.57x10 -1 0.013 M 

Groundwater 

853 8.55 4.84 31.2 1.17x10 -3 0.013 M 

Groundwater 

952 8.55 4.84 31.2 1.05x10 -5 0.013 M 

Groundwater 

908 8.55 4.84 31.2 1.74x10 -6 0.013 M 

Groundwater 

212 8.3 71 646 4.50x10 1 0.002 M 

Groundwater 

1080 8.3 71 646 9.17x10 -1 0.002 M 

Groundwater 

D.23 


Phase ID and 

Information 

Ref 2 









Pomona Basalt 6 










Smectite 6 

Smectite 6

Cesium 


(ml/g) 

Clay 

(wt.% 

) 

Mica 

(%) 

pH CEC a 

(meq/100 g) 

SA 1 

(m 2 /g) 

Aqueous Cs 

(µM) 

Background 

Aqueous 

13042 8.3 71 646 7.66x10 -5 0.002 M 

Groundwater 

9794 8.3 71 646 1.00x10 -6 0.002 M 

Groundwater 

25000 8.3 71 646 7.00x10 -8 0.002 M 

Groundwater 

224 9.2 71 646 4.27x10 -1 0.013 M 

Groundwater 

2136 9.2 71 646 4.68x10 -2 0.013 M 

Groundwater 

5882 9.2 71 646 1.70x10 -4 0.013 M 

Groundwater 

8547 9.2 71 646 1.17x10 -6 0.013 M 

Groundwater 

8333 9.2 71 646 2.40x10 -7 0.013 M 

Groundwater 

D.24 


Phase ID and 

Information 

Ref 2 

Smectite 6 

Smectite 6 

Smectite 6 

Smectite 6 

Smectite 6 

Smectite 6 

Smectite 6 

Smectite 6 

5000 24 4.4 82 6.80x10 -2 1x10 -6 M KCl Batcombe 7 




9000 42 6.2 72 6.80x10 -2 1x10 -6 M KCl Tedburn 7 




1050 42 7.3 54 6.80x10 -2 1x10 -6 M KCl Teigngrace 7 




11000 130 1.00x10 -7 Water Itago Tuff 8 

10000 97 1.00x10 -7 Water Ohya Tuff 8 

5000 2.4 1.00x10 -7 Water Sandstone 8 

2000 1.9 1.00x10 -7 Water Shale 8 

6000 1.9 1.00x10 -7 Water Augite Audesite 8 

500 1.2 1.00x10 -7 Water Plagio Rhyolite 8 

5800 0.75 1.00x10 -7 Water Olivine Basalt 8

Cesium 


(ml/g) 

Clay 

(wt.% 

) 

Mica 

(%) 

pH CEC a 

(meq/100 g) 

SA 1 

(m 2 /g) 

Aqueous Cs 

(µM) 

Background 

Aqueous 


Phase ID and 

Information 

900 0.54 1.00x10 -7 Water Ionada Granite 8 

260 0.35 1.00x10 -7 Water Rokka Granite 8 

80 0.033 1.00x10 -7 Water Limestone 8 

2200 1.2 1.00x10 -7 Water Biotite 8 

1800 0.93 1.00x10 -7 Water Chlorite 8 

630 0.33 1.00x10 -7 Water Hornblende 8 

420 0.11 1.00x10 -7 Water Grossular 8 

460 0.0067 1.00x10 -7 Water Forsterite 8 

30 0.0034 1.00x10 -7 Water K-feldspar 8 

89 0.0032 1.00x10 -7 Water Albite 8 

31 0.00098 1.00x10 -7 Water Quartz 8 

1 0.15849 1.00x10 -1 Calcite 9 

3 0.19953 1.00x10 -1 Apatite 9 

6 1.58489 1.00x10 -1 Hematite 9 

13 1.77828 1.00x10 -1 Orthoclase 9 

16 5.62341 1.00x10 -1 Serpentine 9 

200 7.94328 1.00x10 -1 Hornblende 9 

631 39.8107 1.00x10 -1 Biotite 9 

794 63.0957 1.00x10 -1 Muscovite 9 

100 4.46684 1.00x10 -1 Gneiss 9 

16 6.30957 1.00x10 -1 Diabase 9 

158 10 1.00x10 -1 Stripa Granite 9 

562 11.2202 1.00x10 -1 Finsjo Granite 9 

900 5 1.00x10 -1 Biotite 9 

790 7 1.00x10 -1 Biotite 9 

700 9 1.00x10 -1 Biotite 9 

2 5 1.00x10 -1 Hematite 9 

4 7 1.00x10 -1 Hematite 9 

8 9 1.00x10 -1 Hematite 9 

40 5 1.00x10 -1 Hornblende 9 

100 7 1.00x10 -1 Hornblende 9 

D.25 

Ref 2

Cesium 


(ml/g) 

Clay 

(wt.% 

) 

Mica 

(%) 

pH CEC a 

(meq/100 g) 

SA 1 

(m 2 /g) 

Aqueous Cs 

(µM) 

Background 

Aqueous 


Phase ID and 

Information 

240 9 1.00x10 -1 Hornblende 9 

3 5 1.00x10 -1 Magnetite 9 

5 7 1.00x10 -1 Magnetite 9 

9 9 1.00x10 -1 Magnetite 9 

700 5 1.00x10 -1 Muscovite 9 

810 7 1.00x10 -1 Muscovite 9 

840 9 1.00x10 -1 Muscovite 9 

7 5 1.00x10 -1 Orthoclase 9 

14 7 1.00x10 -1 Orthoclase 9 

7 9 1.00x10 -1 Orthoclase 9 

52000 127 1.67x10 -6 Deionized Water Vermiculite 10 

26000 20 1.67x10 -6 Deionized Water Illite 10 

2500 11.2 1.67x10 -6 Deionized Water Kaolinite 10 

2700 127 1.67x10 -6 0.1 N NaCl Vermiculite 10 

28600 20 1.67x10 -6 0.1 N NaCl Illite 10 

94 11.2 1.67x10 -6 0.1 N NaCl Kaolinite 10 

7 1.00x10 -7 Groundwater Hanford Vadose 

Sediment 


Sediment 

2190 4 9 7.7 8.40x10 -3 Groundwater Sediment CGS-1 12 

7610 5 12 8.2 8.40x10 -3 Groundwater Sediment TBS-1 12 

620 6 9 7.9 8.40x10 -3 Groundwater Sediment Trench-8 12 

1 CEC = cation exchange capacity; SA = surface area. 

2 References: 1 = Lieser and Steinkopff, 1989; 2 = Lieser et al., 1986; 3 =Konishi et al., 1988; 4 = Vine et al., 1980; 

5 = Elprince et al., 1977; 6 = Ames et al., 1982; 7 = Staunton, 1994; 8 = Akiba et al., 1989; 9 = Torstenfelt et al., 1982; 

10 = Tamura, 1972; 11 = Routson et al., 1980; 12 = Serne et al., 1993. 

D.26 

Ref 2 

11 

11

D.4.0 Data Set for Soils 

Table D.14 lists the available cesium K d values identified for experiments conducted with only 

soils. 

Cesium 

K d 

(ml/g) 

Clay 

(wt%) 

Mica 

(% ) 

Table D.14. Cesium K d data set for soils only. 

pH CEC (a) 

(meq/100 

g) 

SA 1 

(m 2 /g) 

D.27 

Cs 

(µM) 

Aqueous 

Phase 

247 6.2 1.90x10 -2 Gorleben 

Groundwater 

Soil ID 

and Information 

Ref. 2 

Gorleben Sediment 1 



4100 6.1 20 130 1.00x10 -1 Water pH 6.1 Takadate Loam 4 

1400 7.7 20 130 1.00x10 -1 Groundwater 

pH 7.7 

Takadate Loam 4 

1100 6.6 70 60 1.00x10 -1 Water pH 6.6 Hachinohe Loam 4 

280 8.3 70 60 1.00x10 -1 Groundwater 

pH 8.3 

Hachinohe loam 4 

95 15 4 4.20x10 -4 0.005 M Na Sav. River Site 

Sediment 

120 15 5.5 4.20x10 -4 0.005 M Na Sav. River Site 

Sediment 


Sediment 

130 15 7 4.20x10 -4 0.005 M Na Sav. River Site 

Sediment 


Sediment 


Sediment 

72 3 4 4.20x10 -4 0.005 M Na 4-Mile Creek Sediment 6 

79 3 5.5 4.20x10 -4 0.005 M Na 4-Mile Creek Sediment 6 

75 3 6.7 4.20x10 -4 0.005 M Na 4-Mile Creek 

Sediment. 

98 3 7 4.20x10 -4 0.005 M Na 4-Mile Creek 

Sediment. 

83 3 8.5 4.20x10 -4 0.005 M Na 4-Mile Creek 

Sediment. 


6 

6 

6 

6 

6 

6 

6 

6 

6

Cesium 

K d 

(ml/g) 

Clay 

(wt%) 

Mica 

(% ) 

pH CEC (a) 

(meq/100 

g) 

SA 1 

(m 2 /g) 

Cs 

(µM) 

Aqueous 

Phase 

Soil ID 













92 4 5.5 4.20x10 -4 0.005 M Na Lower 3 Runs 

Sediment 


Sediment 

85 4 7 4.20x10 -4 0.005 M Na Lower 3 Runs 

Sediment 


Sediment 


Sediment 












Sediment 


Sediment 

D.28 

Ref. 2 

6 

6 

6 

6 

6 

8 

8

Cesium 

K d 

(ml/g) 

Clay 

(wt%) 

Mica 

(% ) 

pH CEC (a) 

(meq/100 

g) 

SA 1 

(m 2 /g) 

Cs 

(µM) 

Aqueous 

Phase 

Soil ID 


3,000 6 7.6 3 8.6 1.00x10 -1 Groundwater Sediment A 10 

4,800 7.5 5.9 4.3 12.2 1.00x10 -1 Groundwater Sediment B 10 

3,100 8 6.6 4.7 14.7 1.00x10 -1 Groundwater Sediment C 10 

3,000 5 8 2.6 6.6 1.00x10 -1 Groundwater Sediment D 10 

2,190 4 9 7.7 8.40x10 -3 Groundwater Sediment CGS-1 11 

7,610 5 12 8.2 8.40x10 -3 Groundwater Sediment TBS-1 11 

620 6 9 7.9 8.40x10 -3 Groundwater Sediment Trench-8 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 

et al., 1992; 11 = Serne et al., 1993. 

D.5.0 References 

Adeleye, S. A., P. G. Clay, and M. O. A. Oladipo. 1994. “Sorption of Caesium, Strontium and 

Europium Ions on Clay Minerals.” Journal of Materials Science, 29:954-958. 

Akiba, D., and H. Hashimoto. 1990. “Distribution Coefficient of Strontium on Variety of 

Minerals and Rocks.” Journal of Nuclear Science and Technology, 27:275-279. 

Akiba, D., H. Hashimoto, and T. Kanno. 1989. “Distribution Coefficient of Cesium and Cation 

Exchange Capacity of Minerals and Rocks.” Journal of Nuclear Science and Technology, 

26:1130-1135. 

Ames, L., and D. Rai. 1978. Radionuclide Interactions with Soil and Rock Media. Volume 1: 

Processes Influencing Radionuclide Mobility and Retention, Element Chemistry and 

Geochemistry, Conclusions and Evaluation. PB-292 460, Pacific Northwest Laboratory, 


Ames, L. L., J. E. McGarrah, B. A. Walker, and P. F. Salter. 1982. “ Sorption of Uranium and 

Cesium by Hanford Basalts and Associated Secondary Smectite.” Chemical Geology, 

35:205-225. 

Comans, R. N. J., J. J. Middelburg, J. Zonderhuis, J. R. W. Woittiez, G. J. De Lange, H. A. Das, 

C. H. Van Der Weijden. 1989. “Mobilization of Radiocaesium in Pore Water in Lake 

Sediments.” Nature, 367-369. 

Cremers, A., A. Elsen. P. De Preter, and A. Maes. 1988. “Quantitative Analysis of 

Radiocaesium Retention in Soils.” Nature, 335:247-249. 

D.29 

Ref. 2

Bruggenwert, M. G. M., and A. Kamphorst. 1979. “Survey of Experimental Information on 

Cation Exchange in Soil Systems.” In Soil Chemistry: B. Physico-Chemical Models, G. H. 

Bolt (ed.), Elsevier Scientific Publishing Company, New York, New York. 

Dahlman, R. C., E. A. Bondietti, and L. D. Eyman. 1976. “Biological Pathways and Chemical 

Behavior of Plutonium and Other Actinides in the Environment.” In Actinides in the 

Environment, A. M. Friedman (ed.), pp. 47-80. ACS Symposium Series 35, American 


Douglas, L. A. 1989. “Vermiculites.” In Minerals in Soil Environments, J. B. Dixon and S. B. 

Week (eds.), Second Edition, pp. 635-674, Soil Science Society of America, Madison, 

Wisconsin. 

Elprince, A. M., C. I. Rich, and D. C. Martens. 1977. “Effect of Temperature and Hydroxy 

Aluminum Interlayers on the Adsorption of Trace Radioactive Cesium by Sediments near 

Water-Cooled Nuclear Reactors.” Water Resources Research, 13:375-380. 

Erten, H. N., S. Aksoyoglu, S. Hatipoglu, and H. Göktürk. 1988. “Sorption of Cesium and 

Strontium on Montmorillonite and Kaolinite.” Radiochimica Acta, 44/45:147-155. 

Evans, D. W., J. J. Alberts, and R. A. Clark. 1983. “Reversible Ion-Exchange Fixation of 

Cesium-137 Leading to Mobilization from Reservoir Sediments.” Geochimica et 


Francis, C. W., and F. S. Brinkley. 1976. “Preferential Adsorption of 137 Cs to Micaceous 

Minerals in Contaminated Freshwater Sediments.” Nature, 260:511-513. 

Fukui, M. 1990. “Desorption Kinetics and Mobility of Some Radionuclides in Sediments.: 

Health Physics, 59:879-889. 

Hartman, M. J., and P. E. Dresel. 1997. Hanford Site Groundwater Monitoring for Fiscal Year 

1996. PNNL-11470, Pacific Northwest National Laboratory, Richland, Washington. 


Water Supply Paper 2254, U.S. Geological Survey, Alexandria, Virginia. 

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. 

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. 

Legoux, Y., G. Blain, R. Guillaumont, G. Ouzounian, L. Brillard, and M. Hussonnois. 1992. “K d 

Measurements of Activation, Fission, and Heavy Elements in Water/Solid Phase Systems.” 


D.30

Li, Y., L. Burkhardt, M. Buchholtz, P. O’Hara, and P. H. Santschi. 1994. “Partition of 

Radiotracers Between Suspended Particles and Seawater.” Geochimica et Cosmochimica 

Acta, 48:2011-2019. 

Lieser, K. H., B. Gleitsmann, and Th. Steinkopff. 1986. “Sorption of Trace Elements or 

Radionuclides in Natural Systems Containing Groundwater and Sediments.” Radiochimica 

Acta, 40:33-37. 

Lieser, K. H., and Th. Steinkopff. 1989. “Sorption Equilibria of Radionuclides or Trace 

Elements in Multicomponent Systems.” Radiochimica Acta, 47:55-61. 

Neter, J., and W. Wasserman. 1974. Applied Linear Statistical Models. Richard D. Irwin, Inc. 

Homewood, Illinois. 

Ohnuki, T. 1991. “Characteristics of Migration of 85 Sr and 137 Cs in Alkaline Solution Through 

Sandy Soil.” Material Research Society Proceedings, 212:609-616. 

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. 



Washington. 

Routson, R. C., G. S. Barney, and R. M. Smith. 1980. Hanford Site Sorption Studies for the 

Control of Radioactive Wastes: A Review. WHO-SA-155, Rev. 1, Rockwell Hanford 


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. 

Sawhney, B. L. 1972. “Selective Sorption and Fixation of Cations by Clay Minerals: A Review.” 

Clays and Clay Minerals, 20:93-100. 




Characterization. PNL-8889, Pacific Northwest National Laboratory, Richland, 

Washington. 

Serne, R. J., and V. L. LeGore. 1996. Strontium-90 Adsorption-Desorption Properties and 

Sediment Characterization at the 100 N-Area. PNL-10899, Pacific Northwest National 


D.31

Shiao, S. Y., P. Rafferty, R. E. Meyer, and W. J. Rogers. 1979. “Ion-Exchange Equilibria 

Between Montmorillonite and Solutions of Moderate-to-High Ionic Strength.” In 

Radioactive Waste in Geologic Storage, S. Fried (ed.), pp. 297 324, ACS Symposium 

Series 100, American Chemical Society, Washington, D.C. 

Smith, J. T., and R. N. J. Comans. 1996. “Modelling the Diffusive Transport and Remobilization 

of 137 Cs in Sediments: The Effects of Sorption Kinetics and Reversibility.” Geochimica et 


Sposito, G. 1984. The Surface Chemistry of Soils. Oxford University Press, New York, New 

York. 

Sposito, G. 1989. The Chemistry of Soils. Oxford University Press, New York, New York. 

Staunton, S. 1994. “Adsorption of Radiocaesium on Various Soils: Interpretation and 

Consequences of the Effects of Soil:Solution Ratio and Solution Composition on the 

Distribution Coefficient.” European Journal of Soil Science, 45:409-418. 


Environmental Pollutant Assessment System. PNL-7145, Pacific Northwest National 


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)

E.1.0 Background 

Appendix E 

Partition Coefficients For Chromium(VI) 

The review of chromium K d data obtained for a number of soils (summarized in Table E.1) 

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: 

C 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. 

C Increasing pH decreases adsorption (decrease in K d) of Cr(VI) on minerals and soils. The 

data are quantified for only a limited number of soils. 

C 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 

K d). Soils containing Mn oxides oxidize Cr(III) into Cr(VI) form thus resulting in lower 

K d values. The relation between oxide/hydroxide contents of iron and manganese and 

their effects on K d have not been adequately quantified except for a few soils. 

C The presence of competing anions reduce Cr(VI) adsorption. The inhibiting effect varies 

2- - 2- 2- - - - 

in the order HPO4 , H2PO4 >>SO4 CO3 /HCO3 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 K d values (478,630 ml/g) at 

lower pH values (5.65), and lower K d values (6,607 ml/g) at higher pH values (7.80). In the 

presence of competing anions (SO 4: 2.5x10 -3 M, solution in equilibrium with 3.5x10 -3 atm CO 2), 

at the same pH values, the observed K d 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, 

Windsor, and Olivier soils with pH values of 5.1, 5.4, and 6.4 exhibited chromium K d values in the 

E.2

ange ~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 d values for 

these soils. The data showed that 3 of the 4 soils studied exhibited decreasing K d values with 

increasing pH. The K d values for these soils were close to 1 ml/g at higher pH values (>8). At 

lower pH values (about 4.5) the K d 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 d # 1 ml/g) for Cr(VI). 

The data regarding the effects of soil organic matter on Cr(VI) adsorption are rather sparse. In 

1 study, Stollenwerk and Grove (1985) evaluated the effects of soil organic matter on adsorption 

of Cr(VI). Their results indicated that organic matter did not influence Cr(VI) adsorption 

properties. In another study, the Cr(VI) adsorption properties of an organic soil was examined by 

Wong et al. (1983). The chromium adsorption measurements on bottom, middle, and top layers 

of this soil produced K d values of 346, 865, and 2,905 ml/g respectively. Also, another K d 

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 K d values that ranged from 6 to 263 ml/g (Sheppard et al., 1987). 

These measurements showed that coarse-textured soils tend to have lower K d values as compared 

to fine-textured soils such as loam (K d ~ 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 K d 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 d value. However, 

removing iron oxide and hydroxide coatings resulted in a K d value of about 0.25 leading the 

authors to conclude that a major fraction of Cr(VI) adsorption capacity of this soil is due to its 

iron oxide and hydroxide content. Desorption experiments conducted on Cr adsorbed soil aged 

for 1.5 yrs indicated that over this time period, a fraction of Cr(VI) had been reduced to Cr(III) 

by ferrous iron and had probably coprecipitated with iron hydroxides. 

Studies by Stollenwerk and Grove (1985) and Sheppard et al. (1987) using soils showed that K d 

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 d 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 K d values or any correlative relationships between K d and the oxide contents. The 

adsorption data obtained by Rai et al. (1988) also showed that quantities of sodium dithionitecitrate-bicarbonate 

(DCB) extractable iron content of soils is a good indicator of a soil’s ability to 

reduce Cr(VI) to Cr(III) oxidation state. The reduced Cr has been shown to coprecipitate with 

ferric hydroxide. Therefore, observed removal of Cr(VI) from solution when contacted with 

E.3

chromium-reductive soils may stem from both adsorption and precipitation reaction. Similarly, 

Rai et al. (1988) also showed that certain soils containing manganese oxides may oxidize Cr(III) 

into Cr(VI). Depending on solution concentrations, the oxidized form (VI) of chromium may also 

precipitate in the form of Ba(S,Cr)O 4. Such complex geochemical behavior chromium in soils 

implies that depending on the properties of a soil, the measured K d values may reflect both 

adsorption and precipitation reactions. 

An evaluation of competing anions indicated that Cr(VI) adsorption was inhibited to the greatest 

2- - - - 

extent by HPO4 and H2PO4 ions and to a very small extent by Cl and NO3 ions. The data 

2- - 2- - 

indicate that Cr(VI) adsorption was inhibited by anions in order of HPO4 , H2PO4 >> SO4 >> Cl , 

- (Leckie et al., 1980; MacNaughton, 1977; Rai et al., 1986; Rai et al., 1988; Stollenwerk and 

NO 3 

Grove, 1985). 

E.4

Table E.1. Summary of K d values for Cr(VI) adsorption on soils. 

Experimental Parameters Reference 


pH CEC 

(meq/100g) 

Iron 

Oxide 

Conten1 (wt.%) 

Organic 

Carbon 

(wt.%) 

Soil Description Clay 

Content 

(wt.%) 

Wong et al. 

(1983) 

NR 

2905 

0.453 

7.1 

NR 

7.05 

NR 

NR 

865 

0.409 

7.2 

NR 

6.71 

NR 

NR 

346 

0.158 

7.3 

NR 

2.79 

NR 

NR 

1729 

0.113 

8.2 

NR 

1.45 

NR 

Organic Soil (Muck) Top 

Layer, Florida 

Organic Soil (Muck) Middle 


Organic Soil (Muck) Bottom 


Hallandale Fine sand, Florida 

Stollenwerk and 

Grove (1985) 

Batch experiment, deionized water , eq. Cr conc. 

1.4 - 0.0004 mmol/l 

Batch experiment, groundwater (pH: 6.8) 

Batch experiment, groundwater, Soil with org matter removed 

Batch experiment, groundwater, Soil with iron oxides removed 

Column experiment, groundwater, initial Cr conc. (0.01 mmol/l) 

Alluvium, Telluride, Colorado 1 0.1 1.2 6.45 NR 1.7 - 52 

E.5 

5.3 

5.6 

0.25 

2.35 

Sheppard and 

Sheppard 

(1987) 

NR 

NR 

1000 

100 

60 

1.6 

NR 

NR 

NR 

NR 

NR 

NR 

NR 

NR 

Loam (Chernozem), Canada 

Sand (Regosol), Canada 

Sheppard et al., 

1987 

Column experiments (lysimeter). Solutions: leachate, groundwater 

263, 6 

8.1 

5.2 

NR 

NR 

NR 

Column experiments (lysimeter). Solutions:leachate, groundwater 



91, 35 

135,160 

53, 9 

0.29 

0.21 

0.17 

5.1 

5.2 

6.2 

NR 

NR 

NR 

NR 

NR 

NR 

NR 

NR 

NR 

Sand (Brunisol) organic surface 

layer (LFH-Ah) 

Sand (Brunisol) upper layer (Ae) 

Sand (Brunisol) middle layer (Bfj) 

Sand (Brunisol) lower layer (Bfjgj) 

1 Total iron oxide (Fe2O 3) content of soils. Values within parenthesis represent DCB extractable Fe content (mmol/g) of soils.

E.2.0 Approach 

The approach used to develop the look-up table was to identify the key parameters that control Cr(VI) 

adsorption reactions. From the data of Rai 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 d values are significantly higher at 

lower pH values and decline with increasing pH. Also, K d 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(III) 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 K d 

values. It is important to note the total iron oxide content is a poor indicator of a soil’s Cr(VI) reducing 

capacity and that DCB extractable iron better represents the fraction of iron content that would reduce Cr(VI) 

to Cr(III). The data indicated that Holton/Cloudland soil with the highest concentrations of DCB extractable 

iron (0.435 mmol/g) exhibited higher K d values than other soils which did not show an observable Cr(VI) 

reduction tendency. 

Based on this information, 4 ranges of pH, which encompass the pH range of most natural soils, were selected 

for the look-up table (Table E.3). Within each pH range, 3 ranges of DCB extractable iron content were 

selected to represent the categories of soils that definitely reduce ($0.3 mmol/g), probably reduce (0.26 to 

0.29 mmol/g), and do not reduce (#2.5 mmol/g) Cr(VI) to Cr(III) form. The range of K d values to be 

expected within each of the 12 categories was estimated from the data listed in Table E.2. The variations of 

K d values as a function of pH and DCB extractable iron as independent variables based on experimental data 

(Table E.2) is also shown as a 3-dimensional graph (Figure E.1). The graph indicates that soils with lower pH 

values and higher DCB extractable iron contents exhibit greater adsorption (higher K d) of Cr(VI). At higher 

pH values (>7), Cr(VI) adsorption tends to be very low (very low K d values) irrespective of DCB extractable 

iron content. Similarly, soils which contain very low DCB extractable iron, adsorb very little Cr(VI) (very 

low K d values) irrespective of soil pH values. 

2- - 

Additionally, Cr(VI) adsorption studies show that the presence of competing anions such as HPO4 , H2PO4, 2- 2- - 

SO4 , CO3 , and HCO3 will reduce the Kd values as compared to a noncompetitive adsorption process. The 

only available data set that can be used to assess the competing anion effect was developed by Rai et al. 

2- 2- - 

(1988). However, they used fixed concentrations of competing anions namely SO4 , CO3 , and HCO3 (fixed 

through a single selected partial pressure of CO2) concentrations (Tables E.4 and E.5). Among these 

2- -3 

competing anions, SO4 at about 3 orders of magnitude higher concentrations (2 x 10 M or 191.5 mg/l) than 

Cr(VI) concentration depressed Cr(VI) K d 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 

2- -3 

Cr(VI) would be reduced as soluble SO4 concentrations increase from 0 to 2x10 M (or 191.5 mg/l). 

E.7

Table E.2. Data from Rai et al. (1988) for the adsorption of Cr(VI) as a function of pH. 

Kenoma Soil Cecil/Pacolet Soil Holton/Cloudland Soil Ocala Soil 

pH -log C 

(mol/m3 -log S Kd pH -log C 

) (mol/kg) (ml/g) (mol/m3 -log S Kd pH -log C 

) (mol/kg) (ml/g) (mol/m3 -log S Kd pH -log C 

) (mol/kg) (ml/g) (mol/m3 -log S Kd ) (mol/kg) (ml/g) 

8.42 3.03 6.25 1 9.26 3.05 5.66 2 9.84 3.03 6.33 1 9.37 3.02 6.56 0 

7.71 3.05 5.84 2 9.29 3.05 5.88 1 8.67 3.04 6.11 1 9.40 3.03 6.05 1 

7.70 3.04 5.97 1 8.57 3.11 5.34 6 8.60 3.08 5.60 3 8.94 3.02 7.71 0 

7.35 3.09 5.54 4 7.80 3.30 5.00 20 8.29 3.09 5.53 4 8.94 3.02 6.67 0 

7.40 3.08 5.59 3 7.41 3.44 4.89 35 8.27 3.07 5.70 2 8.67 3.04 6.00 1 

7.20 3.03 5.36 5 7.38 3.46 4.88 38 8.08 3.11 5.45 5 8.61 3.02 6.36 0 

7.16 3.13 5.37 6 6.99 3.66 4.81 71 7.55 3.30 5.04 18 8.33 3.04 6.00 1 

6.89 3.16 5.27 8 6.94 3.65 4.81 69 7.15 3.44 4.92 33 8.30 3.03 6.07 1 

6.92 3.15 5.29 7 6.67 3.79 4.78 102 7.05 3.51 4.89 42 7.56 3.03 6.14 1 

6.70 3.23 5.13 13 6.49 3.79 4.78 102 6.96 3.60 4.85 56 7.53 3.02 6.48 0 

6.47 3.26 5.09 15 6.19 3.99 4.75 174 6.88 3.61 4.85 58 7.53 3.02 6.86 0 

6.02 3.36 4.98 24 6.16 3.94 4.75 155 6.26 4.26 4.74 331 7.07 3.03 6.25 1 

6.02 3.35 4.99 23 5.89 4.08 4.74 219 6.20 4.25 4.74 324 7.18 3.03 6.19 1 

5.61 3.39 4.95 28 5.84 4.06 4.74 209 5.40 4.55 4.73 661 6.92 3.03 6.21 1 

5.62 3.40 4.95 28 5.46 4.19 4.73 288 5.39 4.63 4.73 794 6.88 3.02 6.48 0 

5.49 4.21 4.73 302 4.90 4.75 4.73 1047 6.61 3.03 6.12 1 

4.98 4.33 4.72 407 4.87 4.74 4.73 1023 5.71 3.02 6.68 0 

4.98 4.32 4.72 398 4.63 4.79 4.72 1175 5.14 3.04 6.01 1 

4.49 4.52 4.71 646 4.63 4.80 4.72 1202 

4.49 4.39 4.72 468 4.51 4.85 4.72 1349 

4.51 4.82 4.72 1259 

4.50 4.88 4.72 1445 

4.45 4.92 4.72 1585 

E.8

Table E.3. Estimated range of K d values for Cr(VI) as a function of soil pH, extractable iron content, and soluble sulfate. 

pH 

4.1 - 5.0 5.1 - 6.0 6.1 - 7.0 $7.1 


(mmol/g) 


(mmol/g) 


(mmol/g) 


(mmol/g) 


#0.25 0.26 - 0.29 $0.30 #0.25 0.26 - 0.29 $0.30 ##0.25 0.26 - 0.29 $0.30 #0.25 0.26 - 0.29 $0.30 

Soluble 

Sulfate 

Conc 

(mg/l) 

Min 25 400 990 20 190 390 8 70 80 0 0 1 

0 - 1.9 

Max 35 700 1770 34 380 920 22 180 350 7 30 60 

E.9 

Min 12 190 460 10 90 180 4 30 40 0 0 1 

2 - 18.9 

Max 15 330 820 15 180 430 10 80 160 3 14 30 

Min 5 90 210 4 40 80 2 15 20 0 0 0 

19 - 189 

Max 8 150 380 7 80 200 5 40 75 2 7 13 

Min 3 40 100 2 20 40 1 7 8 0 0 0 

$190 

Max 4 70 180 3 40 90 2 20 35 1 3 6

Figure E.1. Variation of K d for Cr(VI) as a function of pH and DCB extractable iron 

content without the presence of competing anions. 

E.3.0 Data Set for Soils 

The data set used to develop the look-up table is from the adsorption data collected by Rai et al. (1988). The 

adsorption data for Cr(VI) as a function of pH developed for 4 well-characterized soils were used to calculate 

the K d 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 K d 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. 

pH -log C 

(mol/m 3 ) 

Cr(VI) 1 Cr(VI) + Sulfate 1 Cr(VI) + Carbonate 1 

-log S 

(mol/kg) 

K d 

(ml/g) 

pH -log C 

(mol/m 3 ) 

-log S 

(mol/kg) 

E.11 

K d 

(ml/g) 

pH -log C 

(mol/m 3 ) 

-log S 

(mol/kg) 

9.26 3.05 5.66 2 8.92 3.05 6.27 1 9.62 3.05 6.88 0 

9.29 3.05 5.88 1 8.38 3.07 5.71 2 9.15 3.05 6.79 0 

8.57 3.11 5.34 6 8.38 3.04 5.70 2 9.01 3.06 6.35 1 

7.80 3.30 5.00 20 7.70 3.12 5.28 7 7.92 3.06 6.12 1 

7.41 3.44 4.89 35 7.67 3.12 5.28 7 7.95 3.06 6.10 1 

7.38 3.46 4.88 38 7.37 3.19 5.11 12 7.53 3.08 5.85 2 

6.99 3.66 4.81 71 7.24 3.23 5.09 14 7.52 3.07 6.06 1 

6.94 3.65 4.81 69 6.85 3.34 4.95 24 7.19 3.12 5.55 4 

6.67 3.79 4.78 102 6.76 3.37 4.96 26 7.31 3.10 5.67 3 

6.49 3.79 4.78 102 6.58 3.43 4.92 32 7.22 3.12 5.55 4 

6.19 3.99 4.75 174 6.56 3.34 4.95 25 6.99 3.13 5.48 4 

6.16 3.94 4.75 155 6.15 3.55 4.85 50 6.70 3.22 5.21 10 

5.89 4.08 4.74 219 6.15 3.51 4.88 43 6.68 3.21 5.24 9 

5.84 4.06 4.74 209 5.75 3.58 4.82 58 5.84 3.65 4.87 60 

5.46 4.19 4.73 288 5.79 3.56 4.86 51 6.08 3.54 4.91 43 

K d 

(ml/g) 

5.49 4.21 4.73 302 5.35 3.60 4.83 59 5.12 4.11 4.78 214 

4.98 4.33 4.72 407 5.33 3.59 4.84 57 5.12 4.14 4.78 229 

4.98 4.32 4.72 398 4.68 3.55 4.86 49 4.76 4.20 4.78 263 

4.49 4.52 4.71 646 4.69 3.47 4.86 41 4.75 4.11 4.78 214 

4.49 4.39 4.72 468 4.33 4.39 4.76 427 

1 Cr(VI) concentration: 10 -6 M, Sulfate Concentration: 10 -2.7 M, CO 2 : 10 -1.6 atm. 

4.34 4.37 4.77 398

Table E.5. Data from Rai et al. (1988) on effects of competing anions on 

Cr(VI) adsorption on Kenoma soil. 

pH -log C 

(mol/m 3 ) 

Cr(VI) 1 Cr(VI) + Sulfate + Carbonate 1 

-log S 

(mol/kg) 

K d 

(ml/g) 

E.12 

pH -log C 

(mol/m 3 ) 

-log S 

(mol/kg) 

8.42 3.03 6.25 1 7.49 3.06 6.22 1 

7.71 3.05 5.84 2 7.42 3.06 6.35 1 

7.70 3.04 5.97 1 7.3 3.07 5.98 1 

7.35 3.09 5.54 4 7.38 3.08 5.9 2 

7.40 3.08 5.59 3 7.08 3.08 5.83 2 

7.20 3.03 5.36 5 6.93 3.1 5.64 3 

7.16 3.13 5.37 6 6.49 3.15 5.43 5 

6.89 3.16 5.27 8 6.52 3.16 5.39 6 

6.92 3.15 5.29 7 6.32 3.17 5.33 7 

6.70 3.23 5.13 13 6.32 3.18 5.31 7 

K d 

(ml/g) 

6.47 3.26 5.09 15 5.97 3.23 5.21 10 

6.02 3.36 4.98 24 5.97 3.21 5.25 9 

6.02 3.35 4.99 23 5.7 3.23 5.2 11 

5.61 3.39 4.95 28 5.69 3.24 5.18 11 

5.62 3.40 4.95 28 5.54 3.24 5.19 11 

5.52 3.25 5.18 12 

5.03 3.18 5.32 7 

5.02 3.21 5.26 9 

1 Cr(VI) concentration: 10 -6 M, Sulfate Concentration: 10 -2.7 M, CO2 : 10 -1.6 atm.

E.4.0 References 

Davis, J. A. and J. O. Leckie. 1980. “Surface Ionization and Complexation at the Oxide/Water Interface. 3. 

Adsorption of Anions.” Journal of Colloid Interfacial Science, 74:32-43. 

Griffin, R. A., A. K. Au, and R. R. Frost. 1977. “Effect of pH on adsorption of Chromium form Landfill- 

Leachate by Clay Minerals.” Journal of Environmental Science Health, 12:431-449. 

Korte N. E., J. Skopp, W. H. Fuller, E. E. Niebla and B. A. Alesii. 1976. “Trace Element Movement in 

Soils: Influence of Soil Physical and Chemical Properties.” Soil Science, 122:350-359. 

Leckie, J. O., M. M. Benjamin, K. Hayes, G. Kaufman, and S. Altman. 1980. Adsorption/Coprecipitation of 

Trace Elements from Water with Iron Oxyhydroxides. EPRI-RP-910. Electric Power Research Institute, 

Palo Alto, California. 

MacNaughton, M. G. 1977. “Adsorption of Chromium (VI) at the Oxide-Water Interface.” In Biological 

Implications of Metals in the Environment, H. Drucker and R. F. Wildung (eds.), pp. 244-253, CONF- 

750929, National Technical Information Service, Springfield, Virginia. 

Rai, D., J. M. Zachara, L. E. Eary, C. C. Ainsworth, J. E. Amonette, C. E. Cowan, R. W. Szelmeczka, C. 

T. Resch, R. L. Schmidt, D. C. Girvin, and S. C. Smith. 1988. Chromium reactions in Geological 

Materials. EPRI-EA-5741. Electric Power Research Institute, Palo Alto, California. 

Rai, D., J. M. Zachara, L. E. Eary, D. C. Girvin, D. A. Moore, C. T. Resch, B. M. Sass, and R. L. Schmidt. 

1986. Geochemical Behavior of Chromium Species. EPRI-EA-4544. Electric Power Research Institute, 

Palo Alto, California. 

Ramirez, L. M., J. B. Rodriguez and F. Barba. 1985. “Heavy Metal Concentration in Sludge-Soil Systems as 

a result of Water Infiltration.” In Tropical Hydrology and Caribbean Island Water Resources Congress, 

F. Quinones and A. N. Sanchez (eds.), pp. 20-25, American Water Resources Association, Bethesda, 

Maryland. 

Rhoades, J. D. 1996. “Salinity: electrical Conductivity and Total Dissolved Solids.” In Methods of Soil 

Analysis, Part 3, Chemical Methods, J. M. Bigham (ed.), pp. 417-436. Soil Science Society of America 

Inc. Madison, Wisconsin. 

Richards, L. A. 1954. Diagnosis and Improvement of Saline and Alkali Soils. Agricultural Handbook 60, 

U. S. Department of Agriculture, Washington, D.C. 

Selim, H. M. and M C. Amcher. 1988. “A Second-Order Kinetic Approach for Modeling Solute Retention 

and transport in Soils.” Water Resources Research, 24:2061-2075. 

E.13



Sheppard, M. I., and S. C. Sheppard. 1987. “A Solute Transport Model Evaluated on Two Experimental 

Systems.” Ecological Modeling, 37:191-206. 

Stollenwerk, K. G., and D. B. Grove. 1985. “Adsorption and Desorption of Hexavalent Chromium in an 

Alluvial Aquifer Near Telluride, Colorado.” Journal of Environmental Quality, 14:150-155. 

Wong, K. V., S. Sengupta, D. Dasgupta, E. L. Daly, N. Nemerow, and H. P. Gerrish. 1983. “Heavy Metal 

Migration in Soil-Leachate Systems.” Biocycle, 24:30-33. 

E.14

APPENDIX F 

Partition Coefficients For Lead

F.1.0 Background 

Appendix F 

Partition Coefficients For Lead 

The review of lead K d 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: 

C Lead may precipitate in soils if soluble concentrations exceed about 4 mg/l at pH 4 and 

about 0.2 mg/l at pH 8. In the presence of phosphate and chloride, these solubility limits 

may be as low as 0.3 mg/l at pH 4 and 0.001 mg/l at pH 8. Therefore, in experiments in 

which concentrations of lead exceed these values, the calculated K d values may reflect 

precipitation reactions rather than adsorption reactions. 

C 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. 

C A number of adsorption studies indicate that within the pH range of soils (4 to 11), lead 

adsorption increases with increasing pH. 

C Adsorption of lead increases with increasing organic matter content of soils. 

C Increasing equilibrium solution concentrations correlates with decreasing lead adsorption 

(decrease in K d). 

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/l at pH 4 and about 0.2 mg/l at pH 

8. In the presence of phosphate and chloride, these solubility limits may be as low as 0.3 mg/l at 

pH 4 and 0.001 mg/l at pH 8. Therefore, in experiments in which concentrations of lead exceed 

these values, the calculated K d 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

(CEC) of the soils. Such an anomaly may have resulted from precipitation reactions brought 

about by high initial lead concentrations used in these experiments (20 to 830 mg/l). 

Lead adsorption characteristics of 7 alkaline soils from India were determined by Singh and 

Sekhon (1977). The authors concluded that soil clay, organic matter, and the calcium carbonate 

influenced lead adsorption by these soils. However, the initial lead concentrations used in these 

experiments ranged from 5 to 100 mg/l, indicating that in these alkaline soils the dominant lead 

removal mechanism was quite possibly precipitation. 

In another adsorption study, Abd-Elfattah and Wada (1981) measured the lead adsorption 

behavior of 7 Japanese soils. They concluded that soil mineral components which influenced lead 

adsorption ranged in the order: iron oxides>halloysite>imogolite, allophane>humus, 

kaolinite>montmorillonite. These data may not be reliable because high lead concentrations (up 

to 2,900 mg/l) used in these experiments may have resulted in precipitation reactions dominating 

the experimental system. 

Anionic constituents, such as phosphate, chloride, and carbonate, are known to influence lead 

reactions in soils either by precipitation of minerals of limited solubility or by reducing adsorption 

through complex formation (Rickard and Nriagu, 1978). A recent study by Bargar et al. (1998) 

showed that chloride solutions could induce precipitation of lead as solid PbOHCl. Presence of 

synthetic chelating ligands such as ethylenediaminetetraacetic acid (EDTA) has been shown to 

reduce lead adsorption on soils (Peters and Shem, 1992). These investigators showed that the 

presence of strongly chelating EDTA in concentrations as low as 0.01 M reduced K d for lead by 

about 3 orders of magnitude. By comparison quantitative data is lacking on the effects of more 

common inorganic ligands (phosphate, chloride, and carbonate) on lead adsorption on soils. 

A number of adsorption studies indicate that within the pH range of soils (4 to 11), lead 

adsorption increases with increasing pH (Bittel and Miller, 1974; Braids et al., 1972; Griffin and 

Shimp, 1976; Haji-Djafari et al., 1981; Hildebrand and Blum, 1974; Overstreet and 

Krishnamurthy, 1950; Scrudato 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 and sediments are known to adsorb lead 

(Rickard and Nriagu, 1978; Zimdahl and Hassett, 1977). Additionally, soluble organic matter 

such as fulvates and amino acids are known to chelate soluble lead and affect its adsorption on 

soils (Rickard and Nriagu, 1978). Gerritse et al. (1982) examined the lead adsorption properties 

of soils as a function of organic matter content of soils. Initial lead concentrations used in these 

experiments ranged from 0.001 to 0.1 mg/l. Based on adsorption data, the investigators 

expressed K d value for a soil as a function of organic matter content (as wt.%) and the distribution 

coefficient of the organic matter. The data also indicated that irrespective of soil organic matter 

content, lead adsorption increased with increasing soil pH (from 4 to 8). In certain soils, lead is 

F.3

also known to form methyl- lead complexes (Rickard and Nriagu, 1978). However, quantitative 

relationship between the redox status of soils and its effect on overall lead adsorption due to 

methylation of lead species is not known. 

Tso (1970), and Sheppard et al. (1989) studied the retention of 210 Pb in soils and its uptake by 

plants. These investigators found that lead in trace concentrations was strongly retained on soils 

(high K d values). Lead adsorption by a subsurface soil sample from Hanford, Washington was 

investigated by Rhoads et al. (1992). Adsorption data from these experiments showed that K d 

values increased with decreasing lead concentrations in solution (from 0.2 mg/l to 0.0062 mg/l). 

At a fixed pH of 8.35, the authors found that K d values were log-linearly correlated with 

equilibrium concentrations of lead in solution. Calculations showed that if lead concentrations 

exceeded about 0.207 mg/l, lead-hydroxycarbonate (hydrocerussite) would probably precipitate in 

this soil. 

The K d data described above are listed in Table F.1. 

F.2.0 Approach 

The initial step in developing a look-up table consisted of identifying the key parameters which 

were correlated with lead adsorption (K d values) on soils and sediments. Data sets developed by 

Gerritse et al. (1982) and Rhoads et al. (1992) containing both soil pH and equilibrium lead 

concentrations as independent variables were selected to develop regression relationships with K d 

as the dependent variable. From these data it was found that a polynomial relationship existed 

between K d values and soil pH measurements. This relationship (Figure F.1) with a correlation 

coefficient of 0.971 (r 2 ) could be expressed as: 

K d (ml/g) = 1639 - 902.4(pH) + 150.4(pH) 2 

The relationship between equilibrium concentrations of lead and K d values for a Hanford soil at a 

fixed pH was expressed by Rhoads et al. (1992) as: 

K d (ml/g) = 9,550 C -0.335 

where C is the equilibrium concentration of lead in µg/l. The look-up table (Table F.2) was 

developed from using the relationships F.1 and F.2. Four equilibrium concentration and 3 pH 

categories were used to estimate the maximum and minimum K d values in each category. The 

relationship between the K d values and the 2 independent variables (pH and the equilibrium 

concentration) is shown as a 3-dimensional surface (Figure F.2). This graph illustrates that the 

highest K d values are encountered under conditions of high pH values and very low equilibrium 

lead concentrations and in contrast, the lowest K d values are encountered under lower pH and 

higher lead concentrations. The K d values listed in the look-up table encompasses the ranges of 

pH and lead concentrations normally encountered in surface and subsurface soils and sediments. 

F.4 

(F.1) 

(F.2)

Table F.1. Summary of K d values for lead adsorption on soils. 

Reference 

Kd (ml/g) Experimental 

Parameters 

pH CEC 

(meq/100g) 

Iron 

Oxide 

content 

(wt.%) 

Organic 

Carbon 

(wt.%) 

Soil Description Clay 

Content 

(wt.% ) 

Haji-Djafari et al., 1981 

-- 

-- 

-- 

-- 

20 

100 

1,500 

4,000 

-- 

-- 

-- 

-- 

2.0 

4.5 

5.75 

7.0 

-- 

-- 

-- 

-- 

-- 

-- 

-- 

-- 

-- 

-- 

-- 

-- 

Sediment, Split Rock 

Formation, Wyoming 

Gerritse et al. (1982) 

Batch Experiment 




280 

1295 

3,000 

4,000 

22 

22 

16 

16 

4.5 

5.0 

7.5 

8.0 

-- 

-- 

-- 

-- 

-- 

-- 

-- 

-- 

0 

0 

2 

2 

Sand (Soil C) 

Sand (Soil C) 

Sandy Loam (Soil D) 

Sandy Loam (Soil D) 

F.5 

Sheppard et al. (1989) 





21,000 

19 

30,000 

59,000 

17 

5.8 

120 

8.7 

7.3 

4.9 

5.5 

7.4 

-- 

-- 

-- 

-- 

-- 

-- 

-- 

-- 

15 

2 

Figure F.1. Correlative relationship between K d and pH. 

F.6

Figure F.2. Variation of K d as a function of pH and the equilibrium lead 


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/l) 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/l at a 

fixed pH value. Both these data sets were used for estimating the range of K d values for the range 

of pH and lead concentration values found in soils. 

Table F.2. Estimated range of K d values for lead as a function of soil pH, and 

equilibrium lead concentrations. 

Equilibrium Lead 

Concentration (µg/l) K d (ml/g) 

0.1 - 0.9 

1.0 - 9.9 

10 - 99.9 

100 - 200 

F.8 

Soil pH 

4.0 - 6.3 6.4 - 8.7 8.8 - 11.0 

Minimum 940 4,360 11,520 

Maximum 8,650 23,270 44,580 

Minimum 420 1,950 5,160 

Maximum 4,000 10,760 20,620 

Minimum 190 900 2,380 

Maximum 1,850 4,970 9,530 

Minimum 150 710 1,880 

Maximum 860 2,300 4,410

F.4.0 References 

Abd-Elfattah, A., and K. Wada. 1981. “Adsorption of Lead, Copper, Zinc, Cobalt, and 

Cadmium by Soils that Differ in Cation-Exchange Material.” Journal of Soil Science, 32:71- 

283. 

Bargar, J. R., G. E. Brown, Jr., and G. A. Parks. 1998. “Surface Complexation of Pb(II) at 

Oxide-Water Interface: III. XAFS Determination of Pb(II) and Pb(II)-Chloro Adsorption 

Complexes on Goethite and Alumina.” Geochimica et Cosmochimica Acta, 62(2):193-207. 

Bittel, J. R., and R. J. Miller. 1974. “Lead, Cadmium, and Calcium Selectivity Coefficients on 

Montmorillonite, Illite, and Kaolinite.” Journal of Environmental Quality, 3:250-253. 

Braids, O. C., F. J. Drone, R. Gadde, H. A. Laitenen, and J. E. Bittel. 1972. “Movement of Lead 

in Soil-Water System.” In Environmental Pollution of Lead and Other Metals. pp 164-238, 

University of Illinois, Urbana, Illinois. 

Chow, T. J. 1978. “Lead in Natural Waters.” In The Biogeochemistry of Lead in the 

Environment. Part A. Ecological Cycles., J. O. Nriagu (ed.), pp. 185-218, Elsevier/North 

Holland, New York, New York. 

Forbes, E. A., A. M. Posner, and J. P. Quirk. 1976. “The Specific Adsorption of Cd, Co, Cu, 

Pb, and Zn on Goethite.” Journal of Soil Science, 27:154-166. 

Gerritse, R. G., R. Vriesema, J. W. Dalenberg, and H. P. De Roos. 1982. “Effect of Sewage 

Sludge on Trace Element Mobility in Soils.” Journal of Environmental Quality, 11:359-364. 

Grasselly, G., and M. Hetenyi. 1971. “The Role of Manganese Minerals in the Migration of 

Elements.” Society of Mining Geology of Japan, Special Issue 3:474-477. 

Griffin, R. A., and N. F. Shimp. 1976. “Effect of pH on Exchange-Adsorption or Precipitation of 

Lead from Landfill Leachates by Clay Minerals.” Environmental Science and Technology, 

10:1256-1261. 

Haji-Djafari, S., P. E. Antommaria, and H. L. Crouse. 1981. “Attenuation of Radionuclides and 

Toxic Elements by In Situ Soils at a Uranium Tailings Pond in central Wyoming.” In 

Permeability and Groundwater Contaminant Transport, T. F. Zimmie, and C. O. Riggs 

(eds.), pp 221-242. ASTM STP 746. American Society of Testing Materials. Washington, 

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. Altman. 1980. 

Adsorption/Coprecipitation of Trace Elements from Water with Iron Oxyhydroxides. 

EPRI-RP-910, Electric Power Research Institute, Palo Alto, California. 

Overstreet, R., and C. Krishnamurthy. 1950. “An Experimental Evaluation of Ion-exchange 

Relationships. ” Soil Science, 69:41-50. 

Peters, R. W., and L. Shem. 1992. “Adsorption/Desorption Characteristics of Lead on Various 

Types of Soil.” Environmental Progress, 11:234-240. 

Rhoads, K., B. N. Bjornstad, R. E. Lewis, S. S. Teel, K. J. Cantrell, R. J. Serne, J. L. Smoot, C. 

T. Kincaid, and S. K. Wurstner. 1992. Estimation of the Release 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. 


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 210 Po and 210 Pb 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 


F.10

APPENDIX G 

Partition Coefficients For Plutonium

G.1.0 Background 

Appendix G 

Partition Coefficients For Plutonium 

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 K d 

values for plutonium typically range over 4 orders of magnitude (Thibault et al., 1990). Also, 

from these data a number of factors which influence the adsorption behavior of plutonium have 

been identified. These factors and their effects on plutonium adsorption on soils and sediments 

were used as the basis for generating a look-up table. These factors are: 

C Typically, in many experiments, the oxidation state of plutonium in solution was not 

determined or controlled therefore it would be inappropriate to compare the K d data 

obtained from different investigations. 

C 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 K d values would reflect mainly precipitation reactions and 

not adsorption reactions. 

C Adsorption data show that the presence of ligands influence plutonium adsorption onto 

soils. Increasing concentrations of ligands decrease plutonium adsorption. 

C If no complexing ligands are present plutonium adsorption increases with increasing pH 

(between 5.5 and 9.0). 

C Plutonium is known to adsorb onto soil components such as aluminum and iron oxides, 

hydroxides, oxyhydroxides, and clay minerals. However, the relationship between the 

amounts of these components in soils and the measured adsorption of plutonium has not 

been quantified. 

Because plutonium in nature can exist in multiple oxidation states (III, IV, V, and VI), soil redox 

potential would influence the plutonium redox state and its adsorption on soils. However, our 

literature review found no plutonium adsorption studies which included soil redox potential as a 

variable. Studies conducted by Nelson et al. (1987) and Choppin and Morse (1987) indicated 

that the oxidation state of dissolved plutonium under natural conditions depended on the colloidal 

organic carbon content in the system. Additionally, Nelson et al (1987) also showed that 

plutonium precipitation occurred if the solution concentration exceeded 10 -7 M. 

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 et al. (1975) showed that plutonium adsorption onto mineral surfaces was 

influenced significantly by the type of mineral, the pH and mineral particle size. The reported 

values ranged from zero for quartz (Tamura, 1972) to 4,990 ml/g for montmorillonite (Evans, 

1956). [The K d 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 K d 

defemination was below detection.] These K d 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 K d values cannot be K d 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 K d values between hexavalent 

(250 ml/g) and tetravalent (21,000 ml/g) plutonium species adsorbing on to montmorillonite. 

Bondietti et al. (1975) also demonstrated that natural dissolved organic matter (fulvic acid) 

reduces plutonium from hexavalent to tetravalent state thus potentially affecting plutonium 

adsorption in natural systems. Some of the earlier adsorption experiments also demonstrated that 

complexation of plutonium by various ligands significantly influences its adsorption behavior. 

Increasing concentrations of acetate (Rhodes, 1957) and oxalate (Bensen, 1960) ligands resulted 

in decreasing adsorption of plutonium. Adsorption experiments conducted more recently 

(Sanchez et al., 1985) indicate that increasing concentrations of carbonate ligand also depresses 

the plutonium adsorption on various mineral surfaces. 

Even though the adsorption behavior of plutonium on soil minerals such as glauconite (Evans, 

1956), montmorillonite (Billon, 1982; Bondietti et al., 1975), attapulgite (Billon, 1982), and 

oxides, hydroxides, and oxyhydroxides (Evans, 1956; Charyulu et al., 1991; Sanchez et al., 

1985; Tamura, 1972; Ticknor, 1993; Van Dalen et al., 1975) has been studied, correlative 

relationships between the type and quantities of soil minerals in soils and the overall plutonium 

adsorption behavior of the soils have not been established. 

Adsorption experiments conducted by Billon (1982) indicated K d 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 -6 to 4x10 -6 M of Pu(IV)] 

used in these experiments, it is likely that precipitation and not adsorption resulted in very high K d 

values. Additional experiments conducted with Pu(VI) species on bentonite substrate resulted in 

K d 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

icarbonate concentrations (8.2 x 10 -6 to 2.9 x 10 -4 M)] resulted in distribution coefficients, K d, 

ranging from 170 to 1,400 ml/g. According to Pius et al. (1995), significant removal of Pu(IV) 

from solutions containing 0.1 to 1 M concentrations of sodium carbonate was observed with 

alumina, silica gel, and hydrous titanium oxide as substrates. These investigators also noted that 

the presence of carbonate lowered the sorption distribution coefficient for these adsorbents. 

However, even at 0.5 M carbonate, the coefficients were 60 ml/g, 1,300 ml/g, and 15,000 ml/g, 

respectively, for alumina, silica gel, and hydrous titanium oxide. In another study using 

bicarbonate solutions, the distribution coefficient for Pu(IV) sorption on alumina was lowered to 

about 30 ml/g at 0.5 M bicarbonate (Charyulu 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 PuO 2·xH 2O 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 K d 

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 K d 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 K d values, which showed that plutonium precipitation was not significant in these tests. Also, 

the initial plutonium concentrations were below the measured solubility limits of plutonium 

hydroxide. This experiment demonstrated that in carbonate-free systems, plutonium would be 

adsorbed on alumina substrates. 

Another study of adsorption of Pu(IV) and Pu(V) on goethite was conducted by Sanchez et al. 

(1985). The experimental conditions used by these investigators were evaluated for assessing 

whether the reaction being studied was indeed adsorption. The initial plutonium concentrations 

used in their experiments were 10 -10 and 10 -11 moles per liter. These concentrations are well 

below the equilibrium saturation levels for PuO 2·xH 2O. The equilibrating solutions used in these 

experiments contained salts such as NaNO 3, NaCl, Na 2SO 4, and NaHCO 3 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 NaNO 3 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(IV) and Pu(V) adsorbed onto the 

ionized hydroxyl sites in the form of free ions and their hydrolytic species with metal ion and the 

metal-ion part of the complexes adsorbing onto the surface. The adsorption isotherms obtained at 

the higher initial concentration (10 -10 M) of total soluble Pu(IV) and Pu(V) showed that the 

adsorption edges (pH value at which 50 percent adsorption occurs) increased towards a higher 

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(IV) 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/l) 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/l greatly decreased the adsorption of plutonium in both oxidation 

states. These data demonstrated that practically no Pu(IV) or Pu(V) adsorption occurred on 

goethite when the total carbonate concentration approached 1,000 meq/l (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/l) typically encountered in soils, adsorption of Pu(IV) increased with 

increasing dissolved carbonate concentration. These results indicate that Pu(IV) in these soils 

3+ 

may adsorb in the form of PuHCO3 species. 

Such complete suppression of Pu(IV) and Pu(V) adsorption was attributed to the presence of 

anionic plutonium-hydroxy carbonate species in solution and to the fact that goethite at this pH 

contains mainly negatively charged sites that have negligible affinity to adsorb anionic species. 

This adsorption behavior of Pu(IV) and Pu(V) in the presence of carbonate ions that form strong 

hydroxy carbonate complexes is typical of ligand-like adsorption of metal ions described by 

Benjamin and Leckie (1981). Ligand-like adsorption is described as adsorption of a metal-ligand 

complex that is analogous to adsorption of the free ligand species. Also, the metal-ligand 

complexes may not adsorb at all if these complexes are highly stable. These data clearly 

demonstrate that increasing total carbonate and hydroxyl solution concentrations significantly 

decrease Pu(IV) and Pu(V) on iron oxyhydroxide surfaces. 

Similar suppression of adsorption of higher valence state actinides in the presence of carbonate 

and hydroxyl ions has been observed by a number of investigators. Some of these studies include 

adsorption of U(VI) on goethite (Hsi and Langmuir, 1985; Koehler et al., 1992; Tripathi, 1984), 

ferrihydrite (Payne et al., 1992), and clinoptilolite (Pabalan and Turner, 1992), and Np(V) 

adsorption on ferrihydrite, hematite, and kaolinite (Koehler et al., 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 

6.8x10 -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 K d ranged from 10,000 

ml/g, ~100 to ~10,000 ml/g, and

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 K d values from about 1 order (for silt loam) to 

2 orders (for sand) of magnitude. Using a stronger chelating agent (10 -3 M DTPA) resulted in 

very low K d 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 K d values were 1 to 2 orders of magnitude higher than the values calculated from 

adsorption experiments leading to the conclusion that some fraction of plutonium in soil was 

specifically adsorbed (not exchangeable). These data showed that Pu(IV) adsorption on soils 

would be significantly reduced if the equilibrating solutions contain strong chelating ligands, such 

as EDTA and DTPA. 

The reduction of plutonium adsorption on soils by strong synthetic chelating agents was also 

confirmed by experiments conducted by Delegard et al. (1984). These investigators conducted 

tests to identify tank waste components that could significantly affect sorption of plutonium on 

3 typical shallow sediments from the the DOE Hanford Site. They found that sorption was 

decreased by the chelating agents, 0.05 M EDTA and 0.1 M HEDTA 

(N-2-hydroxyethylethylenediaminetriacetate) but not by low concentrations of carbonate 

(0.05 M). Delegard’s data also showed that roughly a twofold increase in ionic strength caused 

an order of magnitude decrease in plutonium adsorption. 

Based on an adsorption study of plutonium on basalt interbed sediments from the vicinity of 

Hanford site, Barney (1984) reported a K d value of about 500 ml/g. This relatively lower K d 

value may have resulted from the relatively enhanced concentration of 215 mg/l of carbonate 

(a complex forming ligand) which was present in the groundwater used in the experiments. 

Later, sorption of plutonium in +4, +5, and +6 redox states on a Hanford Site shallow sediment 

was studied by Barney (1992) to elucidate any differences in rate and amount of adsorption of 

plutonium in different redox states. The initial plutonium concentrations used in these 

experiments varied between about 10 -11 to 10 -9 M with synthetic ground water as a background 

electrolyte. The data indicated that the K d values ranged from 2,100 to 11,600, 2,700 to 4,600, 

and 1,000 to 4,600 ml/g for plutonium in +4, +5, and +6 redox states, respectively. The data also 

indicated that Pu(V) and Pu(VI) upon adsorption was reduced to the tetravalent state. In these 

experiments, the K d data obtained at lower initial concentrations (~1x10 -11 M) of plutonium are 

reliable because the dominant plutonium removal mechanism from solution was adsorption. 

Using batch equilibration techniques, Bell and Bates (1988) measured K d 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 K d values for plutonium adsorption on river, oceanic, and lake 

sediments range from about 1x10 3 to 1x10 6 ml/g. Duursma and coworkers calculated that K d for 

marine sediments was about 1x10 4 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 K d values from 1.64x10 4 to 

3.85x 10 5 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 K d was ~2.5x10 3 ml/g, and ~2.8x10 6 ml/g for plutonium in reduced states (III, IV). Based on a 

number of observations of lake and sea water samples, Nelson et al (1987) reported that K d values 

for lake particulates ranged from 3,000 to 4x10 5 ml/g, and for oceanic particulates ranged from 

1x10 5 to 4x10 5 ml/g. 

G.2.0 Data Set for Soils 

The most detailed data set on plutonium K d 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 10 -6 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 K d values and the 

measured soil parameters. The data set generated at initial plutonium concentrations of 10 -8 M 

were chosen for statistical analyses because the data sets obtained at higher initial concentrations 

of plutonium may have been affected by precipitation reactions (Table G.1). 

G.3.0 Approach and Regression Models 

The most detailed data set on plutonium K d 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 10 -6 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 K d 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 (K d in ml/g), pH, CEC (in meq/100g), electrical 

conductivity of soil extract (EC in mmhos/cm), dissolved carbonate concentration in soil extract 

(DCARB in meq/l), inorganic carbon content (IC as percent CaCO 3), organic carbon content 

(OC as wt.%), and the clay content (CLAY as wt.%). 

G.9

Table G.1. Plutonium adsorption data for soil samples. [Data taken from results 

reported by Glover et al. (1976) for measurements conducted at an initial 

plutonium concentrations of 10 -8 M.] 

Soil 

Sample 

K d 

(ml/g) 

pH CEC 1 

(meq/100 g) 

EC 1 

(mmhos/cm) 

G.10 

DCARB 1 

(meq/l) 

IC % 1 

CaCO3 

OC 1 

(% 

mass) 

CO-A 2,200 5.7 20.0 3.6 5.97 0.4 2.4 36 

CO-B 200 5.6 17.5 0.4 0.97 0.3 3.4 22 

CO-C 1,900 7.9 29.6 0.4 1.98 2.4 0.7 64 

ID-A 1,700 7.8 15.5 0.5 2.71 17.2 0.8 34 

ID-B 320 8.3 13.8 0.8 2.51 7.9 0.2 32 

ID-C 690 8.0 8.2 1.0 2.52 5.2 0.3 23 

ID-D 2,100 7.5 17.5 1.2 4.90 0.0 0.1 3 

WA-A 100 8.0 6.4 0.9 2.60 0.6 0.3 14 

WA-B 430 8.2 5.8 0.4 2.30 0.0 0.1 14 

SC 280 5.4 2.9 0.4 0.50 0.2 0.7 20 

NY 810 5.4 16.0 1.2 1.40 0.0 2.7 36 

NM 100 6.4 7.0 1.7 2.80 0.2 0.7 18 

AR-A 710 6.2 34.4 0.5 0.10 0.9 3.2 56 

AR-B 80 4.8 3.8 0.4 0.10 0.7 0.6 9 

AR-C 430 2.3 16.2 0.3 0.10 0.6 2.3 37 

IL 230 3.6 17.4 0.5 0.10 0.7 3.6 16 

CLAY 1 

( % 

mass) 

1 CEC: Cation exchange capacity; EC: Electrical conductivity; DCARB: Dissolved carbonate; 

IC: Inorganic carbon; OC: Organic carbon; CLAY: Soil clay content. 

.

The scatterplots are typically displayed in a matrix format with columns and rows representing the 

dependent and independent variables respectively. For instance, the first row of plots shows the 

relationship between K d as a dependent variable and other variables each in turn as selected as 

independent variables. Additionally, histograms displayed in each row illustrate the value 

distribution of each variable when it is being considered as the dependent variable. 

The scatter matrix (Figure G.1) shows that regression relationships may exist between K d 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.11

KD 

PH 

CEC 

EC 

G.12 

DCARB 

Figure G.1. Scatter plot matrix of soil properties and the distribution coefficient (K d) of 

plutonium. 

The variable selection with and without an intercept indicated that the 2 most significant variables 

for reliably forecasting the K d 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 K d 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 K d values are encountered under conditions of high clay content and 

dissolved carbonate concentrations. In contrast, the low K d 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. 

IC 

OC 

CLAY

Table G.2. Regression models for plutonium adsorption. 

Model Type Forecasting Equation R 2 

Linear Regression 

Forward Stepwise 


Forward Stepwise 


Backward Stepwise 


Backward Stepwise 

Piecewise Linear 

Regression 

K d = 284.6 (DCARB) + 27.8 (CLAY) - 594.2 0.7305 

K d = 488.3 (DCARB) + 29.9 (CLAY) - 119.1 (pH) - 356.8 (EC) 0.8930 

K d = 284.6 (DCARB) + 27.8 (CLAY) - 594.2 0.7305 

K d = 351.4 (DCARB) 0.7113 

K d = 25.7 (DCARB) + 12.14 (CLAY) + 2.41 for K d values 767.5 

Polynomial K d = -156.0 (DCARB) + 15.2 (CLAY) +16.1 (DCARB) 2 - 0.04 (CLAY) 2 + 11.3 (DCARB)(CLAY) - 87.0 0.9222 

Polynomial K d = -171.1(DCARB) + 10.5 (CLAY) +17.2(DCARB) 2 + 0.02 (CLAY) 2 + 11.6 (DCARB)(CLAY) 0.9219 

Polynomial K d = -106.1(DCARB) + 11.2 (CLAY) + 12.5 (DCARB)(CLAY) - 72.4 0.9194 

Polynomial K d = -137.9 (DCARB) + 9.3 (CLAY) + 13.4 (DCARB)(CLAY) 0.9190 

Table G.3. Estimated range of K d values for plutonium as a function of the soluble 

carbonate and soil clay content values. 

K d (ml/g) 

Clay Content (wt.%) 

0 - 30 31 - 50 51 - 70 


(meq/l) 

G.13 


(meq/l) 


(meq/l) 

0.1 - 2 3 - 4 5 - 6 0.1 - 2 3 - 4 5 - 6 0.1 - 2 3 - 4 5 - 6 

Minimum 5 80 130 380 1,440 2,010 620 1,860 2,440 

Maximum 420 470 520 1,560 2,130 2,700 1,980 2,550 3,130 

0.9730

Figure G.2. Variation of K d for plutonium as a function of clay content and 

dissolved carbonate concentrations. 

G.14

G.4.0 References 

Barney, G. S. 1984. “Radionuclide Sorption and Desorption Reactions with Interbed Materials 

from the Columbia River Basalt Formation.” In Geochemical Behavior of Radioactive Waste, 

G. S. Barney, J. D. Navratil, and W. W. Schulz (eds.), pp. 1-23. American Chemical Society, 


Barney, G. S. 1992. Adsorption of Plutonium on Shallow Sediments at the Hanford Site, 

WHC-SA-1516-FP, Westinghouse Hanford Company, Richland, Washington. 

Bell, J., and T. H. Bates. 1988. “Distribution coefficients of Radionuclides between Soils and 

Groundwaters and their Dependence on Various test Parameters.” Science of Total 

Environment, 69:297-317. 

Benjamin, M. M., and J. O. Leckie. 1981. “Conceptual Model for Metal-Ligand-Surface 

Interactions during Adsorption.” Environmental Science and Technology, 15:1050-1056. 

Bensen, D. W. 1960. Review of Soil Chemistry Research at Hanford. HW-67201. General 

Electric Company, Richland, Washington. 

Billon, A. 1982. “Fixation D’elements Transuraniens a Differents Degres D’oxydation Sur Les 

Argiles.” In Migration in the Terrestrial Environment of Long-lived Radionuclides from the 

Nuclear Fuel Cycle, pp. 167-176, IAEA-SM-257/32. International Atomic Energy Agency. 

Vienna, Austria. 

Bondietti, E. A., S. A. Reynolds, and M. H. Shanks. 1975. “Interaction of Plutonium with 

Complexing Substances in Soils and Natural Waters.” In Transuranium Nuclides in the 

Environment, pp. 273-287, IAEA-SM-199/51. International Atomic Energy Agency. 

Vienna. 

Charyulu, M. M., I. C. Pius, A. Kadam, M. Ray, C. K. Sivaramakrishnan, and S. K. Patil. 1991. 

“The Behavior of Plutonium in Aqueous Basic Media.” Journal of Radioanalytical and 

Nuclear Chemistry, 152: 479-485. 

Choppin, G. R., and J. W. Morse. 1987. “Laboratory Studies of Actinides in Marine Systems.” 

In Environmental Research on Actinide Elements, J. E. Pinder, J. J. Alberts, K. W. McLeod, 

and R. Gene Schreckhise (eds.), pp. 49-72, CONF-841142, Office of Scientific and Technical 

Information, U. S. Department of Energy, Washington, D.C. 

Dahlman, R. C., E. A. Bondietti, and L. D. Eyman. 1976. “Biological Pathways and Chemical 

Behavior of Plutonium and Other Actinides in the Environment.” In Actinides in the 

Environment, A. M. Friedman (ed.), pp. 47-80. ACS Symposium Series 35, American 


G.15

Delegard, C. H. , G. S. Barney, and S. A. Gallagher. 1984. “Effects of Hanford High-Level 

Waste Components on the Solubility and Sorption of Cobalt, Strontium, Neptunium, 

Plutonium, and Americium.” In Geochemical Behavior of Disposed Radioactive Waste, 

G. S. Barney, J. D. Navratil, and W. W. Schulz (eds.), pp. 95-112. ACS Symposium 

Series 246, American Chemical Society, Washington, D.C. 

Duursma, E. K., and M. G. Gross. 1971. “Marine Sediments and Radioactivity.” In 

Radioactivity in the Marine Environment, pp. 147-160, National Academy of Sciences, 


Duursma, E. K., and D. Eisma. 1973. “Theoretical, Experimental and Field Studies Concerning 

Reactions of Radioisotopes with Sediments and Suspended Particles of the Sea. Part C: 

Applications to Field Studies.” Netherlands Journal of Sea Research, 6:265-324. 

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, 

IAEA-163. International Atomic Energy Agency, Vienna, Austria. 

Evans, E. J. 1956. Plutonium Retention in Chalk River Soil. CRHP-660. Chalk River 

Laboratory, Chalk River, Canada. 

Glover, P. A., F. J. Miner, and W. O. Polzer. 1976. “Plutonium and Americium Behavior in the 

Soil/Water Environment. I. Sorption of Plutonium and Americium by Soils.” In Proceedings 

of Actinide-Sediment Reactions Working Meeting, Seattle, Washington. pp. 225-254, 

BNWL-2117, Battelle Pacific Northwest Laboratories, Richland, Washington. 

Hsi, C. K. D., and D. Langmuir. 1985. “Adsorption of Uranyl onto Ferric Oxyhydroxides: 

Application of the Surface Complexation Site-Binding Model.” Geochimica et 


Koehler M., E.Wieland, and J. O. Leckie. 1992. “Metal-Ligand Interactions during Adsorption 

of Uranyl and Neptunyl on Oxides and Silicates.” In Proceedings of 7th International 

Symposium On Water-Rock Interaction -WRI7. V1: Low Temperature Environment, 

Y. K. Kharaka and A. S. Maest (eds.), A. A. Balkema, Rotterdam, Netherlands. 

Mo, T., and F. G. Lowman. 1975. “Laboratory Experiments on the Transfer Dynamics of 

Plutonium from Marine Sediments to Seawater and to Marine Organisms.” CONF-750503-5, 

Technical Information Center. U.S. Department of Energy, Washington, D.C. 

Montgomery, D. C., and E. A. Peck. 1982. Introduction to Linear Regression Analysis. John 

Wiley and Sons, New York, New York. 

G.16

Nelson, D. M., R. P. Larson, and W. R. Penrose. 1987. “Chemical Speciation of Plutonium in 

Natural Waters.” In Environmental Research on Actinide Elements, J. E. Pinder, J. J. 

Alberts, K. W. McLeod, and R. Gene Schreckhise (eds.), pp. 27-48, CONF-841142, Office of 

Scientific and Technical Information, U.S. Department of Energy, Washington, D.C. 

Nishita, H. 1978. “Extractability of Plutonium-238 and Curium-242 from a Contaminated Soil as 

a Function of pH and Certain Soil Components. CH 3COOH-NH 4OH System.” In 

Environmental Chemistry and Cycling Process, pp. 403-416. CONF-760429, Technical 

Information Center, U.S. Department of Energy, Washington, D.C. 

Nishita, H., M. Hamilton, and A. J. Steen. 1976. “Extractability of Pu-238 and Cm-242 from a 

Contaminated soil as a Function of pH and Certain Soil Components.” Soil Science Society 

of America Abstracts, Madison, Wisconsin. 

Pabalan, R. T., and D. R. Turner. 1992. Sorption Modeling for HLW Performance Assessment. 

Re. On Res. Act. For Calender Year 1991, W. C. Patrick (ed.), pp. 8-1 to 8-66. CNWRA 

91-01A. Center for Nuclear Waste Regulations and Analysis, San Antonio, Texas. 

Payne T. E., K. Sekine, J. A. Davis, and T. D. Waite. 1992. “Modeling of Radionuclide Sorption 

Processes in the Weathered Zone of the Koongarra Ore Body.” In Alligator Rivers Analogue 

Project Annual Report, 1990-1991, P. Duerden (ed.), pp. 57-85. Australian Nuclear Science 

and Technical Organization, Australia. 

Pius, I. C., M. M. Charyulu, B. Venkataramani, C. K. Sivaramakrishnan, and S. K. Patil. 1995. 

“Studies on Sorption of Plutonium on Inorganic Ion Exchangers from Sodium Carbonate 

Medium.” Journal of Radioanalytical and Nuclear Chemistry Letters, 199:1-7. 

Prout, W. E. 1958. “Adsorption of Radioactive Wastes by Savannah River Plant Soil.” Soil 

Science, 13-17. 

Relyea, J. F., and D. A. Brown. 1978. “Adsorption and Diffusion of Plutonium in Soil.” In 

Environmental Chemistry and Cycling Process, CONF-760429. Technical Information 

Center, U.S. Department of Energy, Washington, D.C. 

Rhodes, D. W. 1957. “The Effect of pH on the Uptake of Radioactive Isotopes from Solution by 

a Soil.” Soil Science Society of America Proceedings, 21:389-392. 

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. 


Handbook 60, U.S. Department of Agriculture, Washington, D.C. 

G.17

Rodgers, D. R. 1976. “Behavior of Plutonium-238 Solutions in the Soil and Hydrology System 

at Mound Laboratory.” In Proceedings of Actinide-sediment Reactions Working Meeting, 

Seattle, Washington, pp. 291-497. BNWL-2117, Battelle Pacific Northwest Laboratories, 


Sanchez, A. L., J. W. Murray, and T. H. Sibley. 1985. “The Adsorption of Pu (IV) and (V) of 

Goethite.” Geochimica et Cosmochimica Acta, 49:2297-2307. 



Tamura T. 1972. “Sorption Phenomena Significant in Radioactive Waste Disposal.” In 

Underground Waste Management and Environmental Implications, pp. 318-330. American 

Association of Petroleum Geology Memoirs 18, Tulsa, Oklahoma. 



AECL-10125, Whiteshell Nuclear research Establishment, Pinawa, Canada. 

Ticknor, K. V. 1993. “Actinide Sorption by Fracture-Filling Minerals.” Radiochimica Acta, 

60:33-42. 

Tripathi, V. S. 1984. Uranium (VI) Transport Modeling: Geochemical Data and Submodels. 


Van Dalen, A., F. DeWitte, and J. Wikstra. 1975. Distribution Coefficients for Some 

Radionuclides Between Saline Water and Clays, Sandstones and Other Samples from Dutch 

Subsoil, Report 75-109, Reactor Centrum, Netherlands. 

G.18

APPENDIX H 

Partition Coefficients For Strontium

H.1.0 Background 

Appendix H 

Partition Coefficients For Strontium 

Two simplifying assumptions underlying the selection of strontium K d values included in the lookup 

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/l, ionic strengths greater than about 0.1 M, and 

pH levels greater than approximately 12. 

Based on these assumptions and limitations, strontium K d values and some important ancillary 

parameters that influence cation exchange were collected from the literature and tabulated in 

Section H.3. The tabulated data were from studies that reported K d values (not percent adsorbed 

or Freundlich or Langmuir constants) and were conducted in systems consisting of 

C Natural soils (as opposed to pure mineral phases) 


C pH values between 4 and 10 

C Strontium concentrations less than 10 -4 M 


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 tuff 1 soil collected at Yucca Mountain, Nevada 

(Sample YM-38; Vine et al., 1980). The average strontium K d value was 355 ± 184 ml/g. The 

median 2 strontium K d value was 15.0 ml/g. This is perhaps the single central estimate of a 

strontium K d value for this data set. 

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. 

Table H.1. Descriptive statistics of strontium K d data set for soils. 

Sr K d 

(ml/g) 

Clay 

Content 

(wt.%) 

H.3 

pH CEC 

(meq/100 g) 

Surface 

Area 

(m 2 /g) 

Ca 

(mg/l) 

Mean 355 7.1 6.8 4.97 1.4 56 

Standard Error 183 1.1 0.21 1.21 0 23 

Median 15 5 6.7 0.9 1.4 0 

Mode 21 5 6.2 2 1.4 0 

Standard Deviation 1,458 7.85 1.35 9.66 0.00 134 

Kurtosis 34 10.7 -0.5 11.6 -3 3.4 

Minimum 1.6 0.5 3.6 0.05 1.4 0.00 

Maximum 10,200 42.4 9.2 54 1.4 400 

Number of 

Observations 

63 48 42 63 7.00 32

H.2.0 Approach and Regression Models 

H.2.1 Correlations with Strontium K d Values 

A matrix of the correlation coefficients of the strontium K d 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 K d values was with CEC (r = 0.84). Also significant are the correlation coefficients 

between strontium K d values and clay content (r = 0.82) and CEC and clay content (r = 0.91) 

(Table H.2). 

H.2.2 Strontium K d Values as a Function of CEC and pH 

The CEC and strontium K d 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 K d values are reported ( Keren and O’Connor, 1983; McHenry, 1958; Serne et al., 

1993). The strontium K d 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 K d values may be expected at a given CEC level. 

Strontium K d 

Table H.2. Correlation coefficients (r) of the strontium K d data set for soils. 

Strontium 

K d 

1.00 

Clay Content 0.82 1 

Clay 

Content 

1.00 

pH 0.28 0.03 1.00 

CEC 0.84 1 

0.91 1 

pH CEC Surface 

Area 

0.28 1 

H.4 

1.00 

Surface Area 0.00 -1.00 0.00 1.00 1 

1.00 

Ca Conc. 

Ca Conc. -0.17 0.00 -0.20 0.03 --- 1.00 

1 Correlation coefficients significant at or less than the 5% level of probability (P # 0.05).

Figure H.1. Relation between strontium K d 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 d 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 K d values using 

these data for low CEC soils, such as sandy aquifers, may be especially inaccurate. 

The regression equation for the data in Figure H.1 is presented as Equation 1 in Table H.3. Also 

presented in Table H.3 are the 95 percent confidence limits of the calculated regression 

coefficients, the y-intercepts, and slopes. These coefficients, when used to calculate K d values, 

suggest a K d range at a given CEC by slightly over an order of magnitude. The lower 95 percent 

confidence limit coefficients can provide guidance in selecting lower (or conservative) K d values. 

The large negative intercept in Equation 1 compromises its value for predicting strontium K d 

values in low CEC soils, a potentially critical region of the data, because many aquifers matrix 

have low CEC values. At CEC values less than 2.2 meq/100 g, Equation 1 yields negative 

strontium K d values, which are clearly unrealistic. 1 To provide a better estimate of strontium K d 

1 A negative Kd value is physically possible and is indicative of the phenomena referred to as 

anion exclusion or negative adsorption. It is typically and commonly associated with anions being 

H.5

values at low CEC values, 2 approaches were evaluated. First, the data in Figure H.1 was 

reanalyzed such that the intercept of the regression equation was set to zero, i.e., the regression 

equation was forced through the origin. The statistics of the resulting regression analysis are 

presented as Equation 2 in Table H.3. The coefficient of determination (R 2 ) for Equation 2 

slightly decreased compared to Equation 1 to 0.67 and remained highly significant (F= 2x10 -16 ). 

However, the large value for the slope resulted in unrealistically high strontium K d values. For 

example at 1 meq/100 g, Equation 2 yields a strontium K d 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 K d 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 K d 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 K d Values as a Function of Clay Content and pH 

Because CEC data are not always available to contaminant transport modelers, an attempt was 

made to use independent variables in the regression analysis that are more commonly available to 

modelers. Multiple regression analysis was conducted using clay content and pH as independent 

variables to predict CEC (Equations 7 and 8 in Table H.3) and strontium K d 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 (R 2 = 0.86) and for those data limited to CEC less 

than 15 meq/100 g (R 2 = 0.57). Thus, it is not surprising that clay content and pH were 

correlated to strontium K d 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

Figure H.2. Relation between strontium K d values for soils with 

CEC values less than 15 meq/100 g. 

Figure H.3. Relation between strontium K d values and soil 

clay contents. 

H.7

Table H.3. Simple and multiple regression analysis results involving strontium K d values, 

cation exchange capacity (CEC; meq/100 g), pH, and clay content (percent). 

# Equation n 2 Data 

Range 3 

H.8 

95% Confidence Limits 1 

Intercept Slope First 

Independent 

Parameter 

Slope Second 

Independent 

Parameter 

Lower Upper Lower Upper Lower Upper 

R 2 4 F Value 5 

1 K d = -272 + 126(CEC) 63 All -501 -43 105 147 --- --- 0.70 1x10 -17 

2 K d = 114(CEC) 63 All --- --- 95 134 --- --- 0.67 2x10 -16 

3 K d = 10.0 + 4.05(CEC) 57 CEC

H.2.4 Approach 

Figure H.4. Relation between strontium K d values and soil pH. 

Two strontium K d look up tables were created. The first table requires knowledge of the CEC 

and pH of the system in order to select the appropriate strontium K d value (Table H.4). The 

second table requires knowledge of the clay content and pH to select the appropriate strontium K d 

value (Table H.5). 

A full factorial table was created that included 3 pH categories and 3 CEC categories. This 

resulted in 9 cells. Each cell contained a range for the estimated minimum- and maximum K d 

values. A 2 step process was used in selecting the appropriate K d values for each cell. For the 

first step, the appropriate equations in Table H.3 were used to calculate K d values. The lower and 

upper 95 percent confidence limit coefficients were used to provide guidance regarding the 

minimum and maximum K d values. For the 2 lowest CEC categories, Equation 6 in Table H.3 

was used. For the highest CEC category, Equation 5 was used. For the second step, these 

calculated values were adjusted by “eye balling the data” to agree with the data in Figures 

H.2-H.4. It is important to note that some of the look-up table categories did not have any actual 

observations, e.g., pH

Table H.4. Look-up table for estimated range of K d 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 (

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. 

Equation 1 Y-Intercept Slope CEC 

(meq/100 g) 

H.11 

Clay Content 

(%) 

12 --- 1.34 3 4 

12 --- 1.34 15 20 

11 3.36 1.1.2 15 20 

11 3.36 1.12 50 59 

1 Number of equation in Table H.3.

H.3.0 K d Data Set for Soils 

Table H.7 lists the available K d values identified for experiments conducted with only soils. The K d 

values are listed with ancillary parameters that included clay content, pH, CEC, surface area, solution 

calcium concentrations, and solution strontium concentrations. 

Sr K d 

(ml/g) 

Clay 

Content 

(%) 

pH CEC 

(meq/ 

100 g) 

Table H.7. Strontium K d data set for soils. 

Surface 

Area 

(m 2 /g) 

[Ca] 

ppm 

H.12 

[Sr] Background 

Solution 

Soil 

ID 

Reference 1 , Comments 

21 0.8 5.2 0.9 1.4 0 * NaClO 4 Soil A 1, * = 4.4e2Bq/ml 85-Sr in 

2.4x10 -8 M SrCl 2 


2.4x10 -8 M SrCl 2 


2.4x10 -8 M SrCl 2 


2.4x10 -8 M SrCl 2 


2.4x10 -8 M SrCl 2 


2.4x10 -8 M SrCl 2 


2.4x10 -8 M SrCl 2 

21.4 5 0.47 Groundwater 2 

25 5 0.83 Groundwater 2, CEC was estimated by 

adding exch. Ca,Mg,K 

12.7 5 0.39 Groundwater 2, GW = 7.4Ca, 1.7Mg, 

2.2Na,5.6Cl, 18ppmSO4 

7.9 5 0.46 Groundwater 2, Aquifer sediments 

15.6 5 0.81 Groundwater Chalk River Nat'l Lab, 

Ottawa, Canada 

9.4 5 0.21 Groundwater 2, Described as sand texture 

7.6 5 0.25 Groundwater 2, Assumed 5% clay, mean 

[clay] in sandy soils 



28.1 5 0.76 Groundwater 2

Sr K d 

(ml/g) 

Clay 

Content 

(%) 

pH CEC 

(meq/ 

100 g) 

Surface 

Area 

(m 2 /g) 

[Ca] 

ppm 


Solution 




13 5 0.25 Groundwater 2 








4 10 4 5.5 0 1x10 -8 M 0.01M NaCl Puye 

soil-Na 

15 10 5 5.5 0 1x10 -8 M 0.01M NaCl Puye 

soil-Na 

21 10 6 5.5 0 1x10 -8 M 0.01M NaCl Puye 

soil-Na 

24 10 7.4 5.5 0 1x10 -8 M 0.01M NaCl Puye 

soil-Na 

3 10 3.6 5.5 400 1x10 -8 M 0.01M CaCl Puye 

soil-Ca 

4.5 10 5.2 5.5 400 1x10 -8 M 0.01M CaCl Puye 

soil-Ca 

5.2 10 6.8 5.5 400 1x10 -8 M 0.01M CaCl Puye 

soil-Ca 

5.7 10 7.9 5.5 400 1x10 -8 M 0.01M CaCl Puye 

soil-Ca 

3.5 5.2 2 0 1x10 -10 M NaOH/HCl Hanford 

soil 


soil 


soil 


soil 

H.13 

Soil 

ID 


3 

3, Noncalcareous soils 

3 

3 

3 

3 

3 

3 

4 

4, Carbonate system 

4 

4

Sr K d 

(ml/g) 

Clay 

Content 

(%) 

pH CEC 

(meq/ 

100 g) 

Surface 

Area 

(m 2 /g) 

[Ca] 

ppm 


Solution 

8.3 6 2 0 1x10 -10 M NaOH/HCl Hanford 

soil 

17 7.4 2 0 1x10 -10 M NaOH/HCl Hanford 

soil 


soil 


soil 


soil 


soil 

19.1 4 7.66 10.4 129 100 

µCi/l 

21.5 6 7.87 5.9 58.5 100 

µCi/l 

23.2 5 8.17 4.57 35.1 100 

µCi/l 

48.5 8.24 3 3.8x10 - 

10,200 8.17 54 3.8x10 - 

2,500 8.13 21 3.8x10 - 

3,790 8.24 27 3.8x10 - 

3,820 8.24 27 3.8x10 - 

H.14 

8 M 

8 M 

8 M 

8 M 

8 M 

Hanford 

Groundwater 

Hanford 

Groundwater 

Hanford 

Groundwater 

Yucca 

Groundwater 

Yucca 

Groundwater 

Yucca 

Groundwater 

Yucca 

Groundwater 

Yucca 

Groundwater 

Soil 

ID 


4 

4 

4 

4 

4 

4 

cgs-1 5 

1.6 0.5 6.2 0.05 10x10 -6 M Groundwater Sediments 7 

trench-8 5, Groundwater pH = 8.3 

tbs-1 5, Hanford, Richland, 

Washington surface and 

subsurface sediments 

YM-22 6, Los Alamos, New Mexico 

YM-38 6, Yucca Mountain tuff 

sediments 

YM48 6, Approximate initial pH, 

final pH are presented 

YM-49 6, Final pH 8.1- 8.5 

YM-50 6, Sediments = 106-500 µm 

fractions 

2.6 3 6.2 0.3 10x10 -6 M Groundwater Sediments 7, Added kaolinite to sand 

3.4 5 6.2 0.5 10x10 -6 M Groundwater Sediments 7, CEC estimated based on 

kaolinite = 10 meq/100 g 

4.6 8 6.2 0.8 10x10 -6 M Groundwater Sediments 7 


400 42.4 7.2 34 0 Water Ringhold 

Soil 

8, soil from Richland, 

Washington

Sr K d 

(ml/g) 

Clay 

Content 

(%) 

pH CEC 

(meq/ 

100 g) 

Surface 

Area 

(m 2 /g) 

[Ca] 

ppm 


Solution 

135 26.9 8.3 13.6 0 Water Bowdoin 

Soil 

H.15 

Soil 

ID 


8, soil from Montana 

600 33.5 6.5 26.3 0 Water Hall soil 8, soil from Nebraska 

70 3.5 8.3 5.8 0 Water Composite 

Soil 

8, soil from Hanford Site, 

Richland, Washington 

1 References: 1 = Ohnuki, 1994, 2 = Patterson and Spoel, 1981; 3 = Keren and O'Connor, 1983; 4 = Rhodes and Nelson, 1957; 

5 = Serne et al., 1993; 6 = Vine et al., 1980; 7 = Lieser and Steinkopff, 1989; 8 = McHenry, 1958

H.4.0 K d Data Set for Pure Mineral Phases and Soils 

Table H.8 lists the available K d values identified for experiments conducted with pure mineral phases 

as well as soils. The K d values are listed with ancillary parameters that included clay content, pH, 

CEC, surface area, solution calcium concentrations, and solution strontium concentrations. 

Sr K d 

(ml/g) 

Clay 

Conten 

t (%) 

Table H.8. Strontium K d data set for pure mineral phases and soils. 

pH CEC 

(meq/ 

100 g) 

Surface 

Area 

(m 2 /g) 

[Ca] 

(ppm) 

H.16 


Solution 

Soil ID Reference 1 

and Comments 

21 0.8 5.2 0.9 1.4 0 * NaClO 4 Soil A 1, Ohnuki, 1994 

19 0.8 5.6 0.9 1.4 0 * NaClO 4 Soil A 1, * = 4.4x10 2 Bq/ml 85- 

Sr in 2.4x10 -8 M SrCl 2 











0 5.5 * Quartz 1, * = 4.4x10 2 Bq/ml 85- 


290 5.5 3.3 26.4 0 * Kaolinite 1, * = 4.4x10 2 Bq/ml 85- 


140 5.5 3.6 43.9 0 * Halloysite 1, * = 4.4x10 2 Bq/ml 85- 


17 5.5 0.6 1.4 0 * Chlorite 1, * = 4.4x10 2 Bq/ml 85- 


37 5.5 1.9 2.2 0 * Sericite 1, * = 4.4x10 2 Bq/ml 85- 


8 5.5 0.5 0.7 0 * Oligoclase 1, * = 4.4x10 2 Bq/ml 85- 


6 5.5 0.5 0 * Hornblend 1, * = 4.4x10 2 Bq/ml 85- 

Sr in 2.4x10 -8 M SrCl 2

Sr K d 

(ml/g) 

Clay 

Conten 

t (%) 

pH CEC 

(meq/ 

100 g) 

Surface 

Area 

(m 2 /g) 

[Ca] 

(ppm) 

H.17 


Solution 


and Comments 

16 5.5 0.7 0 * Pyroxene 1, * = 4.4x10 2 Bq/ml 85- 


110 5.5 8.5 19.3 0 * MnO 2 1, * = 4.4x10 2 Bq/ml 85- 


7.7 5.8 24 113 µCi/l Groundwater AA 45/1 2 Jackson and Inch, 1989 

9.9 6.1 25 105 µCi/l Groundwater AA45/3 2, K d = -.38Ca + 0.82. r2 

= 0.19 

12.6 6.1 23 105 µCi/l Groundwater AA45/4 2, Ca not important to Sr 

K d 

13.7 5.8 22 123 µCi/l Groundwater AA45/5 2 

10.1 6 24 99 µCi/l Groundwater AA45/7 2 



11 5.9 21 114 µCi/l Groundwater AA38/3 2 








6.8 6.2 84 µCi/l Groundwater AA27/4 2 









6.5 6.4 17 75 µCi/l Groundwater AA34/5 2

Sr K d 

(ml/g) 

Clay 

Conten 

t (%) 

pH CEC 

(meq/ 

100 g) 

Surface 

Area 

(m 2 /g) 

[Ca] 

(ppm) 


Solution 




H.18 


and Comments 

21.4 0.47 Groundwater 3 Patterson and Spoel, 

1981 

25 0.83 Groundwater 3, CEC was 

approximated by adding 

exch. Ca,Mg,K 

12.7 0.39 Groundwater 3, Groundwater =7.4 

ppm Ca, 1.7 ppm Mg, 2.2 

ppm Na, 5.6 ppm Cl, 18 

ppm SO 4 

7.9 0.46 Groundwater 3 










13 0.25 Groundwater 3 








4 10 4 5.5 0 1x10 -8 M .01M NaCl Puye 

soil-Na 

15 10 5 5.5 0 1x10 -8 M .01M NaCl 4, Noncalcareous soils 

4

Sr K d 

(ml/g) 

Clay 

Conten 

t (%) 

pH CEC 

(meq/ 

100 g) 

Surface 

Area 

(m 2 /g) 

[Ca] 

(ppm) 


Solution 

21 10 6 5.5 0 1x10 -8 M .01M NaCl 4 

24 10 7.4 5.5 0 1x10 -8 M .01M NaCl 4 

3 10 3.6 5.5 400 1x10 -8 M .01M CaCl 2 Puye 

soil-Ca 

4.5 10 5.2 5.5 400 1x10 -8 M .01M CaCl 2 4 

5.2 10 6.8 5.5 400 1x10 -8 M .01M CaCl 2 4 

5.7 10 7.9 5.5 400 1x10 -8 M .01M CaCl 2 4 

7.2 3 0 0.1 ppm 2,000 ppm 

Na 

12.7 5 0 0.1 ppm 2,000 ppm 

Na 

14.9 7 0 0.1 ppm 2,000 ppm 

Na 

12.9 9 0 0.1 ppm 2,000 ppm 

Na 

25.1 11 0 0.1 ppm 2,000 ppm 

Na 

H.19 


and Comments 

4 

Hanford Soil 5 





40.6 0.98 C-27 6 

48.6 0.96 C-27 6 

35 0.88 C-97 6 

39.2 0.8 C-55 6 

25.2 0.73 C-81 6 

16.4 0.39 C-62 6 

10.3 0.36 C-71 6 

8.2 0.32 C-85 6 

7.6 0.25 C-77 6 

7.8 0.51 MK-4 6 

11.2 0.38 TK3 6 

10.5 0.34 RK2 6 

3.7 0.34 NK2 6 

3.5 5.2 2 0 1x10 -10 M NaOH/HCl Hanford soil 7 

4.6 5.6 2 0 1x10 -10 M NaOH/HCl Hanford soil 7

Sr K d 

(ml/g) 

Clay 

Conten 

t (%) 

pH CEC 

(meq/ 

100 g) 

Surface 

Area 

(m 2 /g) 

[Ca] 

(ppm) 


Solution 



8.3 6 2 0 1x10 -10 M NaOH/HCl Hanford soil 7 

17 7.4 2 0 1x10 -10 M NaOH/HCl Hanford soil 7 





140 70 2.4 70 0 1x10 -8 M Water Bentonite 8 

160 70 2.4 70 1x10 -8 M Groundwater Bentonite 8 

1500 70 9.3 70 0 1x10 -8 M Water Bentonite 8 

1100 70 9.3 70 1x10 -8 M Groundwater Bentonite 8 

H.20 


and Comments 

1800 10 6.1 130 0 1x10 -8 M Water Takadate Loam 8, hydrohalloysite=10%, 

70% silt 

950 10 8 130 1x10 -8 M Groundwater Takadate Loam 8, hydrohalloysite=10%, 

70% silt 

550 10 6.5 60 0 1x10 -8 M Water Hachinohe 

Loam 

260 10 8.2 60 1x10 -8 M Groundwater Hachinohe 

Loam 

19.1 4 7.66 10.4 129 100 µCi/l Hanford 

Groundwater 

21.5 6 7.87 5.9 58.5 100 µCi/l Hanford 

Groundwater 

23.2 5 8.17 4.57 35.1 100 µCi/l Hanford 

Groundwater 

48.5 0 8.24 3 3.8x10 -8 M Yucca 

Groundwater 

10200 0 8.17 54 3.8x10 -8 M Yucca 

Groundwater 

2500 0 8.13 21 3.8x10 -8 M Yucca 

Groundwater 

3790 0 8.24 27 3.8x10 -8 M Yucca 

Groundwater 

cgs-1 9 

8, hydrohalloysite = 10%, 

90% silt 

8, hydrohalloysite = 10%, 

90% silt 

trench-8 9, Groundwater pH = 8.3 

tbs-1 9 

YM-22 10, Los Alamos, New 

Mexico 

YM-38 10, Yucca Mt tuff 

sediments 

YM48 10, Approximate initial 

pH, final pH are 

presented 

YM-49 10, Final pH 8.1- 8.5

Sr K d 

(ml/g) 

Clay 

Conten 

t (%) 

pH CEC 

(meq/ 

100 g) 

Surface 

Area 

(m 2 /g) 

[Ca] 

(ppm) 


Solution 

3820 0 8.24 27 3.8x10 -8 M Yucca 

Groundwater 

27000 0 8.4 31 10 3.8x10 -8 M Yucca 

Groundwater 

4850 0 8.63 31 50 3.8x10 -8 M Yucca 

Groundwater 

85 0 8.25 8 10 3.8x10 -8 M Yucca 

Groundwater 

17.7 0 8.5 8 50 3.8x10 -8 M Yucca 

Groundwater 

385 0 8.39 105 10 3.8x10 -8 M Yucca 

Groundwater 

149 0 8.45 105 50 3.8x10 -8 M Yucca 

Groundwater 

H.21 


and Comments 

YM-50 10, Sediments = 106-500 

µm fractions 

JA-18 10 

JA-19 10 

JA-32 10 

JA-33 10 

JA-37 10 

JA-38 10 

25000 12 10 nCi/ml kaolinite 13 

530 12 10 nCi/ml chlorite 13 

71,000 12 10 nCi/ml FeOOH 13 

1.6 0.5 6.2 0.05 10x10 -6 M Groundwater Sediments 14 

2.6 3 6.2 0.3 10x10 -6 M Groundwater Sediments 14, Added Kaolinite to 

sand 

3.4 5 6.2 0.5 10x10 -6 M Groundwater Sediments 14, CEC estimated based 

on kaolinite = 10 

meq/100 g 



17,000 97 1x10 -10 M Ohya tuff 14, Akiba and 

Hashimoto, 1990 

150 3.4 1x10 -10 M Pyrophyllite 14, log K d = log CEC + 

constant: for trace [Sr] 

780 2.4 1x10 -10 M Sandstone 14, pH not held constant, 

ranged from 6 to 9. 

95 1.9 1x10 -10 M Shale 14, 1g solid:50ml 

sol'n,centrifuged,32- 

60mesh 

440 1.9 1x10 -10 M Augite 

Andesite 

39 1.2 1x10 -10 M Plagiorhyolite 14 

14, CEC of Cs and K d of 

Sr

Sr K d 

(ml/g) 

Clay 

Conten 

t (%) 

pH CEC 

(meq/ 

100 g) 

Surface 

Area 

(m 2 /g) 

[Ca] 

(ppm) 


Solution 

380 0.75 1x10 -10 M Olivine Basalt 14 

50 0.57 1x10 -10 M Vitric Massive 

Tuff 

82 0.54 1x10 -10 M Inada granite 14 

22 0.35 1x10 -10 M Rokko Granite 14 

1.3 0.033 1x10 -10 M Limestone 14 

2,000 2 1x10 -10 M Muscovite 14 

140 0.93 1x10 -10 M Chlorite 14 

40 0.36 1x10 -10 M Hedenbergite 14 

20 0.33 1x10 -10 M Hornblende 14 

71 0.11 1x10 -10 M Grossular 14 

150 0.07 1x10 -10 M Microcline 14 

0.92 0.067 1x10 -10 M Forsterite 14 

14 0.034 1x10 -10 M K-Feldspar 14 

30 0.032 1x10 -10 M Albite 14 

3 0.022 1x10 -10 M Epidote 14 

23 0.0098 1x10 -10 M Quartz 14 

H.22 


and Comments 

400 42.4 7.2 34 0 Water Ringhold Soil 11, Soil from Richland 

WA 

135 26.9 8.3 13.6 0 Water Bowdoin Soil 11, from Montana 

600 33.5 6.5 26.3 0 Water Hall Soil 11, from Nebraska 

70 3.5 8.3 5.8 0 Water Composite Soil 11, from Hanford Site 

2.4 4 Groundwater Eolian Sand 12 

4.7 5 Eolian Sand 12, Belgian soils 

6 7 Eolian Sand 12, Composition of 

Groundwater was not 

given 

2.3 4 Mol White 

Sand 


Sand 


Sand 

14 

12, Compared static vs. 

dynamic Kd 

12 

12

Sr K d 

(ml/g) 

Clay 

Conten 

t (%) 

pH CEC 

(meq/ 

100 g) 

Surface 

Area 

(m 2 /g) 

[Ca] 

(ppm) 


Solution 

2.6 4 Mol Lignitic 

Sand 


Sand 


Sand 

H.23 


and Comments 

1 References: 1 = Ohnuki, 1994; 2 = Jackson and Inch ,1989; 3 =Patterson and Spoel ,1981; 4 = Keren and O'Connor, 1983; 5 Nelson, 

1959; 6 = Inch and Killey, 1987; 7 = Rhodes and Nelson, 1957; 8 = Konishi et al., 1988; 9 = Serne et al., 1993; 10 = Vine et al., 1980; 

11 = McHenry, 1958;12 = Baetsle et al., 1964; 13 = Ohnuki, 1991; 14 = Lieser and Steinkopff, 1989 

12 

12 

12

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 


Baetsle, L. H., P. Dejonghe, W. Maes, E. S. Simpson, J. Souffriau, and P. Staner. 1964. 

Underground Radionuclide Movement. EURAEC-703, European Atomic Energy 

Commission, Vienna, Austria. 

Cantrell, K., P. F. Martin, and J. E. Szecsody. 1994. “Clinoptilolite as an In-Situ Permeable 

Barrier to Strontium Migration in Ground Water.” In In-Situ Remediation: Scientific Basis 

for Current and Future Technologies. Part 2., G. W. Gee and N. Richard Wing (eds.). 

pp. 839-850. Battelle Press, Columbus, Ohio. 

Cui, D., and R. E. Eriksen. 1995. “Reversibility of Strontium Sorption on Fracture Fillings.” In 

Scientific Basis for Nuclear Waste Management XVIII, T. Murakami and R. C. Ewing (eds.), 

pp. 1045-1052. Material Research Society Symposium Proceedings, Volume 353, Materials 

Research Society, Pittsburgh, Pennsylvania. 

Del Debbio, J. A. 1991. “Sorption of Strontium, Selenium, Cadmium, and Mercury in Soil.” 


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 90 Sr in a Sand Aquifer at the 

Chalk River Nuclear Laboratories.” Journal of Contaminant Hydrology, 4:27-50. 

Keren, R., and G. A. O’Connor. 1983. “Strontium Adsorption by Noncalcareous Soils - 

Exchangeable Ions and Solution Composition Effects.” Soil Science, 135:308-315. 

H.24

Konishi, M., K. Yamamoto, T. Yanagi, and Y. Okajima. 1988. “Sorption Behavior of Cesium, 

Strontium and Americium Ions on Clay Materials.” Journal of Nuclear Science and 

Technology, 25:929-933. 

Lefevre, R., M. Sardin, and D. Schweich. 1993. “Migration of Strontium in Clayey and 

Calcareous Sandy Soil: Precipitation and Ion Exchange.” Journal of Contaminant 

Hydrology, 13:215-229. 

Lieser, K. H., B. Gleitsmann, and Th. Steinkopff. 1986. “Sorption of Trace Elements or 

Radionuclides in Natural Systems Containing Groundwater and Sediments.” Radiochimica 

Acta, 40:33-37. 

Lieser, K. H., and Th. Steinkopff. 1989. “Sorption Equilibria of Radionuclides or Trace 

Elements in Multicomponent Systems.” Radiochimica Acta, 47:55-61. 

McHenry, J. R. 1958. “Ion Exchange Properties of Strontium in a Calcareous Soil.” Soil 

Science Society of America, Proceedings, 22:514-518. 

Nelson, J. L. 1959. Recent Studies at Hanford on Soil and Mineral Reactions in Waste 

Disposal. HW-SA-2273, Westinghouse Hanford Company, Richland, Washington. 


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 85 Sr and 137 Cs 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



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. 




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, 


Sposito, G. 1984. The Surface Chemistry of Soils. Oxford University Press, New York, 

New York. 


Environmental Pollutant Assessment System. PNL-7145, Pacific Northwest National 


Vine, E. N., R. D. Aguilar, B. P. Bayhurst, W. R. Daniels, S. J. DeVilliers, B. R. Erdal, F. O. 

Lawrence, S. Maestas, P. Q. Oliver, J. L. Thompson, and K. Wolfsberg. 1980. Sorption- 

Desorption Studies on Tuff. II. A Continuation of Studies with Samples form Jackass Flats, 

Nevada and Initial Studies with Samples form Yucca Mountain, Nevada. LA-8110-MS, Los 

Alamos Scientific Laboratory, Los Alamos, New Mexico. 

H.26

APPENDIX I 

Partition Coefficients For Thorium

I.1.0 BACKGROUND 

Appendix I 

Partition Coefficients For Thorium 

Two generalized, simplifying assumptions were established for the selection of thorium K d 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: 

C Thorium adsorption occurs at concentrations less than 10 -9 M. The extent of thorium 

adsorption can be estimated by soil pH. 

C 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 K d value will be used in the look-up table to 

approximate thorium behavior in systems with high thorium concentrations. 


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 K d 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 K d 

values (not percent adsorbed or Freundlich or Langmuir constants) and were conducted in 

systems consisting of: 



C Dissolved thorium concentrations less than 10 -9 M 


descriptive statistics of the thorium K d data set are presented in Table I.1. The lowest thorium K d 

value was 100 ml/g for a measurement made on a pH 10 soil (Rancon, 1973). The largest 

thorium K d value was 500,000 ml/g for a measurement made on a silt/quartz soil of schist origin 

(Rancon, 1973). The average thorium K d value for the 17 observations was 54,000 ± 29,944 

ml/g. 

Table I.1. Descriptive statistics of thorium K d value data set presented in Section I.3. 

Thorium K d 

(ml/g) 

Clay 

Content 

(wt.%) 

pH CEC 

(meq/100 g) 

I.3 

Calcite 

(wt.%) 

Al/Fe- 

Oxides 

(wt.%) 

Mean 54,000 26.8 6.1 13.7 29 -- -- 

Standard Error 29,944 6.3 0.4 11.2 13.4 -- -- 

Median 5,000 30 6 2.9 25 -- -- 

Mode 100,000 40 6 2.9 0 -- -- 

Standard Deviation 123,465 14.1 1.5 29.8 30.1 -- -- 

Sample Variance 1.5x10 10 199.2 2.1 886.2 905 -- -- 

Minimum 100 12 4 1.7 0 -- -- 

Maximum 500,000 40 10 81.2 60 -- -- 

No. Observations 17 5 17 7 5 0 0 

I.2.0 Approach and Regression Models 

I.2.1 Correlations with Thorium K d Values 

Organic 

Matter 

(wt.%) 

A matrix of the correlation coefficients for thorium K d values with soil parameters is 

presented in Table I.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 K d 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 K d 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.

Table I.2. Correlation coefficients (r) of the thorium K d value data set presented in 

Section I.3. 

Thorium K d 

Thorium K d Clay Content pH CEC 

Clay Content -0.79 1 

pH 0.58 2 

1 

(0.14) 3 

-0.84 1 1 

CEC -0.15 -- -0.21 1 

Calcite 0.76 -0.998 2 0.85 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 I.1). Thus, it is possible for thorium K d/clay correlation coefficient of -0.79 to be not 

significant and the thorium K d /pH correlation coefficient of 0.58 to be significant because the former has 4 

degrees of freedom and the latter has 15 degrees of freedom. 

3 Excluding the Kd values at the highest pH value (pH 10), the correlation is 0.58 (n = 16). Including this K d 

value, the correlation coefficient decreases to 0.14. 

I.2.2 Thorium K d Values as a Function of pH 

Thorium K d 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 I.1 and Table I.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 I.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 K d values. Figure I.1 shows an increase in thorium 

K d 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. 

I.4

Figure I.1. Linear regression between thorium K d values 

and pH for the pH range from 4 to 8. [The 

single K d value at pH 10 is identified by the 

filled circle.] 

Table I.3. Calculated aqueous speciation of thorium as a function of pH. [The 

composition of the water and details of the aqueous speciation calculations are 

presented in Chapter 5. Total thorium concentration used in the aqueous 

speciation calculations is 1 ng/ml.] 

pH Dominant 


2+ 

3 ThF2 ThF + 

3 

2- 

7 Th(HPO4) 3 

I.5 

Percent (%) of 

Total Dissolved Thorium 

" 

9 Th(OH) 4 (aq) 99 

54 

42 

98

The regression equation between the pH range of 4 to 8 that is shown in Figure I.1 is 

log (Th K d) = -0.13 + 0.69(pH). (I.1) 

The statistics for this equation are presented in Table I.4. The fact that the P-value for the 

intercept coefficient is $0.05 indicates that the intercept is not significantly (P $ 0.05) different 

than 0. The fact that the P-value for the slope coefficient is #0.05 indicates that the slope is 

significantly (P $ 0.05) different than 1. The lower and upper 95 percent coefficients presented in 

Table I.4 reflect the 95 percent confidence limits of the coefficients. They were used to calculate 

the upper and lower limits of expected thorium K d values at a given pH value. 

I.2.3 Approach 





correlation between clay content and thorium K d 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 K d 

values and the solid phase parameters. Thus, the K d 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 I.4. Regression coefficient and their statistics relating thorium K d values and pH. 

[log (Th K d) = -0.13 + 0.69(pH), based on data presented in Figure I.1.] 

Coefficients Standard 

Error 

I.6 

t- Statistic P-value Lower 

95% 

Upper 

95% 

Intercept Coefficient 2.22 1.06 0.47 0.64 -1.77 2.76 

Slope Coefficient 0.57 0.18 3.24 0.006 0.19 0.95

The look-up table (Table I.5) for thorium K d values was based on thorium concentrations and pH. 

These 2 parameters have an interrelated effect on thorium K d values. The maximum 

concentration of dissolved thorium may be controlled by the solubility of hydrous thorium oxides 

(Felmy et al., 1991; Rai et al., 1995; Ryan 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 K d value was used in 

Table I.5 to describe high thorium concentrations. 

The following steps were taken to assign values to each category in the look-up table. For K d 

values in systems with pH values less than 8 and thorium concentrations less than the estimated 

solubility limits, Equation I.1 was used. This regression equation is for data collected between the 

pH range of 4 to 8 as shown in Figure I.1 [log (Th K d) = -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 K d 

values in the “pH 8 to 10” category were based on the single laboratory experiment conducted at 

pH 10 that had a K d of 200 ml/g. Upper and lower estimates of thorium K d values were 

calculated by adding or subtracting 1 logarithmic unit to the “central estimates” calculated above 

for each pH category (Figure I.2). The 1 logarithm unit estimates for the upper and lower limits 

are based on visual examination of the data in Figure I.1. The use of the upper and lower 

regression coefficient values at the 95 percent confidence limits (Table I.5) resulted in calculated 

ranges that were unrealistically large. At pH 4, for the “pH 3 to 5” category, the lower and upper 

log (Th K d) values were calculated to be 1 and 6.6, respectively; at pH 6.5, this range of K d was - 

0.5 to 9.0). All thorium K d values for systems containing concentrations of dissolved thorium 

greater than their estimated solubility limit (10 -9 M for pH 5 to 10 and 10 -2.6 M for pH < 5) were 

assigned a K d of 300,000 ml/g. 

Table I.5. Look-up table for thorium K d values (ml/g) based on pH and dissolved thorium 

concentrations. [Tabulated values pertain to systems consisting of low ionic 

strength (

I.3.0 K d Data Set for Soils 

Figure I.2. Linear regression between thorium K d values 

and pH for the pH Range 4 to 8. [Values ±1 

logarithmic unit from the regression line are 

also identified. The single K d value at pH 10 

is identified by the filled circle)]. 

The data set of thorium K d values used to develop the look-up table are listed in Table I.6. 

I.8

Thorium 

K d 

(ml/g) 

pH Clay 

(wt.%) 

CEC 1 

(meq/ 

100g) 

Table I.6. Data set containing thorium K d values. 

OM 1 

(wt.%) 

Fe- 

Oxides 

(wt.%) 

Th 

(M) 

Calcite 

(wt.%) 

I.9 

Solution 

Chemistry 

10,0000 7.6 3 Synthetic GW 1 , 

pH 6.6 

500,000 6 40 0 Syn. GW, 232 Th 

Competing Ion 

1,000 4 40 0 Syn. GW, 232 Th 

Competing Ion 

100,000 8 12 60 Syn. GW, 232 Th 

Competing Ion 

150,000 7 30 25 Syn. GW, 232 Th 

Competing Ion 

100 10 12 60 Syn. GW, 232 Th 

Competing Ion 

Soil ID and 

Characteristics 

Ref 2 

Soil A 1 

Silt+Qtz Sed., Schist soil 2 

Silt+Qtz Sed., Schist soil 2 

Silt+Qtz+OM+calcite, 

Schist Soil 

Cadarache Sed. 2 

Silt+Qtz+OM+calcite, 

Schist Soil 

24,000 6 Groundwater Glacial till, Clay 3 

5,800 6 Groundwater Fine Coarse Sand 3 

1,028.6 5.1 2.9 Gleyed Dystric Brunisol, Ae 

Horizon 4-15 cm 

1,271 5.2 2.1 Gleyed Dystric Brunisol, Bf 

Horizon1 5-45 cm 

5,000 4.5 Jefferson City, Wyoming, 

Fine Sandstone and Silty 

Clay 

10,000 5.8 Jefferson City, Wyoming, 


Clay 

15,000 7 Jefferson City, Wyoming, 


Clay 

1,578 5.2 81.2 Groundwater Gleyed Dystric Brunisol, Ah 

Horizon 

1,862.5 5.1 2.9 Groundwater Gleyed Dystric Brunisol, Ae 

Horizon 

1,153.7 5.2 2.1 Groundwater Gleyed Dystric Brunisol, Bf 

Horizon 

206.9 6.2 1.7 Groundwater Gleyed Dystric Brunisol, C 

Horizon 

1 CEC = cation exchange capacity, OC = organic matter, GW = groundwater. 

2 References: 1 =Legoux et al., 1992; 2 =Rancon, 1973; 3 = Bell and Bates, 1988; 4= Sheppard et al., 1987; 5 = Haji-Djafari et al., 

1981; 6 = Thibault et al., 1990. 

2 

2 

4 

4 

5 

5 

5 

6 

6 

6 

6

I.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, 


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 


Felmy, A. R., D. Rai, and D. A. Moore. 1993. “The Solubility of Hydrous Thorium(IV) Oxide in 

Chloride Media: Development of an Aqueous Ion-Interaction Model.” Radiochimica Acta, 

55:177-185. 

Haji-Djafari, S., P. E. Antommaria, and H. L. Crouse. 1981. Attenuation of Radionuclides and 

Toxic Elements by In Situ Soils at a Uranium Tailings Pond in Central Wyoming. In 

Permeability and Groundwater Contaminant Transport, (eds.) T. F. Zimmie and C. O. Riggs, 

pp. 221-242. American Society for Testing and Materials, Philadelphia, Pennsylvania. 


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. “K d 

Measurements of Activation, Fission, and Heavy Elements in Water/Solid Phase Systems.” 


I.10

Rai, D., A. R. Felmy, D. A. Moore, and M. J. Mason. 1995. “The Solubility of Th(IV) and 

U(IV) Hydrous Oxides in Concentrated NaHCO 3 and Na 2CO 3 Solutions.” In Scientific Basis 

for Nuclear Waste Management XVIII, Part 2, T. Murakami and R. C. Ewing (eds.), 

pp. 1143-1150, Materials Research Society Symposium Proceedings, Volume 353, Materials 


Rancon, D. 1973. “The Behavior in Underground Environments of Uranium and Thorium 

Discharged by the Nuclear Industry.” In Environmental Behavior of Radionuclides Released 

in the Nuclear Industry, pp. 333-346. IAEA-SM-172/55, International Atomic Energy 

Agency Proceedings, Vienna, Austria. 

Ryan, J. L., and D. Rai. 1987. “Thorium(IV) Hydrous Oxide Solubility.” Inorganic Chemistry, 

26:4140-4142. 





AECL-10125, Whiteshell Nuclear Research Establishment, Atomic Energy of Canada 

Limited, Pinawa, Canada. 

I.11

APPENDIX J 

Partition Coefficients For Uranium

J.1.0 Background 

Appendix J 

Partition Coefficients For Uranium 

The review of uranium K d values obtained for a number of soils, crushed rock material, and 

single-mineral phases (Table J.5) indicated that pH and dissolved carbonate concentrations are the 

2 most important factors influencing the adsorption behavior of U(VI). These factors and their 

effects on uranium adsorption on soils are discussed below. The solution pH was also used as the 

basis for generating a look-up table of the range of estimated minimum and maximum K d values 

for uranium. 

Several of the studies identified in this review demonstrate the importance dissolved carbonate 

through the formation of strong anionic carbonato complexes on the adsorption and solubility of 

dissolved U(VI). This complexation especially affects the adsorption behavior of U(VI) at 

alkaline pH conditions. Given the complexity of these reaction processes, it is recommended that 

the reader consider the application of geochemical reaction codes, and surface complexation 

models in particular, as the best approach to predicting the role of dissolved carbonate in the 

adsorption behavior of uranium and derivation of K d values when site-specific K d values are not 

available for U(VI). 

J.2.0 Availability of K d Values for Uranium 

More than 20 references were identified that reported the results of K d 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 K d 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 K d 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 K d 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 K d values, those in this report and published elsewhere, 

reported K d 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 K d 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 K d) or up to more 

than 95 percent (very large K d) of its original dissolved concentration. The former 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 K d 

values published in different sources, the reader must remember this source of uncertainty can be 

the major cause for the variability. 

In the following summaries, readers should note that the valence state of uranium is given as that 

listed in the authors’ publications. Typically, the authors describe their procedures and results in 

terms of “uranium,” and do not distinguish between the different valence states of uranium [U(VI) 

and U(IV)] present. In most studies, it is fair for the reader to assume that the authors are 

referring to U(VI) because no special precautions are described for conducting the adsorption 

studies using a dissolved reductant and/or controlled environmental chamber under ultralow 

oxygen concentrations. However, some measurements of uranium sorption onto crushed rock 

materials may have been compromised unbeknownst to the investigators by reduction of U(VI) 

initially present to U(IV) by reaction with ferrous iron [Fe(II)] exposed on fresh mineral surfaces. 

Because a major decrease of dissolved uranium typically results from this reduction due to 

precipitation of U(IV) hydrous-oxide solids (i.e., lower solubility), the measured K d values can be 

too large as a measure of U(VI) sorption. This scenario is possible when one considers the 

geochemical processes associated with some in situ remediation technologies currently under 

development. For example, Fruchter 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(III) present in the sediment minerals to Fe(II). Laboratory experiments have shown 

that dissolved U(VI) will accumulate, via reduction of U(VI) to U(IV) and subsequent 

precipitation as a U(IV) solid, when it contacts such treated sediments. 

J.3

J.2.2 Uranium K d Studies on Soils and Rock Materials 

The following sources of K d values considered in developing the uranium K d look-up table are 

listed in alphabetical order. Due to their extensive length, summary tables that list the uranium K d 

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 

(1.0x10 -4 , 1.0x10 -5 , 1.0x10 -6 , and 1.0x10 -7 M uranium) were used for the measurements. The pH 

values in the final solutions ranged from 7.65 to 8.48. Uranium K d values listed as “D” values in 

Ames et al. (1982, Table III) for the 23 " C sorption measurements are listed in Table J.5. 

Bell and Bates (1988) completed laboratory uranium (and other radionuclides) K d 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 K d 

measurements included a sample of borehole groundwater that was mixed in a solution-to-solid 

ratio of 10 ml/1 g with the

Erikson et al. (1993) determined the K d 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 K d 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 µg/l uranium. The water/soil mixtures were 

sampled at 7 and 30 days. The K d results are given in Table J.5. The K d 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 K d values measured for the Yuma Soil samples 

were attributed to carbonate complexation of the dissolved uranium. 

Giblin (1980) determined the K d 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-SO 4) containing 100 µg/l uranium. Ten milliliters of solution was 

mixed with 0.01 g of kaolinite for a solution-to-solid ratio of 1,000 ml/1 g. The pH of the 

suspension was adjusted to cover a range from 3.8 to 10. Uranium K d 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 K d values 

increase to a maximum of approximately 37,000 ml/g. At pH values from 7 to 10, the uranium 

adsorption decreased. 

Kaplan et al. (1998) investigated the effects of U(VI) concentration, pH, and ionic strength on the 

adsorption of U(VI) to a natural sediment containing carbonate minerals. The sediments used for 

the adsorption measurements were samples of a silty loam and a very coarse sand taken, 

respectively, from Trenches AE-3 and 94 at DOE’s Hanford Site in Richland, Washington. 

Groundwater collected from an uncontaminated part of the Hanford Site was equilibrated with 

each sediment in a ratio of 2 ml/1 g for 14 or 30 days. The Kd values listed in Kaplan et al. 

(1998) are given in Table J.5. The adsorption of U(VI) was determined to be constant for 

2+ 

concentrations between 3.3 and 100 µg/l UO2 at pH 8.3 and an ionic strength of 0.02 M. This 

result indicates that a linear Kd model could be used to describe the adsorption of U(VI) at these 

conditions. In those experiments where the pH was greater than 10, precipitation of 

U(VI)-containing solids occurred, which resulted in apparent Kd values greater than 400 ml/g. 

Kaplan et al. (1996) measured the K d values for U(VI) and several other radionuclides at 

geochemical conditions being considered in a performance assessment for the long-term disposal 

of radioactive low-level waste in the unsaturated zone at DOE’s Hanford Site in Richland, 

1 The uranium Kd values listed in Table J.5 for Giblin (1980) were provided by E. A. Jenne 

(PNNL, retired) based on work completed for another research project. The K d values were 

generated from digitization of the K d values plotted in Giblin (1980, Figure 1). 

J.5

Washington. The studies included an evaluation of the effects of pH, ionic strength, moisture 

content, and radionuclide concentration on radionuclide adsorption behavior. Methods used for 

the adsorption measurements included saturated batch adsorption experiments, unsaturated batch 

adsorption experiments, and unsaturated column adsorption experiments based on the 

Unsaturated Flow Apparatus (UFA). The measurements were conducted using uncontaminated 

pH 8.46 groundwater and the

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 K d 

values for uranium on this sediment material decrease as the sediment becomes less saturated. A 

K d 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 K d 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 K d 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 K d 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 K d values were determined for initial 

concentrations of 1.0x10 -4 , 1.0x10 -5 , 1.0x10 -6 , 1.0x10 -7 , and 2.15x10 -8 M uranium. The K d 

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 K d values for 238 U and 235 U for 48 soil/porewater 

samples. The K d values were determined from analyses of 238 U and 235 U in soil samples and 

associated porewaters taken from contaminated zones downgradient of the seepage basins. It 

should be noted that the mass concentration of 235 U is significantly less than (e.g.,

and then decreases with increasing pH over the pH range from 5.2 to 6.7. Serkiz and Johnson 

found that the field-derived K d values for 238 U and 235 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 K d values and geochemical modeling results, Serkiz and Johnson proposed that the 

uranium was not binding to the clays by a cation exchange reaction, but rather to a mineral 

surface coating with the variable surface charge varying due to the porewater pH. 


100,000 

10,000 

1,000 

100 

10 

1 

2 3 4 5 6 7 8 

Figure J.1. Field-derived K d values for 238 U and 235 U from Serkiz and 

Johnson (1994) plotted as a function of porewater pH for 

contaminated soil/porewater samples. [Square and circle 

symbols represent field-derived K d values for 238 U and 235 U, 

respectively. Solid symbols represent minimum K d values for 

238 U and 235 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.8 

pH


100,000 

10,000 

1,000 

100 

10 

1 

0 10 20 30 40 50 

Clay-Size Particle Content (wt%) 


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 K d values for 238 U and 235 U, respectively. Solid symbols 

represent minimum K d values for 238 U and 235 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


100,000 

10,000 

1,000 

100 

10 

1 

0 5 10 15 20 25 

CEC (meq/kg) 

Figure J.3. Field-derived K d values for 238 U and 235 U 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 K d values for 

238 U and 235 U, respectively. Solid symbols represent minimum Kd values for 

238 U and 235 U that were based on minimum detection limit values for the 

concentrations for the respective uranium isotopes in porewaters associated 

with the soil sample.] 

Serne et al. (1993) determined K d values for uranium and several other radionuclides at 

geochemical conditions associated with sediments at DOE’s Hanford Site in Richland, 

Washington. The K d values were measured using the batch technique with a well-characterized 

pH 8.3 groundwater and the

Sheppard and Thibault (1988) investigated the migration of several radionuclides, including 

uranium, through 3 peat 1 types associated with mires 2 typical of the Precambrian Shield in 

Canada. Cores of peat were taken from a floating sphagnum mire (samples designated PCE, peatcore 

experiment) and a reed-sedge mire overlying a clay deposit (samples designated SCE, sedgecore 

experiment). Uranium K d values were determined by in situ and batch laboratory methods. 

The in situ K d 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 K d values tabulated in Sheppard and 

Thibault (1988) are listed in Table J.5. Because the uranium K d values reported by Sheppard and 

Thibault (1988) represent uranium partitioning under reducing conditions, which are beyond the 

scope of our review, these K d values were not included in Figure J.4. Sheppard and Thibault 

(1988) noted that the uranium K d 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 K d values were 

similar to those determined in situ for the SCE peat sample. 

Thibault et al. (1990) present a compilation of soil K d 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 K d 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 K d values listed by Thibault et 

al. were collected from references that were not available during the course of our review. These 

sources included studies described in reports by M. I. Sheppard, a coauthor of Thibault et al. 

(1990), and papers by Dahlman et al. (1976), Haji-Djafari et al. (1981), Neiheisel (1983), 

Rançon (1973) and Seeley and Kelmers (1984). The uranium K d values, as listed in Thibault et al. 

(1990), taken for these sources are included in Table J.5. 

Warnecke and coworkers (Warnecke et al., 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 (K d values) that were conducted in support 

of Germany’s investigation of the Gorleben salt dome, Asse II salt mine, and former Konrad iron 

ore mine as disposal sites for radioactive waste. Experimental techniques included batch and 

recirculation methods as well as flow-through and diffusion experiments. The experiments were 

designed to assess the effects of parameters, such as temperature, pH, Eh, radionuclide 

concentration, complexing agents, humic substances, and liquid volume-to-soil mass ratio, on 

radionuclide migration and adsorption/desorption. These papers are overviews of the work 

completed in their program to date, and provide very few details on the experimental designs and 

individual results. There are no pH values assigned to the K d values listed in these overview 

1 Peat is defined as “an unconsolidated deposit of semicarbonized plant remains in a water 

saturated environment” (Bates and Jackson, 1980). 

2 A mire is defined as “a small piece of marshy, swampy, or boggy ground” (Bates and 

Jackson, 1980). 

J.11

papers. Warnecke et al. (1984) indicated that the measured pH values for the locations of soil 

and groundwater samples at Gorleben site studies range from 6 to 9. 

Warnecke et al. (1994) summarize experiments conducted during the previous 10 years to 

characterize the potential for radionuclide migration at site-specific conditions at the Gorleben 

site. Characteristic, minimum, and maximum K d 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 K d values. Warnecke et al. 

noted that the following progression in uranium K d values as function of sediment type was 

indicated: 

K d (Clay) > K d (Marl 1 ) > K d (Sandy) . 

Warnecke and Hild (1988) present an overview of the radionuclide migration experiments and 

adsorption/desorption measurements that were conducted for the site investigations of the 

Gorleben salt dome, Asse II salt mine, and Konrad iron ore mine. The uranium K d values listed in 

Warnecke and Hild are identical to those presented in Warnecke et al. (1994). The uranium K d 

values (ml/g) listed by Warnecke and Hild (1988, Table II) for sediments and different water types 

for the Konrad site are: 4 (Quaternary fresh water), 6 (Turonian fresh water), 6 (Cenomanian 

saline water), 20 [Albian (Hauterivain) saline water], 1.4 [Albian (Hils) saline water], 2.6 

(Kimmeridgian saline water), 3 (Oxfordian saline water), and 3 [Bajocian (Dogger) saline water]. 

Warnecke and Hild (1988, Table III) list minimum and maximum uranium K d values (0.54-15.2 

ml/g) for 26 rock samples from the Asse II site. No pH values were assigned to any of the 

tabulated K d 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.1. Uranium K d values (ml/g) listed by Warnecke et al. (1994, Table 1). 

Sediment 

Type 

Typical 

K d Value 

Fresh Water Saline Water 

Minimum 

K d Value 

Maximum 

K d Value 

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 

Typical 

K d Value 

Minimum 

K d Value 

Maximum 

K d Value 

Sandy 27 0.8 332 1 0.3 1.6 

Clayish 17 8.6 100 14 - 1,400 14.1 1,400

Warnecke et al. (1986) present an overview of the radionuclide migration experiments and 

adsorption/desorption measurements that were conducted for the Gorleben salt dome, and 

Konrad iron ore mine. The tabulated K d values for the Gorleben and Konrad site sediments and 

waters duplicate those presented Warnecke et al. (1994) and Warnecke and Hild (1988). 

Warnecke et al. (1984) present a short summary of radionuclide sorption measurements that were 

conducted by several laboratories in support of the Gorleben site investigation. Sediment 

(especially sand and silt) and water samples were taken from 20 locations that were considered 

representative of the potential migration path for radionuclides that might be released from a 

disposal facility sited at Gorleben. The minimum and maximum K d values listed by Warnecke et 

al. (1984, Table III) are 0.5 and 3,000 ml/g, respectively (note that these values are not listed as a 

function of pH). 

Zachara et al. (1992) studied the adsorption of U(VI) on clay-mineral separates from subsurface 

soils from 3 DOE sites. The materials included the clay separates (

Anderson et al. (1982) summarize an extensive study of radionuclides on igneous rocks and 

related single mineral phases. They report K d values for U(VI) sorption on apatite, attapulgite 

(also known as palygorskite), biotite, montmorillonite, and quartz. The K d values were 

determined using a batch technique using 10 -7 -10 -9 mol/l 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 K d values from Anderson 

et al. (1982, Figure 6a) are given in Table J.5. 1 

Ames et al. (1983a,b) investigated the effects of uranium concentrations, temperature, and 

solution compositions on the sorption of uranium on several well-characterized secondary and 

sheet silicate minerals. The secondary phases studied by Ames et al. (1983a, oxide analyses listed 

in their Table 3) included clinoptilotite, glauconite, illite, kaolinite, montmorillonite, nontronite, 

opal, and silica gel. The sheet silicate minerals used by Ames et al. (1983b, oxide analyses listed 

in their Table 1) consisted of biotite, muscovite, and phlogopite. The sorption of uranium on each 

mineral phase was measured with 2 solutions (0.01 M NaCl and 0.01 M NaHCO 3) using 4 initial 

uranium concentrations. The initial uranium concentrations used for the 25 " C experiments 

included 1.0x10 -4 , 1.0x10 -5 , 1.4x10 -6 , and 4.4x10 -7 mol/l uranium. The batch experiments were 

conducted under oxidizing conditions at 5, 25, and 65 " C in an environmental chamber. Solutions 

were equilibrated with the mineral solids in a ratio of 10 ml/1 g. A minimum of 30 days was 

required for the mineral/solution mixtures to reach steady state conditions. Uranium K d values 

calculated from the 25 " C sorption results given in Ames et al. (1983a, Table 6) are listed in Table 

J.5. 

Ames et al. (1983c) studied the effects of uranium concentrations, temperature, and solution 

compositions on the sorption of uranium on amorphous ferric oxyhydroxide. The sorption of 

uranium on amorphous ferric oxyhydroxide was measured with 2 solutions (0.01 M NaCl and 

0.01 M NaHCO 3) using 4 initial uranium concentrations. The initial uranium concentrations used 

for the 25 " C experiments included 1.01x10 -4 , 1.05x10 -5 , 1.05x10 -6 , and 4.89x10 -7 mol/l uranium 

for the 0.01 M NaCl solution, and 1.01x10 -4 , 1.05x10 -5 , 1.53x10 -6 , and 5.46x10 -7 mol/l uranium 

for the 0.01 M NaHCO 3 solution. The batch experiments were conducted under oxidizing 

conditions at 25 and 60 " C. The solutions were equilibrated for 7 days with the amorphous ferric 

oxyhydroxide in a ratio 3.58 l/g of iron in the solid. Uranium K d values calculated from the 25 " C 

sorption results given in Ames et al. (1983c, Table II) are listed in Table J.5. Reflecting the high 

adsorptive capacity of ferric oxyhydroxide, the K d values for the 25 " C measurements range from 

approximately 2x10 6 ml/g for the 0.01 M NaCl solution to approximately 3x10 4 ml/g for the 0.01 

M NaHCO 3 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 K d values 

were generated from digitization of the K d values plotted in Anderson et al. (1982, Figure 6a). 

J.14

Borovec (1981) investigated the adsorption of U(VI) and its hydrolytic complexes at 20 " C and 

pH 6.0 on fine-grained kaolinite, illite, and montmorillonite. The results indicate that the K d 

values increase with decreasing concentrations of dissolved uranium. At uranium concentrations 

less than 10 -4 mol/l, the uranium K d values for the individual minerals were constant. The K d 

values determined at 20 " C and pH 6.0 ranged from 50 to 1,000. The values increased in the 

sequence K d (kaolinite) < K d (illite) < K d (montmorillonite). Borovec presents the following linear 

equations for the maximum sorption capacity of uranium (a m, in meq/100 g) on clays at 20 " C and 

pH 6.0 with respect to CEC (in meq/100 g), 

and specific surface (A, in m 2 /g) of clays, 

a m = 0.90 CEC + 1.56 (r = 0.99522) , 

a m = 0.11 A + 2.05 (r = 0.97232) . 

J.2.4 Published Compilations Containing K d 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 K d, in their leaching 

model. A literature review was completed to evaluate appropriate distributions for K d values for 

various solutes, including uranium. Because Baes and Sharp (1983) are cited frequently as a 

source of K d values in other published K d reviews (e.g, Looney et al., 1987; Sheppard and 

Thibault, 1990), the uranium K d 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 K d values for 

cesium and strontium, they assumed a lognormal distribution for the K d values for all other 

elements in their compilation. Baes and Sharp listed an estimated default K d of 45 ml/g for 

uranium based on 24 uranium K d 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 K d values represent a diversity of soils, pure clays 

(other K d 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 d 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” K d of 39.8 (10 1.6 ) ml/g, and 

a range for its K d values of 0.1 to 1,000,000 ml/g. Looney et al. note that their recommended 

values are specific to the Savannah River Plant site, and they must be carefully reviewed and 

evaluated prior to using them in assessments at other sites. Nonetheless, such data are often used 

as “default values” in radionuclide migration assessment calculations, and are therefore listed here 

for the sake of completeness. It should be noted that the work of Looney et al. (1987) predates 

the uranium-migration and field-derived uranium K d 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 K d sorption databases used by different 

international organizations for performance assessments of repositories for radioactive wastes. 

The uranium K d values listed in McKinley and Scholtis (1993, Tables 1, 2, and 4) are listed in 

Table J.2. The reader should refer to sources cited in McKinley and Scholtis (1993) for details 

regarding their source, derivation, and measurement. Radionuclide K d 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 für Strahlen- und Umweltforschung 

m.b.H., Germany); IAEA (International Atomic Energy Agency, Austria); KBS (Swedish Nuclear 

Safety Board); NAGRA [Nationale Genossenschaft für die Lagerung radioaktiver Abfälle (Swiss 

National Cooperation for Storage of Radioactive Waste), Switzerland]; NIREX (United Kingdom 

Nirex Ltd.); NRC (U.S. Nuclear Regulatory Commission); NRPB (National Radiological 

Protection Board, United Kingdom); PAGIS [Performance Assessment of Geological Isolation 

Systems, Commission of the European Communities (CEC), Belgium; as well as PAGRIS SAFIR 

(Safety Assessment and Feasiblity Interim Report]; PSE (Projekt Sicherheitsstudien Entsorgung, 

Germany); RIVM [Rijksinstituut voor Volksgezondheid en Milieuhygience (National Institute of 

Public Health and Environment Protection), Netherlands]; SKI [Statens Kärnkraftinspektion 

(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 K d values listed by McKinley and Scholtis (1993, Tables 1, 2, and 4) 

from sorption databases used by different international organizations for 

performance assessments of repositories for radioactive wastes. 

Organization 

Argillaceous (Clay) Crystalline Rock Soil/Soil 

Sorbing 

Material 

K d 

(ml/g) 

Sorbing 

Material 

J.17 

K d 

(ml/g) 

Sorbing 

Material 

AECL Bentonite-Sand 100 Granite 5 Soil/Sediment 20 

GSF Sediment 2 

IAEA Pelagic Clay 500 

KBS-3 Bentonite 120 Granite 5,000 

NAGRA Bentonite 1,000 Granite 1,000 Soil/Sediment 20 

NIREX Clay Mudstone 10 

NRC 

Clay 5,000 Soil/Sediment 100 

Clay, Soil Shale 20 Granite 5 

Basalt 4 

Tuff 300 

NRPB Clay 300 Soil/Sediment 300 

PAGIS 

K d 

(ml/g) 

Bentonite 90 Soil/Sediment 1,700 

Subseabed 100 

PAGIS SAFIR Clay 600 

PSE Sediment 0.02 

RIVM Sandy Clay 10 

SKI Bentonite 200 Granite 5,000 

TVO 

UK DoE 

Bentonite 90 Crystalline 

Rock, Reducing 

Baltic Sea 

Sediment 

Ocean Sediment 500 

Lake Sediment 500 

500 Crystalline 

Rock, Real. 

200 Soil/Sediment 500 

Clay 200 Soil/Sediment 50 

Coastal Marine 

Water 

1000 

5

In a similar comparison of sorption databases for use in performance assessments of radioactive 

waste repositories, Stenhouse and Pöttinger (1994) list “realistic” K d values (ml/g) for uranium in 

crystalline rock/water systems of 1,000 (NAGRA), 5,000 [Svensk Kärnbränslehantering AB 

(Nuclear Fuel and Waste Management Company), Sweden; SKB], 1000 (TVO), and 6 (Canadian 

Nuclear Fuel Waste Management Programme, CNFWM). For bentonite/groundwater systems, 

they list 5,000 (NAGRA), 3,000 (SKB), and 500 (TVO). The reader should refer to sources cited 

in Stenhouse and Pöttinger 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 K d 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 K d values from other compilations, 

journal articles, and government laboratory reports for important elements, such as uranium, that 

would be present in the inventory associated with Canada’s nuclear fuel wastes. Because Thibault 

et al. (1990) and Sheppard and Thibault (1990) are frequently cited, their derived uranium K d 

values are reported here for the sake of completeness. The K d values for each element were 

categorized according to 4 soil texture types. These included sand (i.e., contains $70 percent 

sand-size particles), clay (i.e., contains $35 percent clay-size particles), loam (i.e., contains an 

even distribution of sand-, clay-, and silt-size particles, or #80 percent silt-size particles), and 

organic (i.e., contains >30 percent organic matter and are either classic peat or muck sediments, 

or the litter horizon of a mineral sediment). Based on their previous evaluations, Thibault et al. 

ln-transformed and averaged the compiled K d values to obtain a single geometric mean K d value 

for each element for each soil type. The K d values for each soil type and the associated range of 

K d values listed for uranium by Thibault et al. (1990) are given in Table J.3. 

Table J.3. Geometric mean uranium K d values derived by Thibault et al. (1990) for 

sand, loam, clay, and organic soil types. 

Soil Type 

Geometric 

Mean K d 

Values (ml/g) 

Observed Range of 

K d Values (ml/g) 

J.18 

Number of 

K d Values 

Sand 35 0.03 - 2,200 24 

Loam 15 0.2 - 4,500 8 

Clay 1,600 46 - 395,100 7 

Organic 410 33 - 7,350 6

J.3.0 Approach in Developing K d Look-Up Table 

The uranium K d values listed in Table J.5 are plotted in Figure J.4 as a function of pH. The K d 

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 K d values that are less than 10 ml/g. At pH 

values less than 3.5 and greater than 8, this variability includes extremely small K d values of less 

than 1 ml/g. 

Log Kd (ml/g) 

7 

6 

5 

4 

3 

2 

1 

0 

-1 

-2 

2 3 4 5 6 7 8 9 10 11 

Figure J.4. Uranium K d values used for development of K d look-up table. 

[Filled circles represent K d values listed in Table J.5. Open 

symbols (joined by dotted line) represent K d 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 K d 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 

pH

J.3.1 K d Values as a Function ff pH 

Although the uranium K d values in Figure J.4 exhibit a great deal of scatter at any fixed pH value, 

the K d 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 d values 

reported by Serkiz and Johnson (1994) (see Figure J.1), and percent adsorption values measured 

for uranium on single mineral phases as described above and those reported for iron oxides (Duff 

and Amrheim, 1996; Hsi and Langmuir, 1985; Tripathi, 1984; Waite et al., 1992, 1994; and 

others), clays (McKinley et al., 1995; Turner et al., 1996; Waite et al., 1992; and others), and 

quartz (Waite et 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 et al. (1992, 1994), and Duff and Amrheim (1996)], this 

pH-dependent behavior is related to the pH-dependent surface charge properties of the soil 

minerals and complex aqueous speciation of dissolved U(VI), especially near and above neutral 

pH conditions where dissolved U(VI) forms strong anionic uranyl-carbonato complexes with 

dissolved carbonate. 

J.3.2 K d Values as a Function of Mineralogy 

In addition to the sources of error and variability discussed above, the scatter in K d 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 K d values calculated 

from adsorption measurements (reported as percent uranium adsorbed versus pH) for ferrihydrite, 

kaolinite, and quartz by Waite et al. (1992). 

Uranium K d values were estimated 1 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, K d 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 K d 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 K d 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 K d values for ferrihydrite and quartz calculated from data in Waite et al. 

(1992) bound more than 95 percent of the uranium K d values gleaned from the literature. Above 

pH 7, the calculated uranium K d values for ferrihydrite and kaolinite effectively bound the 

maximum uranium K d values reported in the literature.. In terms of bounding the minimum K d 

values, the values calculated for quartz are greater than several data sets measured by Kaplan et 

al. (1996, 1998), Lindenmeirer et al. (1995), and Serne et al. (1993) for sediments from the 

Hanford Site in Richland, Washington which typically contain a significant quality of quartz and 

feldspar minerals. It should also be noted that some of the values listed from these studies 

represent measurements of uranium adsorption on Hanford sediments under partially saturated 

conditions. 

J.3.3 K d Values As A Function Of Dissolved Carbonate Concentrations 

As noted in several studies summarized above and in surface complexation studies of uranium 

adsorption by Tripathi (1984), Hsi and Langmuir (1985), Waite et al. (1992, 1994), McKinley et 

al. (1995), Duff and Amrheim (1996), Turner et al. (1996), and others, dissolved carbonate has a 

significant effect on the aqueous chemistry and solubility of dissolved U(VI) through the 

formation of strong anionic carbonato complexes. In turn, this complexation affects the 

adsorption behavior of U(VI) at alkaline pH conditions. Even differences in partial pressures of 

CO 2 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 CO 2 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 CO 2 or 

carbonate minerals . 

No attempt was made to statistically fit the K d 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 CO 2 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 K d values is through 

the use of surface complexation modeling techniques. 

J.3.4 K d Values as a Function of Clay Content and CEC 

No attempt was made to statistically fit the K d values summarized in Table J.5 as a function of 

CEC or concentrations of clay-size particles. The extent of clay concentration and CEC data, as 

noted from information included in Table J.5, is limited to a few studies that cover somewhat 

limited geochemical conditions. As discussed above, Serkiz and Johnson (1994) found no 

correlation between their uranium in situ K d 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 et al. (1994), McKinley et al. (1995), 

Turner et al. (1996), and others, uranium adsorption onto clay minerals is complicated and 

involves multiple binding sites, including exchange and edge-coordination sites. The reader is 

referred to these references for a detailed treatment of the uranium adsorption on smectite clays 

and application of surface complexation modeling techniques for such minerals. 

J.3.5 Uranium K d Look-Up Table 

Given the orders of magnitude variability observed for reported uranium K d values, a subjective 

approach was used to estimate the minimum and maximum K d values for uranium as a function of 

pH. These values are listed in Table J.4. For K d 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-K d values listed in 

Table J.4. 

Table J.4. Look-up table for estimated range of K d values for uranium based on pH. 

K d 

(ml/g) 

J.22 

pH 

3 4 5 6 7 8 9 10 

Minimum

The minimum and maximum K d 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 K d 

values collected in this review. The curve representing the minimum limit for uranium K d values 

is based on K d values calculated (described above) for quartz from data given in Waite et al. 

(1992) and the K d values reported by Kaplan et al. (1996, 1998), Lindenmeirer et al. (1995), and 

Serne et al. (1993). It is unlikely that actual K d values for U(VI) can be much lower than those 

represented by this lower curve. At the pH extremes along this curve, the uranium K d values are 

already very small. Moreover, if one considers potential sources of error resulting from 

experimental methods, it is difficult to rationalize uranium K d values much lower than this lower 

boundary. 

The curve representing the maximum limit for uranium K d values is based on K d 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 K d 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 K d values plotted in Figure J.4, and the 

assumption that some of the very large K d measurements may have included precipitation of 

uranium-containing solids due to starting uranium solutions being oversaturated. Moreover, as 

noted previously, measurements of uranium adsorption onto crushed rock samples may include 

U(VI)/U(IV) redox/precipitation reactions resulting from contact of dissolved U(VI) with Fe(II) 

exposed on the fresh mineral surfaces. 

J.4.0 Use of Surface Complexation Models to Predict Uranium K d 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 K d 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 d values for the geochemical conditions of interest, the reader is encouraged to 

apply this technology to predict bounding uranium K d values and their functionality with respect 

to important geochemical parameters. 

Numerous laboratory surface complexation studies for uranium have been reported in the 

literature. These include studies of uranium adsorption onto iron oxides (Duff and Amrheim, 

1996; Hsi and Langmuir, 1985; Tripathi, 1984; Waite et al., 1992, 1994; and others), clays 

(McKinley et al., 1995; Turner et al., 1996; Waite et al., 1992; and others), and quartz (Waite et 

J.23

al., 1992; and others). These references include derivation of the surface complexation constants 

for surface coordination sites determined to be important. 

In addition to these laboratory studies, there are numerous examples in the literature of the 

application of surface complexation models and published binding constants to predict and 

evaluate the migration of uranium in soil/groundwater systems. For example, Koß (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 app ) was derived for uranium sorption using the 

MINEQL geochemical code and site-specific geochemical data for soil CEC values, groundwater 

compositions, and measured uranium K d values. Quartz (SiO 2) 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 K d 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. 

Puigdomènech and Bergström (1994) evaluated the use of surface complexation models for 

calculating radionuclide sorption and K d 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 K d 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/l, 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 

["-FeO(OH)]. Their modeling results indicated that pH, inorganic carbon (i.e., dissolved 

carbonate), and Eh (redox conditions) are major parameters that affect uranium K d values. Under 

oxidizing conditions at pH values greater than 6, their derived K d values were approximately 100 

ml/g. At high concentrations of dissolved carbonate, and pH values greater than 6, the K d 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 K d 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 surfacecomplexation 

models (SCMs) [i.e., the diffuse layer model (DLM), constant capacitance model 

(CCM), and triple layer model (TLM)] in the geochemical reaction code MINTEQA2 to simulate 

potentiometric titration and adsorption data published for U(VI) and other radionuclides on 

several single mineral phases. Their studies were conducted in support of developing a uniform 

approach to using surface complexation models to predict radionuclide migration 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., K d) 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 K d 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. 

· Horráth (1960) measured an average enrichment factor of 200 to 350 for the adsorption 

of uranium on peat. 

· Kovalevskii (1967) determined that the uranium content of western Siberian noncultivated 

soils increased as a function of their clay content and that clay soils contained at least 3 

times more uranium than sands. 

1 The full citations listed for these references at the end of this appendix are provided exactly as 

given by Ames and Rai (1978). 

J.25

· Manskaya 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/l 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 K d 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 ( K 2O·6UO 3·11H 2O). Those experiments conducted with Ca-A 

zeolite contained a phase similar to becquerelite ( CaO·6UO 3·11H 2O). 

Ho and coworkers studied the adsorption of U(VI) on a well-characterized, synthetic hematite 

("-Fe 2O 3) 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 µm, surface area of 34 

m 2 /g, isoelectric point 2 of pH 7.6, and composition of >98 percent "-Fe 2O 3 and

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(VI) by Wyoming montmorillonite. The 

experiments consisted of reacting, at room temperature, the

pH 

Table J.5. Uranium K d values selected from literature for development of look-up table. 

U Kd 

(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 

(m 2 /g) Solution Soil Identification Reference / Comments 

8.3 1.98 Hanford Groundwater Trench 8 Loamy Sand Kaplan and Serne (1995, 

Part. Sat. Column, 40%) 



















8.3 2.0 5.2 Hanford Groundwater Trench 8 Loamy Sand Lindenmeir et al. (1995, 

Saturated Column 1) 






Unsat. Column 1, 65%) 


Unsat. UFA 1, 70%) 


Unsat. UFA 2, 24%) 


U nsat. Column 1, 63%) 


Unsat. Column 2, 43%) 


Unsat. UFA 1A, 29%) 


Unsat. UFA 1C, 29%) 

J.29

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 


8.4 0.20 5.3 6.3 Hanford Groundwater Trench 94 Kaplan et al. (1998, Batch) 





7.92 1.99 6.4 14.8 Hanford Groundwater Trench AE-3 Kaplan et al. (1998, Batch) 










8.46 0.90 6.4 14.8 Hanford Groundwater Trench AE-3 Kaplan et al. (1996, 100% 

Unsaturated Batch) 

8.46 1.70 5.3 6.3 Hanford Groundwater Trench 94 Kaplan et al. (1996, 100% 


8.46 1.00 6.0 6.3 Hanford Groundwater TSB-1 Kaplan et al. (1996, 100% 




8.46 2.10 6.0 6.3 Hanford Groundwater TSB-1 Kaplan et al. (1996, Batch) 

8.46 0.24 6.4 14.8 Hanford Groundwater Trench AE-3 Kaplan et al. (1996) 






8.46 0.23 5.3 6.3 Hanford Groundwater Trench 94 Kaplan et al. (1996) 





J.30

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 


2 8 Sand Neiheisel [1983, as listed 

in Thibault et al. (1990)] 





750 36 Clayey Sand Neiheisel [1983, as listed 






2.00 100 Fine Sandstone and 

Silty Sand 

4.50 200 Fine Sandstone and 

Silty Sand 

5.75 1,000 Fine Sandstone and 

Silty Sand 

7.00 2,000 Fine Sandstone and 

Silty Sand 

J.31 

Haji-Djafari et al. [1981, as 

listed in Thibault et al. 

(1990)] 



(1990)] 



(1990)] 



(1990)] 

5.6 25,000 Red-Brown Clayey Seeley and Kelmers [1984, as 


(1990)] 

5.6 250 Red-Brown Clayey Seeley and Kelmers [1984, as 


(1990)] 

5.20 58.4 Thibault et al. (1990, values 

determined by coworkers) 



5.20 160 Thibault et al. (1990, values 




7.00 450 36 28.0 Silty Loam Clay Thibault et al. (1990, values 

determined by coworkers)

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 


7.30 1.2 15 17.0 Loam Thibault et al. (1990, values 


4.90 0.03 2 5.8 Medium Sand Thibault et al. (1990, values 


5.50 2900 1 120.0 Organic Thibault et al. (1990, values 


7.40 1.9 10 9.1 Fine Sandy Loam Thibault et al. (1990, values 


7.40 2.4 11 8.7 Fine Sandy Loam Thibault et al. (1990, values 


6.60 590 10 10.8 Fine Sandy Loam Thibault et al. (1990, values 






7.00 16 Sand Rancon [1973, as listed in 

Thibault et al. (1990)] 

7.00 33 Organic Peat Rancon [1973, as listed in 

Thibault et al. (1990)] 

6.50 4400 Clay Fraction Dahlman et al. [1976, as listed 


2.80 200 Abyssal Red Clay Erickson (1980) 

7.10 790,000 Abyssal Red Clay Erickson (1980) 

8.3 1.70 2.6 Hanford Groundwater CGS-1 sand (coarse 

gravel sand) 

8.3 2.30 5.2 Hanford Groundwater Trench 8 Loamy Sand 

(medium/coarse sand) 

8.3 79.30 6.0 Hanford Groundwater TBS-1 Loamy Sand 

(Touchet Bed sand) 

J.32 

Serne et al. (1993, Batch) 



8.00 56.0 Hanford Groundwater, GR-1 Umtanum Basalt Salter et al. (1981) 





8.00 13.0 Hanford Groundwater, GR-1 Flow E Basalt Salter et al. (1981)

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 


8.00 2.7 Hanford Groundwater, GR-1 Flow E Basalt Salter et al. (1981) 




8.00 16.0 Hanford Groundwater,GR-1 Pomona Basalt Salter et al. (1981) 





10.00 2.8 Hanford Groundwater,GR-2 Umtanum Basalt Salter et al. (1981) 





10.00 1.0 Hanford Groundwater,GR-2 Flow E Basalt Salter et al. (1981) 










7.66 7.5 1.83 17.7 Hanford Groundwater,GR-1 Umtanum Basalt Ames et al. (1982) 

7.66 13 1.83 17.7 Hanford Groundwater,GR-1 Umtanum Basalt Ames et al. (1982) 







7.65 2.7 1.5 10.3 Hanford Groundwater,GR-1 Flow E Basalt Ames et al. (1982) 


J.33

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 








7.90 2.2 4.84 31.2 Hanford Groundwater,GR-1 Pomona Basalt Ames et al. (1982) 








7.7 27 71.66 646 Hanford Groundwater,GR-1 Smectite, secondary Ames et al. (1982) 

7.7 39 4.84 31.2 Hanford Groundwater,GR-1 Smectite, secondary Ames et al. (1982) 







6.85 477,285 0.01 NaCl Amor Fe(III) 

Hydroxide 

6.80 818,221 0.01 NaCl Amor Fe(III) 

Hydroxide 

6.90 1,739,87 

7 

6.90 1,690,52 

2 

0.01 NaCl Amor Fe(III) 

Hydroxide 

0.01 NaCl Amor Fe(III) 

Hydroxide 

8.60 4,313 0.01 NaHCO 3 Amor Fe(III) 

Hydroxide 


Hydroxide 


Hydroxide 


Hydroxide 

J.34 

Ames et al. (1983c) 







Ames et al. (1983c)

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 


7.15 8.4 15.3 1.59 0.01 NaCl Biotite Ames et al. (1983b) 




7.15 113.7 0.95 1.88 0.01 NaCl Muscovite Ames et al. (1983b) 




7.15 67.9 1.17 1.22 0.01 NaCl Phlogopite Ames et al. (1983b) 



8.65 0.9 15.3 1.59 0.01 NaHCO 3 Biotite Ames et al. (1983b) 




8.65 2.2 0.95 1.88 0.01 NaHCO 3 Muscovite Ames et al. (1983b) 




8.65 0.6 1.17 1.22 0.01 NaHCO 3 Phlogopite Ames et al. (1983b) 



7 544.5 25 116.1 0.01 NaCl Illite, only lowest U 

conc 

8.5 90.5 25 116.1 0.01 NaHCO 3 Illite, only lowest U 

conc 

7 657.8 12.2 68.3 0.01 NaCl Kaolinite, only lowest 

U conc 

8.5 400.8 12.2 68.3 0.01 NaHCO 3 Kaolinite, only lowest 

U conc 

7 542.0 120 747 0.01 NaCl Montmorillonite, only 

lowest U conc 

8.5 1.8 120 747 0.01 NaHCO 3 Montmorillonite, only 

lowest U conc 

7 299.9 95 861 0.01 NaCl Nontronite, only lowest 

U conc 

J.35 

Ames et al. (1983a) 






Ames et al. (1983a)

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 


8.5 4.1 95 861 0.01 NaHCO 3 Nontronite, only lowest 

U conc 

7 138.0 16.03 137.3 0.01 NaCl Glauconite, only lowest 

U conc 

8.5 114.2 16.03 137.3 0.01 NaHCO 3 Glauconite, only lowest 

U conc 

7 66.5 140.2 20 0.01 NaCl Clinoptilolite, only 

lowest U conc 

8.5 0.6 140.2 20 0.01 NaHCO 3 Clinoptilolite, only 

lowest U conc 

7 225.7 3.18 46.8 0.01 NaCl Opal, only lowest U 

conc 

8.5 1.7 3.18 46.8 0.01 NaHCO 3 Opal, only lowest U 

conc 

7 300.5 2.79 626.3 0.01 NaCl Silica Gel,, only lowest 

U conc 

8.5 639.9 2.79 626.3 0.01 NaHCO 3 Silica Gel,, only lowest 

U conc 

J.36 










7.3 4200.0 4.36 Spesutie (silt loam) Erikson et al. (1993) 

6.2 136.0 1.29 Transonic (silt loam) Erikson et al. (1993) 

8.0 44 9.30 Yuma (sandy loam) Erikson et al. (1993) 

6.8 4360 4.36 Spesutie (silt loam) Erikson et al. (1993) 

5.6 328 1.29 Transonic (silt loam) Erikson et al. (1993) 

8.0 54 9.30 Yuma (sandy loam) Erikson et al. (1993) 

39 River Sediment 

(Quartz, clay, calcite, 

organic matter) 

Rancon (1973) as cited 

by Ames and Rai (1978) 

33 River Peat Rancon (1973) as cited 


16 River Sediment 

(Quartz, clay, calcite) 

270 Soil (Quartz and Clay, 

from Altered Schist) 





0 Quartz Rancon (1973) as cited 


7 Calcite Rancon (1973) as cited 


139 Illite Rancon (1973) as cited 

by Ames and Rai (1978)

pH 


(ml/g) 

27 

(0.8- 

332) 

1 

(0.3-1.6) 

17 

(8.5- 

100) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 


Fresh Water Gorleben Salt Dome, 

Sandy Sediment 

Fresh Water Gorleben Salt Dome, 

Sandy Sediment 

Saline Water Gorleben Salt Dome, 

Clayish Sediment 

14-1,400 Saline Water Gorleben Salt Dome, 

Clayish Sediment 

4 Quaternary fresh water Former Konrad Iron 

Ore Mine 

6 Turonian fresh water Former Konrad Iron 

Ore Mine 

6 Cenomanian saline water Former Konrad Iron 

Ore Mine 

20 Albian (Hauterivain) saline 

water 

J.37 

Former Konrad Iron 

Ore Mine 

1.4 Albian (Hils) saline water Former Konrad Iron 

Ore Mine 

2.6 Kimmeridgian saline water Former Konrad Iron 

Ore Mine 

3 Oxfordian saline water Former Konrad Iron 

Ore Mine 

3 Bajocian (Dogger) saline 

water 

3.83 310 Synthetic Groundwater, 

function of pH 















Former Konrad Iron 

Ore Mine 

Warnecke et al. (1984, 1986, 

1994), Warnecke and Hild 

(1988) 



(1988) 



(1988) 



(1988) 

Warnecke et al. (1986), 

Warnecke and Hild (1988) 















Kaolinite Giblin (1980) 







Kaolinite Giblin (1980)

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 











4.81 1,216 Synthetic Groundwater, 




































Surface 

Area 


J.38 
























pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 















































Surface 

Area 


J.39 
























pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 











































Surface 

Area 


J.40 







Quartz Andersson et al. (1982) 













Biotite Andersson et al. (1982) 

Biotite Andersson et al. (1982)

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 











































Surface 

Area 


J.41 











Apatite Andersson et al. (1982) 










Attapulgite 

(Palygorskite) 

Andersson et al. (1982)

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 







































Surface 

Area 


J.42 

Attapulgite 


Attapulgite 


Attapulgite 


Attapulgite 


Attapulgite 


Attapulgite 


Attapulgite 


Attapulgite 


Attapulgite 


Attapulgite 


5.1 7,391 45 99 Ca Electrolyte, CO 2 Free Kenoma Clay,

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 


5.1 2,180 45 99 Ca Electrolyte, CO 2 Free Kenoma Clay,

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 


5.1 508 59 112 Ca Electrolyte, CO 2 Free Ringold Clay Isolate, 

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 


4.9 918 59 112 Ca Electrolyte, CO 2 Free Ringold Clay Isolate, 

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 


6.09 9,400 Reducing Conditions PCE Deep Core, 25-32 

cm 

6.09 12,500 Reducing Conditions PCE Deep Core, 33-40 

cm 

5.94 3,000 Reducing Conditions SCE Surface Core, 0-5 

cm 

6.82 8,800 Reducing Conditions SCE Surface Core, 

6-20 cm 


21-25 cm 


26-30 cm 

7.28 700 Reducing Conditions SCE Surface Core, 

31-40 cm 

1,300 Reducing Conditions PCE Surface Core, 

0-40 cm 

2,100 Reducing Conditions PCE Deep Core, 40-80 

cm 

2,000 Reducing Conditions SCE Surface Core, 

1-10 cm 

2,900 Reducing Conditions SCE Surface Core, 

10-30 cm 

870 Reducing Conditions SCE Surface Core, 

30-40 cm 

5.7 46 2.3 Site Borehole Groundwater Clay (Glacial Till, Less 

Than 5 mm) 

5.7 46 3.0 Site Borehole Groundwater C1:2 (Brown, Slightly 

Silty, Less Than 5 mm) 

5.7 900 2.7 Site Borehole Groundwater C3 (Dark Brown 

Coarse Granular 

Deposit, Less Than 5 

mm) 

5.7 2,200 2.9 Site Borehole Groundwater C6 (Brown Coarse 

Granular Deposit, Less 

Than 5 mm) 

5.7 560 0.8 Site Borehole Groundwater Sand (Light Brown 

Coarse Granular 

Deposit, Less Than 5 

mm) 

J.46 

Sheppard and Thibault 

(1988, In Situ) 


(1988, In Situ) 


(1988, In Situ) 


(1988, In Situ) 


(1988, In Situ) 


(1988, In Situ) 


(1988, In Situ) 


(1988, Batch) 


(1988, Batch) 


(1988, Batch) 


(1988, Batch) 


(1988, Batch) 

Bell and Bates (1988) 





4.16 85.0 0.5 1.11 A12 Serkiz and Johnson (1994) 


3.42 5.3 3 3.74 A13R Serkiz and Johnson (1994) 

3.19 2.1 1.5 1.39 A22 Serkiz and Johnson (1994)

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 




3.5 1.4 3.1 1 A32 Serkiz and Johnson (1994) 


5.42 2,200.0 2.5 0.68 A52 Serkiz and Johnson (1994) 

3.72 2.3 2 0.42 A53 Serkiz and Johnson (1994) 

3.24 2.7 2.8 4.71 B13 Serkiz and Johnson (1994) 


3.86 10.1 4.9 B23 Serkiz and Johnson (1994) 

4.02 5.2 2.5 3.8 B23R Serkiz and Johnson (1994) 







3.61 5.3 B52 Serkiz and Johnson (1994) 


3.68 6.4 C13 Serkiz and Johnson (1994) 

3.75 23.0 6.4 C14 Serkiz and Johnson (1994) 


4.17 980.0 6.4 6.12 C23 Serkiz and Johnson (1994) 

5.53 3,600.0 5.5 2.54 C32 Serkiz and Johnson (1994) 



4.51 13,000.0 3 5.04 C43 Serkiz and Johnson (1994) 

6.78 11,000.0 5.3 1.96 D13 Serkiz and Johnson (1994) 

4.14 13.0 D13RA Serkiz and Johnson (1994) 

9.3 2 2.55 D13RB Serkiz and Johnson (1994) 

4 320.0 10.5 11.4 E13 Serkiz and Johnson (1994) 

4.04 310.0 4.5 8.5 E14 Serkiz and Johnson (1994) 

5.85 2,700.0 6.4 15.5 E23 Serkiz and Johnson (1994) 

4.32 980.0 3.9 13.3 E23R Serkiz and Johnson (1994) 




J.47

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 




4.3 7,500.0 15.5 16.1 F12 Serkiz and Johnson (1994) 

4.9 830.0 8.51 F13 Serkiz and Johnson (1994) 

4.69 160.0 8.1 7.48 F22 Serkiz and Johnson (1994) 

6.48 16,000.0 13 11.6 F23 Serkiz and Johnson (1994) 









3.42 9.46 3 3.74 A13R Serkiz and Johnson (1994) 




3.5 1.38 3.1 1 A32 Serkiz and Johnson (1994) 



3.72 16.0 2 0.42 A53 Serkiz and Johnson (1994) 



3.86 10.7 4.9 B23 Serkiz and Johnson (1994) 








3.61 2.6 B52 Serkiz and Johnson (1994) 


3.68 5.6 C13 Serkiz and Johnson (1994) 


J.48

pH 


(ml/g) 

Clay 

Cont. 

(wt.%) 

CEC 

(meq/100g) 

Surface 

Area 







4.51 460.0 3 5.04 C43 Serkiz and Johnson (1994) 

6.78 690.0 5.3 1.96 D13 Serkiz and Johnson (1994) 

4.14 26.6 D13RA Serkiz and Johnson (1994) 

22.6 2 2.55 D13RB Serkiz and Johnson (1994) 

4 650.0 10.5 11.4 E13 Serkiz and Johnson (1994) 


4.32 310.0 3.9 13.3 E23R Serkiz and Johnson (1994) 







4.9 2,200.0 8.51 F13 Serkiz and Johnson (1994) 


6.48 950.0 13 11.6 F23 Serkiz and Johnson (1994) 







J.49

J.6.0 References 

Ames, L. L., J. E. McGarrah, B. A. Walker, and P. F. Salter. 1982. “Sorption of Uranium and 

Cesium by Hanford Basalts and Associated Secondary Smectite.” Chemical Geology, 

35:205-225. 

Ames, L. L., J. E. McGarrah, B. A. Walker, and P. F. Salter. 1983c. “Uranium and Radium 

Sorption on Amorphous Ferric Oxyhydroxide.” Chemical Geology, 40:135-148. 

Ames, L. L., J. E. McGarrah, and B. A. Walker. 1983a. “Sorption of Trace Constituents from 

Aqueous Solutions onto Secondary Minerals. I. Uranium.” Clays and Clay Minerals, 

31(5):321-334. 

Ames, L. L., J. E. McGarrah, and B. A. Walker. 1983b. “Sorption of Uranium and Radium by 

Biotite, Muscovite, and Phlogopite.” Clays and Clay Minerals, 31(5):343-351. 

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. Conclusions and Evaluation. EPA 520/6-78-007 (Volume 1 of 2), U.S. 

Environmental Protection Agency, Las Vegas, Nevada. 

Amonette, J. E., J. E. Szecsody, H. T. Schaef, J. C. Templeton, Y. A. Gorby, and J. S. Fruchter. 

1994. “Abiotic Reduction of Aquifer Materials by Dithionite: A Promising In-Situ 

Remediation Technology.” In In-Situ Remediation: Scientific Basis for Current and Future 

Technologies. Thirty-Third Hanford Symposium on Health and the Environment, November 

7-11, 1994, Pasco, Washington, G. W. Gee and N. R. Wing (eds.). Battelle Press, Richland, 

Washington. 

Andersson, K., B. Torstenfelt, and B. Allard. 1982. "Sorption Behavior of Long-Lived 

Radionuclides in Igneous Rock." In Environmental Migration of Long-Lived Radionuclides 

Proceedings of an International Symposium on Migration in the Terrestrial Environment of 

Long-Lived Radionuclides from the Nuclear Fuel Cycle Organized by the International 

Atomic Energy Agency, the Commission of the European Communities and the OECD 

Nuclear Energy Agency and held in Knoxville, United States, 27-31 July 1981., Knoxville, 

Tennessee. IAEA-SM-257/20. pp. 111-131. International Atomic Energy Agency, Vienna, 

Austria. 

Baes, C. F., III, and R. D. Sharp. 1983. “A Proposal for Estimation of Soil Leaching and 

Leaching Constants for Use in Assessment Models.” Journal of Environmental Quality, 

12:17-28. 

Bates, R. L., and J. A. Jackson (eds.). 1980. Glossary of Geology. American Geological 

Institute, Falls Church, Virginia. 

J.50

Barney, G. S. 1982a. Radionuclide Sorption on Basalt Interbed Materials FY 1981 Annual 

Report. RHO-BW-ST-35 P, Rockwell Hanford Operations, Richland, Washington. 

Barney, G. S. 1982b. Radionuclide Sorption of Columbia River Basalt Interbed Materials. 

RHO-BW-SA-198 P, Rockwell Hanford Operations, 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 


Borovec, Z. 1981. “The Adsorption of Uranyl Species by Fine Clay.” Chemical Geology, 

32:45-58. 

Borovec, Z., B. Kribek, and V. Tolar. 1979. “Sorption of Uranyl by Humic Acids.” Chemical 

Geology, 27:39-46. 

Brindley, G. W., and M. Bastovanov. 1982. “Interaction of Uranyl Ions with Synthetic Zeolites 

of Type A and the Formation of Compreignacite-Like and Becquerelite-Like Products.” 


Chisholm-Brause, C., S. D. Conradson, C. T. Buscher, P. G. Eller, and D. E. Morris. 1994. 

“Speciation of uranyl Sorbed at Multiple Binding Sites on Montmorillonite.” Geochimica et 

Cosmochimica Acta, 58(17):3625-3631. 

Dahlman, R. C., E. A. Bondietti, and L. D. Eyman. 1976. Biological Pathways and Chemical 

Behavior of Plutonium and Other Actinides in the Environment. In Actinides in the 

Environment, (ed.) A. M. Friedman, pp. 47-80. ACS Symposium Series 35, American 


Dement’yev, V. S., and N. G. Syromyatnikov. 1968. “Conditions of Formation of a Sorption 

Barrier to the Migration of Uranium in an Oxidizing Environment.” Geochemistry 

International, 5:394-400 

Doi, K., S. Hirono, and Y. Sakamaki. 1975. “Uranium Mineralization by Ground Water in 

Sedimentary Rocks, Japan.” Economic Geology, 70:628-646. 

Duff, M. C., and C. Amrhein. 1996. “Uranium(VI) Adsorption on Goethite and Soil in 

Carbonate Solutions.” Soil Science Society of America Journal, 60(5):1393-1400. 

Erickson, K. L. 1980. Radionuclide Sorption Studies on Abyssal Red Clays. In Scientific Basis 

for Nuclear Waste Management. Volume 2, (ed.) C. J. M. Northrup, Jr., pp. 641-646. 

Plenum Press, New York, New York. 

J.51

Erikson, R. L., C. J. Hostetler, R. J. Serne, J. R. Divine, and M. A. Parkhurst. 1993. 

Geochemical Factors Affecting Degradation and Environmental Fate of Deleted Uranium 

Penetrators in Soil and Water. PNL-8527, Pacific Northwest Laboratory, Richland, 

Washington. 

Fruchter, J. S., J. E. Amonette, C. R. Cole, Y. A. Gorby, M. D. Humphrey, J. D. Isok, F. A. 

Spane, J. E. Szecsody, S. S. Teel, V. R. Vermeul, M. D. Williams, and S. B. Yabusaki, 1996, 

In Situ Redox Manipulation Field Injection Test Report - Hanford 100-H Area. 

PNNL-11372, Pacific Northwest National Laboratory, Richland, Washington. 

Giblin, A. M. 1980. "The Role of Clay Adsorption in Genesis of Uranium Ores." Uranium in the 

Pine Creek Geosyncline. In Proceedings of the International Uranium Symposium on the 

Pine Creek Geosyncline Jointly Sponsored by the Bureau of Mineral Resources, Geology, 

and Geophysics and the CSIRO Institute of Earth Resources in Co-operation with the 

International Atomic Energy Agency and Held in Sydney, Australia 4-8 June, 1979, eds. J. 

Ferguson and A. B. Goleby, pp. 521-529. International Atomic Energy Agency, Vienna, 

Austria. 

Goldsztaub, S. and R. Wey. 1955. “Adsorption of Uranyl Ions by Clays.” Bull. Soc. Franc. 

Mineral. Crist., 78:242. 

Haji-Djafari, S., P. E. Antommaria, and H. L. Crouse. 1981. Attenuation of Radionuclides and 

Toxic Elements by In Situ Soils at a Uranium Tailings Pond in Central Wyoming. In 

Permeability and Groundwater Contaminant Transport, (eds.) T. F. Zimmie and C. O. Riggs, 

pp. 221-242. American Society for Testing and Materials, Philadelphia, Pennsylvania. 

Ho, C. H., and N. H. Miller. 1986. “Adsorption of Uranyl Species from Bicarbonate Solution 

onto Hematite Particles.” Journal of Colloid and Interface Science, 110:165-171. (Note 

paper issued under report number AECL-8433, Atomic Energy of Canada Limited, Whiteshell 

Nuclear Research Establishment, Pinawa, Manitoba, Canada.) 

Ho, C. H., and N. H. Miller. 1985. “Effect of Humic Acid on Uranium Uptake by Hematite 

Particles.” Journal of Colloid and Interface Science, 106:281-288. (Note paper issued under 

report number AECL-8432, Atomic Energy of Canada Limited, Whiteshell Nuclear Research 

Establishment, Pinawa, Manitoba, Canada.) 

Ho, C. H., and D. C. Doern. 1985. “The Sorption of Uranyl Species on a Hematite Sol.” 

Canadian Journal of Chemistry, 63:1100-1104. (Note paper issued under report number 

AECL-8038, Atomic Energy of Canada Limited, Whiteshell Nuclear Research Establishment, 

Pinawa, Manitoba, Canada.) 

J.52

Horráth, E. 1960. “Investigations of Uranium Adsorption to Peat in Natural Waters Containing 

U-Traces.” Magyar Tudomanyos Akad. Atommag Kutató Intézete, Közlemenyek, 2:177-183 

(in Hungarian). 

Hsi, C-K. D., and D. Langmuir. 1985. “Adsorption of Uranyl Onto Ferric Oxyhydroxides: 

Application of the Surface Complexation Site-Binding Model.” Geochimica et 


Johnson, W. H., S. M. Serkiz, L. M. Johnson, and S. B. Clark. 1994. Uranium Partitioning 

Under Acidic Conditions in a Sandy Soil Aquifer. WSRC-MS--94-0528, Westinghouse 

Savannah River Company, Savannah River Site, Aiken, South Carolina. 

Kaplan, R. J. Serne, A. T. Owen, J. Conca, T. W. Wietsma, and T. L. Gervais. 1996. 

Radionuclide Adsorption Distribution Coefficient Measured in Hanford Sediments for the 

Low Level Waste Performance Assessment Project. PNNL-11385, Pacific Northwest 

National Laboratory, Richland, Washington. 

Kaplan, D. I., T. L. Gervais, and K. M. Krupka. 1998. “Uranium(VI) Sorption to Sediments 

Under High pH and Ionic Strength Conditions.” Radiochimica Acta, 80:201-211. 

Kaplan, D. I., and R. J. Serne. 1995. Distribution Coefficient Values Describing Iodine, 

Neptunium, Selenium, Technetium, and Uranium Sorption to Hanford Sediments.” 

PNL-10379 (Supplement 1), Pacific Northwest Laboratory, Richland, Washington. 

Kent, D. B., V. S. Tripathi, N. B. Ball, J. O. Leckie, and M. D. Siegel. 1988. Surface- 

Complexation Modeling of Radionuclide Adsorption in Subsurface Environments. 

NUREG/CR-4807, U.S. Nuclear Regulatory Commission, Washington, D.C. 

Kohler, M., G. P. Curtis, D. B. Kent, and J. A. Davis. 1996. “Experimental Investigation and 

Modeling of Uranium(VI) Transport Under Variable Chemical Conditions.” Water Resources 

Research, 32(12):3539-3551. 

Koß, V. 1988. “Modeling of Uranium(VI) Sorption and Speciation in a Natural Sediment 

Groundwater System.” Radiochimica Acta, 44/45:403-406. 

Kovalevskii, A. L. 1967. “Dependence of the Content of Some Trace Elements on the Clayiness 

of Soils.” Mikroelem. Biosfere Ikh Primen. Scl. Khoz. Med. Sib. Dal’nego Vostoka, Dokl. 

Sib. Knof., 2nd. 1964. O. V. Makew. Buryat. Khizhn. lzd. Ulan-Ude, USSR. 

Krupka, K. M., D. Rai, R. W. Fulton, and R. G. Strickert. 1985. "Solubility Data for U(VI) 

Hydroxide and Np(IV) Hydrous Oxide: Application of MCC-3 Methodology," pp. 753-760. 

In Scientific Basis for Nuclear Waste Management VIII, eds. C. M. Jantzen, J. A. Stone, and 

J.53

R. C. Ewing. Materials Research Society Symposium Proceedings, Volume 44, Materials 


Lindenmeier, C. W., R. J. Serne, J. L. Conca, A. T. Owen, and M. I. Wood. 1995. Solid Waste 

Leach Characteristics and Contaminant-Sediment Interactions Volume 2: Contaminant 

Transport Under Unsaturated Moisture Contents. PNL-10722, Pacific Northwest 


Looney, B. B., M. W. Grant, and C. M. King. 1987. Estimating of Geochemical Parameters for 

Assessing Subsurface Transport at the Savannah River Plant. DPST-85-904, Environmental 

Information Document, E. I. du pont de Nemours and Company, Savannah River Laboratory, 

Aiken, South Carolina. 

Manskaya, S. M., G. V. Drozdora, and M. P. Yelmel’yanova. 1956. “Fixation of Uranium by 

Humic Acids and Melanoidins.” Geokhimiya, No. 4. 

Masuda, K., and T. Yamamoto. 1971. “Studies on Environmental Contamination by Uranium. 

II. Adsorption of Uranium on Soil and Its Desorption.” Journal of Radiation Research, 

12:94-99. 

McKinley, J. P., J. M. Zachara, S. C. Smith, and G. D. Turner. 1995. “The Influence of Uranyl 

Hydrolysis and Multiple Site-Binding Reactions on Adsorption of U(VI) to Montmorillonite.” 

Clays and Clay Minerals, 43(5):586-598. 

McKinley, G., and A. Scholtis. 1993. “A Comparison of Radionuclide Sorption Databases Used 

in Recent Performance Assessments.” Journal of Contaminant Hydrology, 13:347-363. 

Morris, D. E., C. J. Chisholm-Brause, M. E. Barr, S. D. Conradson, and P. G. Eller. 1994. 

2+ 

“Optical Spectroscopic Studies of the Sorption of UO2 Species on a Reference Smectite.” 


Neiheisel, J. 1983. Prediction Parameters of Radionuclide Retention at Low-Level Radioactive 

Waste Sites. EPA 520/1-83-025, U.S. Environmental Protection Agency, Washington, D.C. 

Payne, T. E., and T. D. Waite. 1991. “Surface Complexation Modelling of Uranium Sorption 

Data Obtained by Isotope Exchange Techniques.” Radiochimica Acta, 52/53:487-493. 

Puigdomènech, I., and U. Bergström. 1994. Calculated Distribution of Radionuclides in Soils 

and Sediments. SKB Technical Report 94-32, Swedish Nuclear Fuel and Waste Management 

Company, Stockholm, Sweden. 

J.54

Puls, R. W., L. L. Ames, and J. E. McGarrah. 1987. Sorption and Desorption of Uranium, 

Selenium, and Radium in a Basalt Geochemical Environment. WHC-SA-0003-FP, 

Westinghouse Hanford Company, Richland, Washington. 

Rançon, D. 1973. The Behavior in Underground Environments of Uranium and Thorium 

Discharge by the Nuclear Industry. In Environmental Behavior of Radionuclides Released in 

the Nuclear Industry, pp. 333-346. IAEA-SM-172/55, International Atomic Energy Agency 

Proceedings, Vienna, Austria. 

Ritchie, J. C., P. H. Hawks, and J. R. McHenry. 1972. “Thorium, Uranium, and Potassium in 

Upper Cretaceous, Paleocene, and Eocene Sediments of the Little Tallahatchie River 

Watershed in Northern Mississippi.” Southeast Geology, 14:221-231. 

Rozhkova, Ye.V., Ye. G. Razumnaya, M. B. Serebrayakova and O. V. Shchebak. 1959. “Role 

of Sorption in Concentration of Uranium in Sedimentary Rocks.” Tr. II. Mezhdunar, knof. 

po miro nmu ispol’z. atom. energii. 3. 

Rubtsov, D. M. 1972. “Thorium and Uranium Content in the Clay Fraction of Podzolic 

Mountain Soils of Thin Forests.” Radioekol. Issled Prir. Biogeotsenozakh, 53-66 (in 

Russian). 

Salter, P. F., L. L. Ames, and J. E. McGarrah. 1981. The Sorption Behavior of Selected 

Radionuclides on Columbia River Basalts. RHO-BWI-LD-48, Rockwell Hanford 


Seeley, F. G., and A. D. Kelmers. 1984. Geochemical Information for the West Chestnut Ridge 

Central Waste Disposal Facility for Low-Level Radioactive Waste. ORNL-6061, Oak Ridge 

National Laboratory, Oak Ridge, Tennessee 

Serkiz, S. M. And W. H. Johnson. 1994. Uranium Geochemistry in Soil and Groundwater at the 

F and H Seepage Basins (U). EPD-SGS-94-307, Westinghouse Savannah River Company, 

Savannah River Site, Aiken, South Carolina. 




Characterization. PNL-8889, Volume 1, Pacific Northwest Laboratory, Richland, 

Washington. 

Sheppard, M. I., D. I. Beals, D. H. Thibault, and P. O’Connor. 1984. Soil Nuclide Distribution 

Coefficients and Their Statistical Distribution. AECL-8364, Chalk River Nuclear Labs, 

Atomic Energy of Canada Limited, Chalk River, Canada. 

J.55

Sheppard, M. I., and D. H. Thibault. 1988. “Migration of Technetium, Iodine, Neptunium, and 

Uranium in the Peat of Two Minerotrophic Mires.” Journal of Environmental Quality, 

17:644-653. 

Sheppard, M. I., and D. H. Thibault. 1990. “Default Soil Solid/Liquid Partition Coefficients, 

K ds, for Four Major Soil Types: A Compendium.” Health Physics, 59(4)471-482. 

Starik, I. Ye., F. Ye Starik and A. N. Apollonova. 1958. “Adsorption of Traces of Uranium on 

Iron Hydroxide and Its Desorption by the Carbonate Method.” Zh. Neorgan. Khimii. 3(1). 

Stenhouse, M. J., and J. Pöttinger. 1994. “Comparison of Sorption Databases Used in Recent 

Performance Assessments Involving Crystalline Host Rock.” Radiochimica Acta, 

66/67:267-275. 

Stumm, W., and J. J. Morgan. 1981. Aquatic Chemistry. An Introduction Emphasizing 

Chemical Equilibria in Natural Waters. John Wiley and Sons, New York, New York. 

Szalay, A. 1954. “The Enrichment of Uranium in Some Brown Coals in Hungary.” Acta Geol. 

Acad. Sci. Hungary, 2:299-311. 

Szalay, A. 1957. “The Role of Humus in the Geochemical Enrichment of U in Coal and Other 

Bioliths.” Acta Phys. Acad. Sci. Hungary, 8:25-35. 



AECL-10125, Whiteshell Nuclear Research Establishment, Atomic Energy of Canada 

Limited, Pinawa, Canada. 

Tripathi, V. S. 1984. Uranium(VI) Transport Modeling: Geochemical Data and Submodels. 


Tsunashima, A., G. W. Brindley, and M. Bastovanov. 1981. “Adsorption of Uranium from 

Solutions by Montmorillonite: Compositions and Properties of Uranyl Montmorillonites.” 


Turner, D. R. 1993. Mechanistic Approaches to Radionuclide Sorption Modeling. CNWRA 

93-019, Center for Nuclear Waste Regulatory Analysis, San Antonio, Texas. 

Turner, D. R. 1995. Uniform Approach to Surface Complexation Modeling of Radionuclide 

Sorption. CNWRA 95-001, Center for Nuclear Waste Regulatory Analysis, San Antonio, 

Texas. 

J.56

Turner, D. R., T. Griffin, and T. B. Dietrich. 1993. “Radionuclide Sorption Modeling Using the 

MINTEQA2 Speciation Code.” In Scientific Basis for Nuclear Waste Management XVI, 

(eds.) C. G. Interrante and R. T. Pabalan, Materials Research Society Symposium 

Proceedings, Volume 294, p. 783-789. Materials Research Society, Pittsburgh, Pennsylvania. 

Turner, G. D., J. M. Zachara, J. P. McKinley, and S. C. Smith. 1996. “Surface-Charge 

2+ 

Properties and UO2 Adsorption of a Subsurface Smectite.” Geochimica et Cosmochimica 

Acta, 60(18):3399-3414. 

Vochten, R. C., L. van Haverbeke, and F. Goovaerts. 1990. “External Surface Adsorption of 

Uranyl-Hydroxo Complexes on Zeolite Particles in Relation to the Double-Layer Potential.” 

Journal of the Chemical Society. Faraday Transaction, 86:4095-4099. 

Waite, T. D., T. E. Payne, J. A. Davis, and K. Sekine. 1992. Alligators Rivers Analogue 

Project. Final Report Volume 13. Uranium Sorption. ISBN 0-642-599394 

(DOE/HMIP/RR/92/0823, SKI TR 92:20-13. 

Waite, T. D., J. A. Davis, T. E. Payne, G. A. Waychunas, and N. Xu. 1994. “Uranium(VI) 

Adsorption to Ferrihydrite: Application of a Surface Complexation Model.” Geochimica et 

Cosmochimica Acta, 58(24):5465-5478. 

Warnecke, E., G. Tittel, P. Brennecke, G. Stier-Friedland, and A. Hollman. 1986. 

“Experimental Investigations of Possible Radionuclide Releases from the Planned Repositories 

in the Gorleben Salt Dome and Konrad Iron ore Mine as Part of the Long-Term safety 

Assessment.” In Site, Design and Construction of Underground Repositories for Radioactive 

Wastes, IAEA-SM-289/49, p. 401-416, International Atomic Energy Agency, Vienna, 

Austria. 

Warnecke, E., A. Hollman, G. Tittel, and P. Brennecke. 1994. “Gorleben Radionuclide 

Migration Experiments: More Than 10 Years of Experience.” In Fourth International 

Conference on the Chemistry and Migration Behavior of Actinides and Fission Products in 

the Geosphere, p. 821-827, R. Oldenbourg Verlag, München, Germany. 

Warnecke, E., and W. Hild. 1988. “German Experience in the Field of Radionuclide Migration in 

the Geosphere.” Radioactive Waste Management and the Nuclear Fuel Cycle, 

10(1-3):115-144. 

Warnecke, E., A. Hollman, and G. Stier-Friedland. 1984. “Migration of Radionuclides: 

Experiments Within the Site Investigation Program at Gorleben.” In Scientific Basis for 

Nuclear Waste Management VII, (ed.) G. L. McVay, Materials Research Society Symposium 

Proceedings, Volume 26, p. 41-48. North-Holland, New York, New York. 

Yakobenchuk, V. F. 1968. “Radioactivity and Chemical Properties of Sod-Podzolic Soils in the 

Ukrainian Western Polesie.” Visn. Sil’s Kogosped. Nauki, 11:45-50 (in Ukrainian). 

J.57

Yamamoto, T., E. Yunoki, M. Yamakawa, and M. Shimizu. 1973. “Studies on Environmental 

Contamination by Uranium. 3. Effects of Carbonate Ion on Uranium Adsorption to and 

Desorption from Soils.” Journal of Radiation Research, 14:219-224. 

Zachara, J. M., C. C. Ainsworth, J. P. McKinley, E. M. Murphy, J. C. Westall, and P. S. C. Rao. 

1992. "Subsurface Chemistry of Organic Ligand-Radionuclide Mixtures.” In Pacific 

Northwest Laboratory Annual Report for 1991 to the DOE Office of Energy Research. Part 

2: Environmental Science, pp. 1-12. PNL-8000 Pt. 2, Pacific Northwest Laboratory, 


J.58

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