v>EPA
United Stcilos
Environmental Protection
Agency
Office of Hatiiatiun programs
Las Vegas Facility ^
P.O. Box 15027
Las Vegas NV 89114
'ZfH jtu/ ui- /'
August 1978
Radiation
Radionuclide interactions
with Soil and Rock Media
Volume 1:
Processes Influencing Radio-
nuclide Mobility and Retention
Element Chemistry and
Geochemistry
Conclusions and Evaluation
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EPA 520/6-78-007
August 1978
Volume 1 of 2
RADIONUCLIDE INTERACTIONS WITH SOIL AND ROCK MEDIA
Volume 1: Processes Influencing Radionuclide Mobility and Retention
Element Chemistry and Geochemistry
Conclusions and Evaluation
by
L. L. Ames
Dhanpat Rai
Battelle
Pacific Northwest Laboratories
Richland, Washington 99352
Final Report for Contract 68-03-2514
Project Officer
Robert F. Kaufmann
Evaluation Branch
Office of Radiation Programs-Las Vegas Facility
U.S. Environmental Protection Agency
Las Vegas, Nevada 89114
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PREFACE
The Office of Radiation Programs carries out a national program designed
to evaluate the exposure of man to ionizing and nonionizing radiation, and to
promote the development of controls necessary to protect the public health
and safety and assure environmental quality. A current area of great interest
and effort within the EPA is the development of environmental protection cri-
teria and standards for the management of radioactive wastes. Accidental or
deliberate interactions of such wastes with various geologic media dictate
that the Agency have access to current knowledge concerning the occurrence and
mobility of selected radionuclides in the lithosphere. Herein lies the purpose
and scope of the present study.
Office of Radiation Programs technical reports allow comprehensive and
rapid publishing of the results of intramural and contract projects. The
reports are distributed to groups who have known interests in this type of
information. The study reported on herein -is expected to be of considerable
interest to the Department of Energy, the U.S. Geological Survey, the Nuclear
Regulatory Commission, counterpart organizations in foreign countries facing
similar nuclear-related issues, selected private consulting and environmental
groups, and factions within industry. Ready availability of technical reports
to the scientific community as a whole and to the public is made possible by
distribution through the National Technical Information Service.
Readers of this report are encouraged to inform the Office of Radiation
Programs of any omissions or errors by contacting the Director, Office of
Radiation Programs - Las Vegas Facility, U.S. Environmental Protection
Agency, Las Vegas, Nevada 89114.
W. D. Rowe, Ph.D.
Deputy Assistant Administrator
for Radiation Programs
111
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CONTENTS
Volume 1
EPA REVIEW NOTICE , • . 11
PREFACE - 111
ABSTRACT . . iv
TABLE OF CONTENTS v
LIST OF FIGURES x11
LIST OF TABLES ..... xv
SECTION 1 - INTRODUCTION 1-1
METHODS AND CONSTRAINTS- 1-2
Selection of Radionuclides 1-2
Collection of Pertinent Literature ..1-4
REFERENCES. . . 1-5
SECTION 2 - PROCESSES INFLUENCING RADIONUCLIDE MOBILITY
AND RETENTION 2-1
NATURAL SOIL AND ROCK DISTRIBUTIONS 2-1
SOLID PHAS£ AND SOLUTION SPECIES 2-2
Uncomplexed, Mononuclear and Polynuclear Solution Species . . .2-4
Construction and Interpretation of the Diagrams 2-5
ION EXCHANGE 2-6
Ion Exchange Properties of Soil' and Rock Components ...... 2-8
ANION EXCLUSION 2-13
DIFFUSION 2-13
REPLACEMENT REACTIONS 2-15
PHYSICAL TRANSPORT AND FILTRATION 2-17
SATURATION EFFECTS 2-19
SPECIFIC RETENTION 2-25
THE DISTRIBUTION COEFFICIENT OR Kd 2-27
Kd Relationship to Migration 2-30
Field Determination of Kd Values 2-31
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CONTENTS (contd)
REFERENCES 2-33
SECTION 3 - ELEMENT CHEMISTRY AND GEOCHEMISTRY 3-1
AMERICIUM 3-1
Natural Soil and Rock Distributions 3-1
Brief Chemistry 3-1
Solid Phase and Solution Equilibria 3-2
Experimental Adsorption Results 3-4
Migration Results .... 3-8
Summary 3-10
References. . .3-11
ANTIMONY 3-12
Natural Soil and Rock Distributions 3-12
Brief Chemistry 3-13
Solid Phase and Solution Equilibria 3-14
Experimental Adsorption Results 3-15
Migration Results 3-22
Summary. 3-24
References 3-25
CERIUM 3-27
Natural Soil and Rock Distributions 3-27
Brief Chemistry 3-28
Solid Phase and Solution Equilibria 3-28
Experimental Adsorption Results 3-30
Migration Results 3-32
Summary 3-34
References 3-35
CESIUM . . . '. 3-37
Natural Soil and Rock Distributions 3-37
Brief Chemistry . 3-38
Solid Phase and Solution Equilibria 3-38
Experimental Adsorption Results 3-38
Migration Results 3-55
Summary 3-62
References 3-63
vi
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CONTENTS (contd)
COBALT 3-69
Natural Soil and Rock Distributions 3-69
Brief Chemistry 3-69
Solid Phase and Solution Equilibria 3-70
Experimental Adsorption Results 3-72
Migration Results 3-74
Summary. ..... 3-75
References 3-76
-\
CURIUM 3-78
Natural Soil and Rock Distributions 3-78
Brief Chemistry 3-78
Solid Phase and Solution Equilibria 3-79
Experimental Adsorption Results 3-81
Migration Results 3-82
Summary. 3-83
References. 3-84
EUROPIUM ........... . . . 3-85
Natural Soil and Rock Distributions . 3-85
Brief Chemistry 3-85
Solid Phase and Solution Equilibria 3-86
Experimental Adsorption Results 3-88
Migration Results 3-91
Summary 3-91
References 3-92
IODINE - 3-94
Natural Soil and Rock Distributions . . . 3-94
Brief Chemistry 3-95
Solid Phase and Soil Solution Equilibria 3-96
Experimental Adsorption Results 3-96
Migration Results 3-97
Summary 3-98
References 3-99
vii
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• CONTENTS (contd)
NEPTUNIUM . . : 3-101
Natural Soil and Rock Distributions 3-101
Brief Chemistry 3-101
Solid Phase and Solution Equilibria 3-101
Experimental Adsorption Results 3-103
Migration Results .... 3-106
Summary 3-106
References. . • 3-107
PLUTONIUM 3-108
Natural Soil and Rock Distributions 3-108
Brief Chemistry 3-108
Solid Phase and Solution Equilibria . . • 3-109
Experimental Adsorption Results 3-115
.Migration Results . . 3-126
Summary 3-132
References." .".... 3.-133
PROMETHIUM " .' . . 3-138
Natural Soil and Rock Distributions '3-138
Brief Chemistry 3-138
Solid Phase and Soil Solution Equilibria 3-138
Experimental Adsorption Results 3-139
Migration Results 3-140
Summary 3-140
References . . . 3-141
RADIUM 3-142
Natural Soil and Rock Distributions 3-142
Brief Chemistry 3-142
Solid Phase and Solution Equilibria "... 3-143
Experimental Adsorption Results 3-143
Migration Results 3-145
Summary 3-147
References 3-147
vi i i
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Volume 2 - Part A
EPA REVIEW NOTICE 11
PREFACE . . . . 111
ABSTRACT 1v
TABLE OF CONTENTS v
LIST OF ABBREVIATIONS vl
BIBLIOGRAPHIC CITATIONS (A THROUGH Me) -. . A-l
Volume 2 - Part B
EPA REVIEW NOTICE 11
PREFACE Ill
ABSTRACT . . 1v
TABLE OF CONTENTS v
BIBLIOGRAPHIC CITATIONS (M THROUGH Z) ... M-l
ALPHABETICAL INDEX TO JUNIOR AUTHORS . Index-!
DOCUMENT INDEX Index-121
GEOGRAPHIC INDEX ...... Index-181
SUBJECT INDEX Index-201
xi
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FIGURES
Number Page
2-1 Types of column breakthrough curves .2-25
3-1 The relative stability of various americium solids in an
oxidizing soil environment 3-3
3-2 The activity of various americium species in equilibrium
with Am02(s) in em oxidizing soil environment . 3-4
3-3 The activity of SbO in equilibrium with various antimony
solids in an oxidizing soil solution environment 3-14
3-4 The activity of various antimony species in equilibrium
with Sb203(s) (cubic) with pNH4+ = 3.0 and pp- = 4.5 3-16
3-5 The activity of Ce in equilibrium with phosphate
levels from Variscite and Gibbsite (V and G), Dicalcium
Phosphate Dihydrate (DCPD) and Octacalcium Phosphate
(OCR). ..'.... . . .3-29
3-6 The activity of various cerium species in equilibrium
with CeP04(s) in an oxidizing soil 'environment 3-30
3-7 The relative stability of cobalt solids in an oxidizing soil
environment [p02(g) = 0.68 atm], pCa2+ = 2.5 and phosphate
levels in equilibrium with Variscite and Gibbsite (V and G),
Dicalcium Phosphate Dihydrate (DCPD) and Octacalcium
Phosphate (OCP) 3-71
3-8 The activity of various cobalt ions in equilibrium with
CoC03(s) in an oxidizing soil environment 3-71
3-9 The activity of various curium ion species in equilibrium
with Cm(OH)3(s) 3-80
3-10 The relative stability of various europium solids at
pSOf- = 2.5 ' 3-87
3-11 The .activity of various europium species in equilibrium
with Eu(OH)3(s) 3-87
3-12 The influence of pH on the activity of europium in
solution 3-89
3-13 The relative stability of various neptunium solids in
an oxidizing soil environment 3-102
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CONTENTS (contd)
RUTHENIUM 3-150
Natural Soil and Rock Distributions 3-150
Brief Chemistry 3-151
Solid Phase and Solution Equilibria 3-152
Experimental Adsorption Results 3-154
Migration Results 3-157
Summary. 3-162
References. 3-163
STRONTIUM ....... 3-166
Natural Soil and Rock Distribution 3-166
Brief Chemistry 3-166
Solid Phase and Solution Equilibria 3-167
Experimental Adsorption Results 3-169
Migration Results . 3-188
Summary. 3-197
References 3-198
TECHNETIUM. .' .. - 3-204
Natural Soil and Rock Distribution. 3-204
Brief Chemistry 3-204
Solid Phase and Solution Equilibria 3-205
Experimental Adsorption Results 3-205
Migration Results . 3-207
Summary 3-209
References 3-210
THORIUM 3-211
Natural Soil and Rock Distributions 3-211
Brief Chemistry 3-212
Solid Phase and Solution Equilibria . . 3-213
Experimental Adsorption Results 3-215
Migration Results 3-216
Summary 3-217
References. 3-218
ix
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TRITIUM 3-221
Natural Soil and Rock Distributions 3-221
Brief Chemistry 3-221
Solid Phase and Solution Equilibria 3-222
Experimental Adsorption Results 3-222
Migration Results 3-223
Summary 3-226
References 3-226
URANIUM 3-228
Natural Soil and Rock Distributions 3-228
Brief Chemistry 3-228
'Solid Phase and Solution Equilibria 3-230
Experimental Adsorption Results. . 3-232
Migration Results 3-235
Summary 3-238
References 3-239
ZIRCONIUM . . 3-243
Natural Soil and Rock Distributions . . . 3-243
Brief Chemistry . 3-243
Solid Phase and Solution Equilibria 3-245
Experimental Adsorption Results 3-247
Migration Results 3-248
Summary 3-250
References 3-251
SECTION 4 - CONCLUSIONS AND EVALUATION 4-1
FACTORS INFLUENCING ADSORPTION OF RADIONUCLIDES ON
GEOLOGIC MEDIA 4-2
ADDITIONAL RESEARCH NEEDS 4-6
ACKNOWLEDGMENTS . Ack-1
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FIGURES (contd)
Number Page
3-14 The activity of various neptunium species in equilibrium with
Np02(s) in an oxidizing soil environment ......... 3-103
3-15 The relative stability of various plutonium solids in an
oxidizing environment Cp02(g) = 0.68 atm], pC02(q) =
3.52 atm, pF- = 3.5 and phosphate levels in equilibrium
with Variscite and Gibbsite (V&G) Dicalcium Phosphate
Dihydrate (DCPD) and Octacalcium Phosphate (OCP) ...... 3-110
3-16 The relative stability of various plutonium solids in a
very reducing environment Cp02(g) = 80 atm], pC02(q) =
3.52 atm, pF° = 3.5 and phosphate levels in equilibrium
with Variscite and Gibbsite (V&G) Dicalcium Phosphate
Dihydrate (DCPD) and Octacalcium Phosphate (OCP) ..... : 3-110
3-17 The activity of various plutonium species in soil solution
in equilibrium with Pu02(s) at pH 8, pC02 - 3.52 atm, pCl" =
- = 2.5 and pH2P04- = 5.0 ............ 3-112
3-18 The activity of various plutonium species in soil solution
in Pu02(s) in a mildly oxidizing soil environment ...... 3-113
3-19 The relative stability of various promethium solids ..... 3-139
3-20 The relative stability of various ruthenium solids in an
oxidizing soil environment ............. 3-153
3-21 The activity of various ruthenium species in equilibrium
with RuOg (amorphous hydrate) in an oxidizing soil
environment .................. 3-153
3-22 The relative stability of various strontium solids at
pCa2+ = pS042- = 2.5, pH4Si04 =3.1, and pC02(g) =1-52
atm in equilibrium with Variscite and Gibbsite (V and G),
Dicalcium Phosphate Dihydrate (DCPD) and Octacalcium
Phosphate (OCP) .................. 3-168
3-23 The activities of various strontium species in equilibrium
with SrC03(s) in the soil at pCl = pS042' =2.5, pNOa- =
3.0, pF- = 4.5, pC02(g) » 1.52 atm and pH2P04" = 5.0 ..... 3-169
3-24 The relative stability of various thorium solids in
equilibrium with Variscite and Gibbsite (V&G) Dicalcium
Phosphate Dihydrate (DCPD) and Octacalcium Phosphate (OCP). . . 3-213
3-25 Activity of various thorium species in equilibrium
with Th02(s), pCl" = pNOs" = pS042' = 3.0, pF~. =
4.5 and HP0- = 5.0 ............... 3-214
3-26 The relative stability of various uranium solids in an
oxidizing soil environment [p02(g) = 0.68 atm], PC02(g)
=3.52 atm, pK* = pNa+ = pNfy* = 3.0 and phosphate levels
in equilibrium with Variscite and Gibbsite ........ 3-231
3-27 Activity of various uranium species in equilibrium with
in an oxidizing soil environment ......... 3-232
xi i i
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FIGURES (contd)
Number Page
3-28 The stability of zirconium solids at 25°C in equilibrium
with soil silica (oH4Si04 =3.1) 3-245
3-29 The activity of various zirconium species in equilibrium
with zircon [ZrSi04(s)] and soil solution silica (pH4$i04
=3.1) at pCT = pS042- = 2,5, pNOs" = 3.0 and pF~ = 4.5. . . . 3-246
xiv
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TABLES
Number Page
1-1 LWR Plant Waste from Uranium Fuels Reprocessing 1-3
2-1 Cation Exchange Capacities . 2-10
2-2 Increase in Cation Exchange Capacities with Kaolinite
Particle Size ' 2-10
2-3 Increase in Cation Exchange Capacities for Some
Smectites with Particle Size 2-11
2-4 Cation Exchange Capacity in meq-100 g Due to Grinding . . . 2-11
2-5 Examples of Replacement Reactions 2-16
2-6 Soil Capacities for Saturated and Unsatured Flow
Conditions 2-20
2-7 Dispersion Coefficients in the Unsaturated Zone of
a Yellow Eolian Soil 2-23
3-1 Americium Radionuclide Data. . 3-2
3-2 Americium Kd Values from Several Organic Solutions .... 3-5
3-3 Americium 50-Day Distribution 3-8
3-4 Leaching of Americium Adsorbed on the Soil by
Organic Wastes 3-10
3-5 Antimony Concentration in Igneous and Sedimentary Rocks . . 3-13
3-6 Antimony Radionuclide Data 3-13
3-7 Radiochemical Analysis of Test Shot Debris Sampled
about 200-ft Above the Shot Point at about 2 Years
in Age 3-17
3-8 Composition of Safford Copper Ore 3-17
•toe
3-9 Solutions Used in Sb Equilibrations with Glacioflu-
viatile Sediments from Wells 3-18
3-10 Sediments Used in Obtaining Equilibrium Distribution
Coefficients with 125$b and the Solutions Given in
Table 3-9 3-19
3-11 Equilibrium Distribution Coefficient Values (ml/g)
Between Seven Glaciofluviatile Sediments and the
Solutions of Table 3-9 Containing 125Sb 3-19
xv
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TABLES (contd)
Number Page
•joe
3-12 Sb Analyses of Water and Soil Samples from a
Seep Near Trench 7 '. . 3-20
3-13 Chemical Compositions of Melt Glass and Primary
Copper Ore. . . 3-21
3-14 Composition of Leaching Solution Influent . . 3-22
3-15 Cerium Abundance in Rocks . . . . ._ 3-27
3-16 Stable Isotopes of Cerium 3-28
144
3-17 Adsorption of Ce on Suspended Solids ........ 3-34
3-18 Cesium Content of Rocks 3-37
3-19 Cesium Distribution Coefficients for Several Rocks,
Minerals.and Soil Types 3-40
3-20 ' The Equilibrating Solution Composition Plus Trace
137cs That Was Utilized in All of Berak's (1963)
Kd Work With 24 Hour Equilibrations at 20°C 3-49
3-21 Effect of Addition of Potassium to the Influence on
the Sorption of Cesium by Vermiculite from 0.5M NaN03
Containing the Mass Equivalent of 2 uc 13'Cs/ml 3-49
3-22 Effect of pH on 137Cs Adsorption 3-50
1 "37
3-23 Effect of NaNOs Concentration on Cs Adsorption 3-50
3-24 Summary of Cesium Kd Values 3-51
3-25 Variation of Trace Cesium Kd Values with pH 3-52
3-26 Solution:Soil Ratio Effect-on Trace Cesium
Equilibrium Distribution Coefficients 3-53
3-27 Cesium Equilibrium Distribution Coefficients for Sand,
Silt and Clay Fractions of Burbank Loamy Fine Sand in
0.5N NaCl 3-53
3-28 Trace Cesium Adsorption by Three Hanford Soils from
a 0.2N NaCl Solution 3-54
3-29 Cesium Kd Values Between Seven Glaciofluviatile Sediments
and the Solutions of Tables 3-9 and 3-10 3-55
3-30 K Values of Cesium in the Saturated Subsoils at
Mol, Belgium 3-59
3-31 Cesium Distribution Coefficients and Relative Migration
Rates in Houthulst Clays at pH 3 3-59
3-32 Properties of Natural Water Samples 3-61
3-33 Adsorption of Radionuclides on Suspended Solids 3-61
xvi
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TABLES (contd)
Number Page
3-34 Cobalt Concentrations in Igneous and Sedimentary
Rocks in ppm 3-69
3-35 Curium Radionuclide Data 3-79
3-36 Curium 50-Day Distribution Computed from
Sheppard et al 3-82
?4?
3-37 Kd Values for Cm in Untreated Aiken Clay Loam ..... 3-83
3-38 Europium Content of Igneous Rocks in ppm 3-85
3-39 Europium Content of Sedimentary Rocks in ppm 3-85
3-40 Europium Radionuclide Data 3-85
3-41 Variation of Trace Europium Kd Values on Mol Soils .... 3-88
3-42 Effect of pH and Europium Concentration on
Europium Kd by Burbank Sand 3-90
3-43 K Values of Europium in the Saturated Subsoils at
Mol Belgium . . 3-91
3-44 Iodine Concentration in Volcanic Rocks 3-94
3-45 Iodine in Sedimentary Rocks 3-95
3-46 Range of Surface Soil Properties Used in the Methyl
Iodide and Iodide Adsorptions 3-97
3-47 Kd Values and Correlations Relating Soil Properties
Iodide and Methyl Iodide Adsorption 3-98
3-48 Properties of Soil Samples 3-104
3-49 Neptunium Kd (ml/g) as a Function of Soil and Solution . . . 3-104
3-50 Neptunium 50-Day Distribution Values Computed from
Sheppard et al 3-105
3-51 Plutonium Radionuclide Data 3-108
3-52 Plutonium Adsorption Versus Particle Size . . .' . . . .3-116
3-53 Plutonium Kd Values for Quartz-of Various Particle
Sizes from A 0.001 N HN03" Plutonium Solution 3-116
3-54 Plutonium Kd Values for Different Minerals of Coarse
Clay Size from 0.001 N HN03-Plutonium Solution 3-117
3-55 Rhodes' Plutonium Kd Data as Reported by Evans 3-117
3-56 Removal of Plutonium from pH 7 Water by Several Soil
Minerals 3-120
3-57 Percent Plutonium Leached by Extractants 3-125
3-58 Plutonium Kd as a Function of .Oxidation State 3-125
xv n
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TABLES (contd) -
Number Page
3-59 . Kd- Values of 238Pu in Untreated Aiken Clay Loam 3-129
3-60 Average Radium Content of Various Rock Types 3-142
3-61 Radium Radionuclide Data '.. 3-143
3-62 Radium Kd Values from Lime-Neutralized Waste;
1.25 g Exchanger/1 of Waste ...'.'. 3-144
3-63 Simulated.River Water Composition .--3-144
3-64 Radium Distribution Coefficients with the Solution
of Table 3-63 3-145
3-65 Abundances of Ruthenium in Igneous Rocks and
Rock-Forming Minerals in ppm :..... .3-150
3-66 Ruthenium Contents of Shales . . . 3-151
3-67 Atomic Percentages of Stable Ruthenium Isotopes 3-151
3-68 Ruthenium Distribution Coefficients on a Hanford Soil
as a Function of pH . . : 3-154
3-69 Equilibrium Distribution Coefficients Between Seven
Glaciofluviatile Sediments and the Solutions in
Table 3-9 Containing 106RU . . .3-156
3-70 Properties of Natural Water Samples 3-160
3-71 Adsorption of 106Ru on Suspended Solids in
Natural Waters 3-160
3-72 Average Strontium Content of Igneous Rocks . ." . . . .3-166
3-73 Average Strontium Content in Sedimentary Rocks -
and Sediments 3-166
3-74 Strontium Kd Values for Western Soils 3-171
3-75 The Influence of Anions on the Strontium Kd in the
Hanford Composite Soil 3-171
3-76 Summary of Strontium Kd Values 3-17-3
3-77 Strontium Distribution Coefficients for Several Rock,
Mineral and Soil Types . . . : 3-174
3-78 Microstrontium-Macrocalcium Discrimination of Some
0.1 to 0.2 mm Rocks and Minerals . . .3-181
3-79 Variation of Trace Strontium Kd Values with pH 3-182
3-80 Composition of Burns, Mississippi, Montmorillonite .... 3-184
3-81 Comparison of Experimental and Computed Strontium
Kd Values for Burns Montmorillonite 3-184
xvm
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TABLES (contd)
Number Page
3-82 Strontium Kd Values 3-186
3-83 Characteristics of the Soils Used in Strontium Kd
Determinations .3-187
3-84 Trace Strontium Adsorption Characteristics from
0.2N NaCl , . 3-187
3-85 K Values of Strontium in the Saturated Subsoils
at Mol, Belgium 3-193
3-86 Strontium Distribution Coefficients and Relative
Migration Rates in Houthulst Clays at pH 3 . . . • . . .3-194
3-87 Technetium Content of Some Natural Materials . . . . . . 3-204
3-88 Range of Properties of 22 Surface Soils Used in the
Pertechnetate Adsorption Studies . 3-206
3-89 Properties of Soils Used in the Technetium
Adsorption Studies 3-206
3-90 Properties of Soil Samples 3-207
3-91 Technetium Kd as a Function of NaHCOi Concentration for
a South Carolina Soil Characterized in Table 3-90 .... 3-207
3-92 Properties of Bergland and Arveson Soils Used in
the "Tc Extraction Work 3-208
3-93 Extractability of 99Tc from Bergland and
Arveson Soils 3-208
3-94 Thorium Content of Common Rocks and Soils 3-211
3-95 Thorium Radionuclide Data 3-212
3-96 Abundances of Uranium in Natural Materials 3-229
3-97 Nuclear Properties of Natural Uranium Isotopes 3-230
3-98 Uranium Radionuclide Data > 3-230
3-99 Uranium Kd Values 3-235
3-100 Desorption of Uranium from Soils with Distilled Water
and Carbonate Solutions ' 3-236
3-101 The Average Zirconium Content of Rocks 3-243
3-102 Natural Isotopes of Zirconium and Their Abundances .... 3-244
3-103 Zirconium-Niobium Kd Values as a Function of pH 3-247
3-104 Adsorption of 95Zr and 95Nb on Suspended Solids
in Natural Waters 3-250
xix
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TABLES (contd)
Number Page
4-1 Factors Reported to Effect Adsorption of Radioelements
Over the pH Range of 4 to 9 4-3
4-2 Predominant Solution Species of Elements in a pH 4 to 9,
p02 0.68 to 80, pC02 1.52 to 3.52, pCl" = pNOs" - pS042' =
3.0, pF~ 4.5 and pH2P04~ 5.0 Environment Without Organic
Ligands . . 4-4
xx
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SECTION 1
INTRODUCTION
The nuclear industry, like many other industries, produces wastes. Part
of these wastes are radioactive and biologically hazardous without management.
In general, the hazard represented by the radioactive wastes decreases with
isolation and storage, due to radioactive decay with passing time. However,
some of these wastes contain radioisotopes of americium, curium, iodine,
neptunium, plutonium, technetium and uranium with half-lives of many thousands
of years, so that these isotopes will persist.
Selected radioactive wastes have been, and are being, disposed by land
burial. Barriers around storage and disposal sites are generally constructed
to safeguard against radioactive contamination of the environment by wind
erosion or by trespassing of unauthorized personnel. However, in the case of
accidental spills or breach of barriers, the radioactive elements can enter
the surrounding environment. Water is one of the main transporting agents in
terrestrial environments and will control dissemination of the radionuclides
below the ground surface. The biological retention and mobility of the radio-
nuclides in geologic regimes will be governed by the physical and chemical
characteristics of the geohydrologic system. The composition of the water in
the geohydrologic system is determined by the adsorption-desorption and
precipitation*dissolution reactions between the water and the solid phases
making up the soils, sediments and rocks. As one of the principal controls
of the concentration of radionuclides in the geohydrologic system at a given
time, the interactions of the radionuclides with soils and rocks in an aqueous
system become of primary interest to understanding radionuclide migration and
retention. As a consequence, the present work was undertaken to critically
review the literature available through 1976 concerned with radionuclide
interactions with soils and rocks. The objectives of this study were to:
1-1
-------
1. summarize and critically review the available information on radionuclide
interactions with the geologic media, and
2. determine the deficiencies in the available data and outline future work
required to correct these deficiencies.
Several reviews already are available that include radionuclide inter-
actions with geologic materials (Gera, 1975; Borg et al., 1976a; Dames and
Moore, 1976; Routson, 1973; Bensen, 1960). However, these reviews do not
cover but a limited number of elements, refer to work done only in specific
geographical areas or their emphasis is not primarily the interactions between
geologic media and radionuclides.
METHODS AND CONSTRAINTS
The fission and activation products that are present in nuclear process-
ing and other radioactive wastes and are most biologically hazardous were
selected for study. For each selected element, pertinent available literature
on natural distribution in soils and rocks, thermodynamic data for solid and
solution species, chemistry of the element, radionuclide adsorption-desorption,
radionuclide precipitation and dissolution and the mobility of the element in
laboratory and field situations was collected. These data were organized and
used to critically review the present state of knowledge of the interactions
of radionuclides with geologic media.
Due to time (14 months) and level of effort (one man year) constraints,
the present study was limited in scope. Soil and rock reactions with a
selected group of radionuclides were thoroughly reviewed. Less effort was
spent on elemental chemistry and natural distribution. Available thermodynamic
data for various solids and solution species used in this study were not
critically evaluated by the present writers. However, the thermodynamic data
were checked for internal consistency. Laboratory work, field work and model-
ing were excluded from the study.
Selection of Radionuclides
Based upon radionuclide concentrations in radioactive wastes (Table 1-1),
19 potentially biologically hazardous elements were chosen for critical review
including americium, antimony, cerium, cesium, cobalt, curium, europium, iodine,
1-2
-------
TABLE 1-1. LWR PLANT WASTE FROM URANIUM FUELS REPROCESSING*
(SCHNEIDER AND PLATT, 1974)
SELECTED ACTINIDES. curies/metric ton of U
Radlonucl Ide
225Ra
226Ra
229Th
23°Th
233U
234U
237NP
239Np
239Pu
24°Pu
"'Am
243Am
242Cn,
245Cn
Separation
3.48 x 10"8
1.97 x 10"8
3.49 x 10"8
2.04 x 10"5
2.31 x 10"7
3.77 x 10"3
3.40 x 10"'
1.76 x lO1
1.62 x 10°
2.37 x 10°
1.52 x 102
1.76 x 101
1.93 x 10*
2.76 x 10"'
1 Year
3.50 x 10"8
2.85 x 10"8
'3.50 x 10"8
2.05 x 10"5
1.84 x 10"6
3.95 x ID"3
3.40 x 10"'
1.76 x 101
1.62 x 10°
2.63 x 10°
1.52 x 102
1.76 x 101
4.08 x 103
2.76 x 10"'
SELECTED FISSION
3H
79Se
90Sr
93Zr
95Zr.
99Tc
103Ru
"KRu
107Pd
124Sb
125Sb
126Sn
129j
'34Cs
135Cs
137Cs
141Ce
'47Pm
148Pm
151 Sm
152Eu
">
6.92 x 102
3.98 x 10"'
9.72 x 10*
7.68 x 104
1.46 x 10°
2.77 x 10S
1.43 x 101
8.83 x 104
4.10 x 10S
1.10 x 10"'
8.53 x 101
7.99 x 103
5.46 x 10"'
3.74 x 10"2
2.15 x 105
4.96 x 10"'
1.07 x 105
S.64 x 104
7.71 x 10S
9.80 x 104
3.19 x 101
1.47 x 103
1.14 x 101
6.76 x 103
9.84 x 103
6.54 x 102
3.98 x 10"'
7.48 x 102
7.50 x 104
1.46 x 10°
5.65 x 103
1.43 x 101
1.48 x 102
2.06 x 105
1.10 x 10"'
1.26 x 10°
6.18 x 103
5.46 x 10"'
3.75 x 10"2
1.53 x 105
4.96 x 10"'
1.04 x 105
2.29 x 101
3.16 x 105
7.52 x 104
7.71 x 10"2
1.46 x 103
1.07 x 101
6.47 x 103
6.71 x 103
10 Years
4.22 x 10"8
1.09 x 10"7
4.22 x 10"8
2.09 x 10"5
1.49 x 10"5
6.64 x 10"3
3.41 x 10"'
1.75 x 101
1.62 x 10°
4.52 x 10°
1.55 x 102
1.75 x 101
3.16 x 10°
2.76. x 10"'
100 Years
7.31 x 10"7
1.11 x 10"6
7.31 x 10"7
3.42 x 10"5
1.47 x 10"4
2.56 x 10"-2
3.45 x 10"'
1.74 x 101
1.67 x 10°
8.90 x 10°
1.44 x 102
1.74 x 101
2.09 x 10°
2.74 x 10"'
1 ,000 Years
6.94 x 10"5
7.22 x 10"5
6.94 x 10"S
3.86 x 10"4
1.52 x 10"3
4.56 x 10"2
3.67 x 10"'
1.60 x 101
2.05 x 10°
8.25 x 10°
3.43 x 101
1.50 x 101
3.46 x 10"2
2.56 x 10"'
10,000 Years
5.54 x 10"3
2.91 x 10"3
5.54 x 10"3
3.74 x 10"3
1.56 x ID"2
4.45 x 10"2
3.74 x 10"'
7.10 x 10?-
4.02 x 10°
3.28 x 10°
1.31 x 10"'
7.10 x 1.0°
5.24 x ID"20
1.31 x 10"'
PRODUCTS, curies/metric ton of U
3?94 x 102
3.98 x 10"'
7.05 x 10"17
6.01 x 104
1.46 x 10°
3.40 x 10"'2
1.43 x 101
1.54 x 10"23
4.15 x 102
1.10 x 10"'
4.08 x 10"17
6.14 x 102
5.46 x 10*'
3.75 x 10"2
7.30 x 103
4.96 x 10"'
8.48 x 104
0.00
1.04 x 102
6.96 x 103
2.15 x 10"25
1.36 x 103
6.38 x 10°
4.38 x 103
2.14 x 102
2.47 x 10°
3.98 x 10*'.
0.00
6.53 x 103
1.46 x 10°
0.00
1.43 x 101
0.00
4.58 x 10"25
1.10 x 10"'
0.00
5.70 x 10"8
5.46 x 10"'
3.75 x 10*2
4.47 x 10"'°
4.96 x 10"'
1.06 x 104
0.00
0.00
3.19 x 10"7
0.00
6.46 x 102
3.53 x 10"2
8.89 x 101
2.32 x 10"'3
2.35 x 10"22
3.94 x 10"1
0.00.
1.50 x 10"6
1.46 f 10°
0.00
1.43 x 101
0.00
0.00
1.10 x 10"1
0.00
0.00
5.42 x 10"'
3.75 x 10"2
0.00
4.95 x 10"'
9.92 x 10"6
0.00
0.00
0.00
0.00
5.12 x 10"'
9.45 x 10"25
1.05 x 10"15
0.00
0.00
1.37 x 10"'
0.00
0.00
1.45 x 10°
0.00
1.39 x 101
0.00
0.00
1.10 x 10''
0.00
0.00
5.10 x 10"'
3.74 x 10"2
0.00
4.94 x 10"'
0.00
0.00
0.00
0.00
0.00
0.00
0.00
0.00
0.00
•Assumptions: power 30 megawatts/metric ton, exposure 33,000 megawatt-days/metric ton,
3.3X enriched U, flux 2.92 x 10'J N/cm'-sec, spent fuel reprocessed
150 days after discharge, 0.5J fuel loss to waste.
1-3
-------
neptunium, plutonium, promethium, radium, ruthenium, strontium, technetium,
thorium, tritium, uranium and zirconium. This list includes most of the
elements that are of both immediate and long-term concern. Cobalt is not
listed in Table 1-1 because it is not an actinide or a fission product. It
is included here because as Co it is widely used in gamma irradiation
sources and occurs in other cobalt-containing materials that have been exposed
to a neutron flux. Cobalt can often be found in industrial and medical radio-
active wastes. Radium, a uranium and thorium decay chain product, is listed
with the parent actinides for convenience in Table 1-1. The selected radio-
nuclides, for example, also include the fission products and actinides calcu-
lated to be the most radiotoxic in the leachate from German high level waste-
borosilicate glasses (Scheffler et al., 1977).
Collection of Pertinent Literature
Because of limited time and resources available for the literature col-
lection and abstracting, an attempt was made to use already available biblio-
graphical data to critically review radionuclide interactions with geological
media. However, it soon became apparent that the available bibliographies
were not adequate (Environmental Plutonium Data Base Group, 1972, 1973;
Martin et aT.7~T97~
-------
fauna. The term soil is used in an engineering sense, and is not tied to any
specific depth such as the root zone. A sediment has a different origin than
a soil, but is considered here to be composed of essentially the same mate-
rials. A rock is a consolidated material composed of several mineral phases
and may also contain organic matter.
REFERENCES
Ames, L. L., Dhanpat Rai and R. J. Seme. 1976. A Review of Actinide-Sediment
Reactions with an Annotated Bibliography. BNWL-1983.
Bensen, D. W. 1960. Review of Soil Chemistry Research at Hanford. HW-67201.
Borg, I. Y., R. Stone, H. B. Levy and L. D. Ramspott. .1976a. Information
Pertinent to the Migration of Radionuclides in Ground Water at the Nevada
Test Site. Part 1: Review and Analysis of Existing Information. UCRL-52078,
Pt. 1.
Borg, I. Y., R. Stone, H. B. Levy and L. D. Ramspott. 1976b. Information
Pertinent to the Migration of Radionuclides in Ground Water at the Nevada
Test Site. Part 2: Annotated Bibliography. UCRL-52078, Pt. 2.
Dames and Moore. 1976. Development of Monitoring Programs for ERDA Owned
Radioactive Low-Level Waste Burial Sites.
Environmental Plutonium Data Base Group. 1972. Environmental Aspects of
Plutonium. A Selected, Annotated Bibliography. ORNL-EIS-72-21.
Environmental Plutonium Data Base Group. 1973. Environmental Aspects of
Plutonium and Other Elements. A Selected, Annotated Bibliography. ORNL-EIS-
73-21 (Suppl. 1).
Faust, R. A., F. M. Martin, C. T. Sanders, and S. S. Talmage (comps. and eds.).
1975. Environmental Aspects of the Transuranics: A Selected, Annotated
Bibliography. ORNL-EIS-75-21-5.
Francis, C. W., S. S. Talmage and B. B. McMullin. 1975. Radionuclide Move-
ment in Soils and Uptake by Plants. A Selected, Annotated Bibliography.
ORNL-EIS-75-77.
Gera, F. "1975. GeochemicaT Behavior of Long-Lived Radioactive Wastes.
ORNL-TM-4481.
Martin, F. M., C. T. Sanders and S. S. Talmage (eds.). 1974. Environmental
Aspects of the Transuranics. A Selected, Annotated Bibliography. ORNL-EIS-
74-21 (Suppl. 3).
Routson, R. C. 1973. A Review of Studies on Soil-Waste Relationships on the
Hanford Reservation from 1944 to 1967. BNWL-1464.
1-5
-------
Scheffler, K., U. Riege, K. Louwrier, Hj. Matzke, I. Ray and H. Thiele. 1977.
Long Term Leaching of Silicate Systems: Testing Procedure, Actinides Behavior
and Mechanism. KFK 2456. EUR 5509e, p. 29.
Schneider, K.
Alternatives.
2.D.22.
J. and A. M. Platt (eds.). 1974. High-Level Waste Management
BNWL-1900, Volume I, Appendix 2.D, pp. 2.D.7-2.D-8, 2.0.21-
1-6
-------
SECTION 2
PROCESSES INFLUENCING RADIONUCLIDE
MOBILITY AND RETENTION
This section is concerned with a general discussion of some of the factors
that are expected to influence the radionuclide-geologic media interactions.
All of the factors discussed here may or may not influence the behavior of a
specific element. The influence of these factors on various selected elements
is the subject of the next section. The main purpose of this section is to:
1) briefly enumerate the factors and processes involved in radionuclide-
geologic media interactions in the next section,
2) define various terms, and
3) outline-the reasons for developing, limitations and the method of construc-
tion of various solid phase and solution species diagrams presented in
the next sections.
NATURAL SOIL AND ROCK DISTRIBUTIONS
It is important to determine the concentration of the stable isotopes
of a given element in the soils and rocks because the chemical behavior of
the stable isotopes would not differ from that of the radioisotopes of the
same element. Hence, the presence of stable isotopes can affect the mobility
and retention of the radioisotopes. For example, limestone contains an
90
average of 610 ppm strontium. Disposal or an accidental spill o'f Sr-contain-
ing waste in a limestone area would very likely add much less strontium to the
environment than was already present as stable isotopes. The stable isotopes
of strontium can be present in such abundance that they control the soil reac-
tions and mobility of the radioactive strontium.
2-1
-------
The average elementary abundances in rocks and soils are given in the
next section. The actual elemental concentrations may range ..widely from the
average values given due to local conditions. It is always necessary to
establish local base level concentrations to avoid disposal or storage of the
radioisotope of an element in environments with high concentrations of stable
isotopes of the same element. Such an environment could enhance mobility of
the radionuclide.
SOLID PHASE AND SOLUTION SPECIES
Among factors that influence radionuclide behavior in the environment,
the nature of the solid phases and solution species of the radionuclide is
one of the most important. Through their precipitation and dissolution solid
phases would cause changes in solution concentration which, in turn, would
affect adsorbed ions. The environmental concentration of most radionuclides
is expected to be low so that discrete solid phases of the element may not
be present in the terrestrial environment. In this case the element may be
distributed throughout the soil and rock matrix and adsorption reactions may
control solution concentrations. However, if the concentration of the element
in solution is such that a solid phase of the element can precipitate, the
knowledge of this solid phase and solution ions would help in predicting the
solution concentrations available for transport (Dutt, Shaffer, and Moore,
1972; Jurinak and Medrano, 1974; Rai and Franklin, 1973; Tanji and Doneen,
1966). Granted the importance of the total solution concentrations of the
element, the nature of the predominant solution species are important since
they affect 1) adsorption through their charge; 2) adsorption because of
changes in the nature of the species due to alteration in solution properties
such as pH, Eh, competing ions and complexing ions; 3) movement through the
soil and rock matrix because of their physical size; and 4) plant uptake.
Laboratory information on possible solid phase and solution species of
elements of interest in radioactive waste is almost nonexistent. What little
information is available is very site and solution specific with such poorly
defined systems that no extrapolations can be made to other situations. In order
to bridge this gap, thermodynamic data are generally used to construct Eh-pH
2-2
-------
diagrams to predict stable, phases and solution species (Pourbaix, 1966;
Garrels and Christ, 1965). This approach also has several limitations, but
is the best method until extensive laboratory data become available. Some of
the limitations of a theoretical approach are:
1. The reliability of the prediction is dependent upon the accuracy of the
thermodynamic data which, in a large number of elements of concern in
radioactive waste, contain a large error term.
2. The approach is limited to those compounds for which information is
available at the present time.
3. This approach does not consider the kinetics of reactions.
4. The predictions apply only under equilibrium conditions.
5. The influence of organic ligands and colloids which may be very important
cannot be evaluated because the thermodynamic data are lacking.
In order to develop solid phase-solution species diagrams, accurate values
of the equilibrium constants of all the reactants and products are required.
The time and money available for this study allowed neither an extensive search
nor a critical evaluation or tabulation of the available thermodynamic data.
In a related DOE sponsored document, Rai and Seme (1978), the equations and
thermodynamic data used to construct the solid phase-solution species diagrams
are tabulated. Therefore, equilibrium constants were selected from several
readily available sources, the most important of these sources being Sillen and
Martell (1964); NBS publications (Rossini et al., 1952; Wagman et al., 1968,
1969, 1971; Parker et al., 1971; Schumm, 1973), Latimer (1952); Pourbaix (1966);
Cleveland (1970); Keller (1971); Burney and Harbour (1974); and Baes and Mesmer
(1976). The thermodynamic data were used to develop equilibrium constants for
chemical equations that describe the behavior of various species. Generally
these equations are used to develop Eh-pH diagrams (Pourbaix, 1966; Garrels and
Christ, 1965) that provide the information regarding the stable compounds and
solution species in the whole range of pH and Eh. However, Eh-pH diagrams
can be used only for elements that exist in more than one oxidation state, and
these diagrams do not indicate the total amount of the element and the relative
amount of various species in solution under specified conditions. Therefore,
activity-pH diagrams rather than Eh-pH diagrams were developed. The solid
2-3
-------
phase-solution species diagrams of elements affected by oxidation-reduction
were developed at an assumed partial pressure of oxygen (0.2 atmospheres in
most cases). However, the changes in solution concentration or the stability
of the solid phases that would occur due to the decrease in oxygen pressures
are also discussed. This approach is similar to the one used by Lindsay (1972),
Baes and Mesmer (1976), and Rai and Serne (1977).
Uncomplexed, Mononuclear and Polynuclear Solution Species
An element in solution can exist as free ions or as mononuclear and poly-
nuclear species (complexed). The total concentration of the element in solution
can be obtained by summing the concentrations of all the free and complexed
ions. Mononuclear species are of the type ML, ML9, ML., ML , where M is
u Oo n
the metal ion and L is the complexing ligand such as C03~, HCOZ, F", NOl, Cl",
OH", CM", CgHcOy " (citrate) and C2H302~ (acetate). Some examples of mononu-
clear species are CoOH+, Co(OH)2 , Co(OH)3" and Co(OH)4 . Polynuclear species
are of the type MmLn, where m and n are integers. Some examples of polynuclear
species are Co2OH3+ and Co4(OH)44+.
In general, mononuclear species of the type ML maintain higher concentra-
tions in equilibrium with metal ions than the ML2 ML species. Hence, in
most cases mononuclear species of the type ML2 MLn do not contribute sig-
nificantly to total metal ion solution concentrations. Hydroxide bridging is
apparently responsible for polynuclear species of M(II), M(III) and M(IV) ions
(Baes and Mesmer, 1976). The polyanions of M(V) and M(VI) usually involve
oxygen bridging. In most cases, the polynuclear species do not contribute
significantly to total metal ion solution concentrations.
Polymerization of a radionuclide is a special case of hydrolysis and
precipitation, and is primarily a function of Eh, pH, temperature, radionuclide
concentration and the presence or absence of complexing ligands. For example:
T) Pu(IV) polymerizes but Pu(VI) and Pu(III) do not significantly polymerize
(Cleveland, 1970),
2) at 10 M Pu(IV), 91% of the plutom'um was present as polymer at pH 8.5
(Lindenbaum and Westfall, 1965), but true colloid was not present at
pH values of <7.5 for 10"8M Pu(IV) (Qrebenshchikova and Davydov, 1961),
and
2-4
-------
p
3) the presence of complexing ligands such as F", S04 , and oxidizing agents,
reduce the polymerization of plutonium (Cleveland, 1970).
The metal ions of interest in:this study that might be expected to poly-
merize include americium, antimony, cerium, cobalt, curium, europium, plutonium,
promethium, ruthenium, thorium, uranium and zirconium. Usually the polymeric
o o
units are colloidal in size range (MO A to 10,000 A) and possess a surface
charge that tends to keep the units from coagulating and precipitating from
solution. Rhodes (1957b) showed with laboratory studies of Kd versus solution
pH that 239Pu, 106Ru, 95Zr-95Nb and 144Ce Kd values all tended to diminish in
the pH region of 9 to 13, probably due to reduced surface change on the polymers
produced by the increasing amounts of NaOH added to reach the higher pH values,
resulting in reduced polymer adsorption by the soil. Prediction of whether
or not colloidal size polymeric units will form during a precipitation reaction
and remain stable in solution is difficult because of the large number of sys-
tem variables that effect possible polymer formation. However, the metal
cations with high charges (M , M ) show the greatest tendency to form
polymers.
The stability constants that were available in the literature,and appli-
cable to the soil and rock environments selected, were utilized in this study
to estimate solid phases and solution species equilibria. It is unfortunate
that reliable hydrolysis constants are not available for many of the elements
including plutonium. Also missing are reliable formational constants for many
other complex ion species such as the nitrosyl ruthenium complexes. In some
cases, there is a reasonable doubt as to the identity of some of the species,
including those formed with organic materials found in many soils and surface
waters.
Construction and Interpretation of the Diagrams
The details of the method of constructing the diagrams presented in this
report can be found in Rai and Lindsay (1975) and Rai and Serne (1977). In
these diagrams the pCOgCg), K , and pNa , etc., refer to the negative logarithm
of the activity (in moles/1 for solution species and atmospheres for gases).
The oxidation-reduction conditions are denoted by partial pressures of oxygen.
2-5
-------
In the solid phase diagrams, any compound that lies below another com-
pound at a given pH is the more stable. Thus, for any two solids at a specific
pH, the solid that maintains lower ionic activity is more stable. In addition
to stability, information regarding weathering and formation can also be
obtained from these diagrams. If, in a given soil, the soil solution composi-
tion is below the solid line, the soil is under-saturated with respect to
the solid represented by the line, and the mineral will dissolve or weather.
On the other hand, if a soil solution point is above the solid line, the soil
is supersaturated with respect to the mineral represented by the line, and
the mineral can precipitate.
The solution species diagrams were developed assuming an oxidation-
reduction environment, an equilibrium with a compound, and ionic activities
of various cations (Ca , Na+, NH/1") and anions (Cl~, HCCL, CO, , NO,", F",
n • , f J J J
SO, ", and P0» ") that may be- present in the terrestrial environment. As men-
tioned earlier, the organic ligands and colloids that may be present in the
solutions were not considered. The solution species diagrams depict anionic,
cationic, and several neutral or uncharged species for which the data were
available.
ION EXCHANGE
Ion exchange is a reversible interchange of ions between two compounds,
one of which is insoluble in the medium in which the exchange is carried out
(Amphlett, 1964). It is one of the most common mechanisms responsible for
radionuclide adsorption on geologic materials.
If X" is a solid, negatively-charged medium, an exchange reaction can
occur between two cations, -A and B as follows:
X"A+ + B"1" -JX'B** A+< .
The anion does not enter into the reaction in this case because the solid
medium has a residual negative charge that is inherent in the mineral frame-
work. In the case of clay minerals and zeolites, the negative charge is
+3 +4
acquired by substitutions of Al for Si in the mineral lattice.
2-6
-------
.A selectivity quotient or constant can be experimentally derived from
contacting the exchange medium, based in cation A , with a known concen-
tration of cation B . For univalent-univalent cation exchanges, the
selectivity coefficient,
KB _ [XV] [A+]
A [B+] [XV]. '
where the brackets denote solution and solid medium concentrations. For a
univalent-divalent exchange, the reaction would be 2X~A+ + B t(X") B + 2A+,
with the rational selectivity coefficient,
B *2
=
[A
A [B+2] [X V]2 ' .
If the capacity of the solid medium is known in meq/g (m ) and the initial
solution concentration in meq/ml (C ), a relationship between the selectivity
B +
coefficient, KjJ , and a distribution coefficient for B , KdB, can be derived.
Assuming that [A*] + [B*] = CQ and [xV] + [X~B+] = mQ then
B . CX"B*3 CA*1 _ [XV] (Co - [B+]) = . Co • [B+] .
. A [B+] [XV] [B+] (m0 - [X'B+]) m0 - [X'B]
If B is a trace component ([B ] much less than CQ, [X"B ] much less than
m ), then the above can be written as:
P m
KB % K(JB _?_ and KdB £ KB £ ,
0 0
or the Kd is inversely proportional to [A*] + [B+]. If Kd were plotted
versus log [A*], a straight line would result with a slope of minus one.
+2
A divalent cation, A , would yield a slope of minus two when plotted in
the above manner. The same treatment applies equally well to anion exchange
reactions.
One of the key parameters in the above treatment is the requirement that
the component for which the Kd value is determined be present in the system
2-7
-------
in trace' quantities. In practice, a trace quantity should be defined experi-
mentally by determining Kd values at several concentrations of the traced ion.
The Kd will remain constant, all other things being equal, over the traced ion
concentration range for which the Kd concept remains valid.
An implied limitation can be found in the original exchange reaction,
-++'-+ +
XA + B £ X B + A .. The expression requires that the solid medium, X , be
based in cation A before contacting the solution containing B . A Kd value
that has been experimentally determined for cation B without first. basing
the solid medium (soil, clay mineral, zeolite) in cation A , has been deter-
mined for an undefined system. Also, because the dimensions of m were meq/g,
and of C , meq/ml , the dimensions of the Kd value are ml/g.
If the selectivity quotient is modified with the activity coefficients
for A and B on the solid medium, f.-and fg, and in the solution y. and yg, a
rational thermodynamic equilibrium constant, K, can be derived (Gaines and
Thomas, 1953):
B
The thermodynamic quantities are sometimes of value in establishing and
studying reaction mechanisms, but are limited in practical applications. The
D
distribution coefficient (Kd) and the selectivity quotient (1C) where trace
concentrations are exceeded are experimentally determined values that are
characteristic of a given exchange system, and can be applied in several
practical situations.
Ion Exchange Properties of Soil and Rock Components
As far as soils, sediments and altered rocks are concerned, ion exchange
means chiefly cation exchange. The soil particles have an amphoteric charac-
ter, but in general they carry a net negative charge (Winklander, 1964). For
cation exchange, the negative charge on the soil colloids increases with rising
pH. The positive charge tends to increase with diminishing pH, giving rise to
an anion exchange capacity on the acidic side. Chloride and sulphate ions,
for example, show no adsorption at pH 7 on kanditic or smectitic soil minerals
(Mattson, 1931). Lowering of the pH activates basic groups, R-OH, to R-OHL,
2-8
-------
with the OH originating in clay mineral broken bonds and hydrous oxides of
iron, aluminum and manganese. They take part in exchange reactions such as,
R-OH^Cl" + NO^ -" R-OH^ NO^ + Cl~. However, the anion exchange capacities of
most soils are relatively minor compared to their cation exchange capacities
because the pH.of most soils is seldom as low as 4.0, and the number of sites
that can be activated to function as anion exchange sites usually are relatively
small as well.
There are at least two types of anion exchange in clay minerals. The
first involves replacement of the hydroxyl (OH") ions by other anions.
McAuliffe et al. (1947) showed that the OH ions of clays can enter into
exchange reactions by using clays with deuterium-tagged hydroxyls. The replace-
ment of hydroxyls by fluoride ions also was reported (Dickman and Bray, 1941). '
Another type of anion exchange involves the substitution of anions such
as phosphate, arsenate and borate on the edges of the (SiO,) (silica tetra-
hedron) sheets, perhaps growing as extensions of these sheets. Other anions
such as sulfate, nitrate and chloride are not adsorbed because their geometries
are not comparable to the silica tetrahedra.
A third, type of anion exchange site could occur as a positive site on
basal surfaces as a result of a local electrostatic imbalance in the clay
framework such as an excess of aluminum in the octahedral position.
In the cases of both OH ion replacement and anion exchange due to geo-
metrical similarities between the anion and silica tetrahedra, exchange would
occur around clay mineral edges. There would be very little exchange on basal
surfaces of expandable clay minerals. Investigations of anion exchange are
difficult because of the low pH requirement. However, for clay minerals the
cation and anion exchange capacities should be about equal at optimum pH
values for each.
The anion exchange capacities of expanding clays were 12 to 31 meq/100 g
for smectites and 4 meq/100 g for vermiculite, while kandites varied from
6 to 20 meq/100 g (Hofmann et al., 1956). Compare these anion exchange
capacity data with those of Table 2-1 for cation exchange at pH 6 to 8.
The cation exchange capacity ranges of some common clay minerals are
given in Table 2-1. The ranges vary due to differences in crystallinity and
2-9
-------
particle size. In soils, rocks, weather rocks and sediments, cation exchange
capacities are due mainly to the types and quantities of minerals in the silt
(2 to 62 urn) and clay (<2 urn) size fractions.
TABLE 2-1. CATION EXCHANGE CAPACITIES (CARROLL,. 1960, 1970)
: Cation Exchange Capacity,
Mineral meq/lOOg
. Kaolinite 3-15
Halloysite (2H20) - 5-10
Halloysite (4H20) ' 40-50
Smectites 36-100
Illite 10-40
Vermiculite 100-150
Chlorite 4-47
GlauconUe 11-20
Palygorskite 20-30
Allophane . 70
Zeolites 100-300
Organic Matter in Soils. ' 130-350
Feldspar <1
Quartz <1
There is a considerable variation in cation exchange capacity with particle
size shown by the non-expandable soil clays. Kaolinite, for example, shows an
increase in exchange capacity as outlined in Table 2-2. The expanding layer
silicates, on the other hand, show little change in cation exchange capacity
with particle size as shown for some smectites in Table 2-3.
TABLE 2-2.' INCREASE IN CATION EXCHANGE CAPACITIES WITH KAOLINITE
PARTICLE SIZE (HARMON AND FRAULINI, 1940)
Particle Size, urn
Cation Exchange 10-205-TOZ-41-0.50.5-0.250.25-0.10.1-0.05
Capacity, meq/lOOg 2.4 2.6 3.6 3.8 3.9 5.4 9.5
2-10
-------
TABLE 2-3. INCREASE IN CATION EXCHANGE CAPACITIES FOR SOME
SMECTITES WITH PARTICLE SIZE (CALDWELL AND
MARSHALL, 1942)
Particle Size, um
Cation Exchange Smectite 2-1 1-0.5 0.5-0.2 0.2-0.05 -0.05
Capacity, meq/IOOg Nontronite 60.8 61.0 64.35.70
Saponite 69.3 76.0 8.15 86.3 81.4
An increase is to be expected with decreasing particle size when the
cation exchange capacity primarily occurs from broken bonds in minerals such
as illite and kaolinite. In the .case of expanding layer silicates where most
of the capacity is on basal plane surfaces, particle size has little effect
on cation exchange capacity.
Excessive grinding of clay and nonclay minerals, causes an increase in
cation exchange capacity, probably a result of the same decrease in particle
size and number of broken bonds, as shown in Table 2-4. A rubber-lined ball
mill and polished agate balls were used for the grinding.
TABLE 2-4. CATION EXCHANGE CAPACITY IN meq/100 g DUE TO GRINDING
(KELLY AND JENNY, 1936)
Grinding Time, Days
100 Mesh ~2~ 3 7
Muscovite 10.5 - 76.0
Biotite 3.0 62.0 72.5
Kaolinite 8.0 57.5 70.4 100.5
Montmorillonite 126 - 238.0
The clay minerals are relatively sensitive to pH. Chatterjee and Paul
(1942) and Mukherjee et al. (1942) have shown that so-called hydrogen clays
are really hydrogen-aluminum or hydrogen-magnesium-aluminum systems, the mag-
nesium and aluminum being removed from the clay framework with resulting clay
damage. Other things being equal, higher valence cations more easily replace
lower valence cations, and are more difficult to replace when already present
on a clay or zeolite. Hydrogen ions are an exception because they behave
more like divalent or trivalent cations. With ions of the same valence, the
2-11
-------
smaller diameter hydrated cations are more tightly held than the larger diam-
eter hydrated cations. The usual lyotropic replacement series on smectites,
for example, is Cs > K > Na > Li and Ba > Sr > Ca > Mg . An excep-
tion arises when the cation just fits the size and charge requirements of the
anionic site, as does K into the basal sites of illite.
These cations can cause structural modifications in the clay or zeolite
structure that leads to their so-called "fixation" (Bray and DeTurk, 1939;
Wood and DeTurk, 1941; Nishita et al., 1956; Frysinger, I960; Tamura. and Jacobs,
1960; Tamura and Jacobs, 1961).
•In contrast to clay minerals, soil organic matter has no well-defined
ability to bind exchangeable cations, although the exchange displacement fol-
lows the lyotropic order, as do most of the clays (Bartlett and Norman, 1938).
The cation exchange capacity increases with increasing pH values, arising from .
the increasing dissociation of phenols and carboxylic acids at higher pH levels
(Broadbent and Bradford,.1952).
The cation exchange capacity of soil organic substances are reported as
between 150 and 450 meq/100 g depending on the origin of the fractions. Posner
(1966) showed-that the cation exchange capacities of humic acids extracted
from a red-brown soil under different conditions yielded substances with rela-
tively similar cation exchange capacities. Yields of humic acids varied from
2.7 to 12.0 g/kg soil and cation exchange capacities varied from 217 to 409
meq/100 g, with the averages at 7.0 g/kg soil for yield and 310 meq/100 g
capacity. Cation exchange isotherms between humic acids and alkaline earth
and alkali metals were investigated in detail by Zadmard (1939). The role of
carboxyl and hydroxyl groups in metal cation fixation by organic soil materials
was investigated by Schnitzer and Skinner (1965). Specific blocking of the
functional groups indicated two reaction types: 1) a main reaction with par-
ticipation of both carboxyl and phenolic groups, and 2) a lesser reaction
involving only carboxyl groups. 'No evidence was found for exchange reactions
of the carbonyl groups.
2-12
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ANION EXCLUSION
Reactions between an inorganic anion and a soil or rock exchanger can
involve both adsorption and repulsion. Adsorption of anions is due to a
limited number of positive charges on soil minerals (anion exchange capacity).
Anion exchange and the capacities of various minerals are discussed in the
ion exchange section above. Anion exclusion involves the repulsion of anions
by the predominantly negatively charged clay minerals and zeolites in soil
(Schofield, 1947; de Haan and Bolt, 1963; de Haan, 1964; de Haan, 1965). The
higher the cation exchange capacity (negative charge) of the soil or rock,
the more pronounced the anion exclusion (Thomas and Swoboda, 1970). The high
cation exchange capacity smectites exhibit anion exclusion to a greater degree
than do the kandites (kaolinite group). Thus anions, such as chloride, iodide
and pertechnetate, can be repelled by the negative charges of a high cation
exchange capacity soil into the higher velocity soil solution flow channels
and break through the soil column in less than one pore volume of displacing
solution (Thomas and Swaboda, 1970; Cassel et al., 1975). The implication
for anionic radionuclides and anionic complexes of radionuclides is that their
displacement through high cation capacity soil columns can take place at less
than one displacing solution volume, with a higher apparent migration velocity
than the displacing solution.
DIFFUSION
One of the parameters of aqueous transport of ions is that of aqueous
diffusion. The transport of matter in the absence of bulk flow is referred
to as the process of diffusion. The flux of matter due to diffusion is pro-
portional to the concentration gradient and is a molecular process. In
general terms, the flux, J. , of component i in the z direction, is (Fick's
dfy 1Z
first law) 0. = -D^-1- , where D is the proportionality constant and diffusion
1Z . QZ n
coefficient, with the dimensions of cm /sec, J. has the dimensions-of moles/
?
cm /sec, n. is the concentration of material i and the negative sign indicates
that flow of material i in the z direction is also in the direction of lower
i concentration. In a binary electrolyte, at infinite dilution, the aqueous
diffusion coefficient is:
2-13
-------
D =
'o (M-, + u2) F
where 'u, and Uo are the mobilities of the two ions (K and Cl", for example,
as a binary electrolyte former), RT is the gas constant times the degrees
o
Kelvin and F is a faraday. The dimensions of D are once again cm /sec. The
ionic mobilities of many ions are known (Harned and Owen, 1958; Dean, 1973)
in aqueous solutions. Scott et al . (1974) defined an apparent or measured.
difusion coefficient (D.) and its relation to the aqueous diffusion coefficient
(D ) as (the change of concentration with time in the x direction)
*r D« ^2r
9C o 3 C
- 3t " R 3X2 '
where R is a retardation factor.
The retardation factor, R, was defined as
R = 1 + £ Kd
£
where p is the soil bulk density, e is the pore fraction (volume occupied by
the solution phase) and Kd is the equilibrium distribution coefficient with
dimensions of ml/g. Note that the above expression is the same one that is
used' to evaluate the relationship between groundwater velocity and radionu-
clide migration velocity. The same retardation factor serves in both cases.
The apparent diffusion coefficient is defined by DA = DQ/R.
-8
Relyea (1977) obtained D values for 10" M plutonium nitrate in 0.01M
HNO,, 0.1M KNO, and 1M Ca(NO,), of 4.8 x 10~6 ± 2.5 x 10"6 cm2/sec, 7.8 x 10"5
5 ? 5 52
± 3.7 x 10 cm /sec and 3.0 x 10" ± 2.0 x 10" cm /sec, respectively.
These D values may be compared to the limiting diffusion coefficient of KC1 ,
-52
for example, of 2.0 x 10 cm /sec. Subsequent measurement of the plutonium
particle size in the 0.01M HNO, by means of the Stokes-Einstein equation
0 o
showed that it had rapidly hydrolyzed to particles about 8 A in diameter at
25°C. Increasing the nitric acid level to 0.1M reduced the effective radius of
o
plutonium particles to less than 1 A. The measured DQ from the Fuquay soil
supernate was 2.7 x 10"5 ± 1.2 x 10"5 cm2/sec at pH 2.0, 25°C and 8 x 10"10M
plutonium nitrate, and from the Burbank soil at pH 7.1 and plutonium 5 x 10" M,
2-14
-------
D = 3.6 x 10" ± 1.0 x 10~6 cm /sec. The effective plutonium radius for the
0 o
Burbank sample was 6 A. The D values of various molecules are remarkably
similar in the absence of polymerization, or some other complicating factor.
The plutonium D» values obtained by Relyea for the Burbank sandy loam at
' -in 7
e = 0.30, T = 21°C and 791 hours was 5 x 10 cnT/sec and at e = 0.14, T =
-11 2
21°C and 3600 hours was 8 x 10 cm /sec. The above DA values can be com-
pared to the previously-given D value for Burbank supernate to illustrate
the influence of the retardation factor (and Kd) on the mobility of high Kd
radionuclides. If, on the other hand, the radionuclide Kd value is very low,
the apparent diffusion coefficient (D.) approaches the aqueous diffusion
coefficient (D ) in value.
The simple expression, x" = ,/Vt (Jost, 1960) can be utilized to determine
the effect of D, the diffusion coefficient and t, time, on the distance (root
mean square distance) that a radionuclide will migrate. The dimension of D
2
is cm /sec and t also is in seconds with x in cm. If D. is substituted for D
in Jost's expression above, it can be seen that the high Kd radionuclides such
as plutonium must have very high t values in order to diffuse a significant
distance. However, the low Kd value radionuclides such as tritium or. techni-
tium, can diffuse distances of up to 25.cm/yr, due to D. values of close to
-52
1 x 10 cm /sec. Diffusion can be an important migration mechanism for the
low Kd radionuclides. Its potential contribution to migration should be
examined for each radionuclide.
REPLACEMENT REACTIONS
Replacement is a type of precipitation reaction in which one of the ions
of a final replacement product is contributed by a solid that is more soluble
in the contacting solution than is the final product. Replacement reactions
can be divided into two kinds: 1) anion replacement reactions in which only
the anion in the solid is replaced by an anion in solution, with cations of
both solution and solid being common, and 2) cation replacement reactions in
which only the cation in the solid is replaced by the a cation in solution,
with anions of both solution and solid being common. Some examples of anion
and cation replacement reactions are given in Table 2-5. The principle govern-
ing replacement reactions, assuming that the solution-solid system obeys Raoult's
law, is a relatively simple solubility relationship.
2-15
-------
TABLE 2-5. EXAMPLES OF REPLACEMENT REACTIONS
(AMES, 1961a; AMES, 1961b)
Original . Ion in
Solid Solution
Reaction
Solid
Product
Anion Replacement
CaCO^ F"
0
CaCO, P0«
•J *T
CaC03 Ba+2
CaC03 Sr+2
CaS04 2H20 Ba+2
CaS04 2H20 Sr+2
CaC03 + 2F" ^ CaF2 + C03"2
5 CaC03 + 3P04"3' + OH" -^
Ca5(P04)3OH + 5 C03"2
Cation Replacement
CaC03 + Ba+2 t=; BaC03 + Ca+2
CaC03 + Sr ^ SrC03 + Ca+
CaS04'2H2.0 + Ba+2 f^ BaS04 + Ca+2
CaS04-2H20 + Sr 2 f^ SrS04 + Ca+2
CaF2
Ca5(P04)3OH .
BaC03
SrC03
BaS04
SrS04
2H20
The difference in solubility between. the original so'lid and replacement
product is the driving force for the reactions shown in Table 2-5. The reac-
tion will cease as the ratio of replaced anion or cation to replacing anion or
cation approaches the same ratio as the activity products of the original solid
and solid product with the common anion or cation. For example, if the
activity products are 1.5 x 10"9 for BaS04 and 3.6 x 10"5 for CaS04«2H20 during
the replacement reaction CaSO,,*2H90 + Ba+2 ^ BaSO, + Ca*2 + 2H,0,.the
*
* 1.0 »
reaction rate approaches zero as the activity ratio of Ca to Ba+2 approaches
3 6 x 10~5 4
— - Q, or 2.4 x 10 . If excess calcium is added, the equilibrium can be
1.5 x 10-a
made to go to the left and form CaS04*2H20) (gypsum). When several solid
products are theoretically possible, only the least soluble will form.
Replacement reactions involve the dissolution of the original solid at
an interface between solid product and original solid (Ames, 1963). Bonds
between cation and anion in the original solid must actually be broken, the
ions diffuse into solution and recrystallize on the advancing solid product.
Ions must be able to diffuse through the solid product to the reaction face or
2-16
-------
the replacement reaction soon ceases. It is during the portion of the reaction
when ions are traveling between the more soluble original solid to the less
soluble final product that trace radionuclides able to satisfy the ion size
and charge requirements of the final product are included in the growing
reaction product (Ames, McHenry and Honstead, 1958; Ames, 1959; Ames, 1960).
90
Replacement type reactions that removed Sr from high salt wastes containing
0..01M P04~ were first reported by Rhodes (1957). The presence of the phosphate
in the solution caused a reaction with the calcite in the soil to form apatite,
90
Ca5(P04)3OH, with the Sr substituted for calcium in the apatite crystal
lattice. Other cations that can be included in the apatite lattice during
the replacement reaction include promethium, uranyl ions and plutonium (Ames,
1960). The cations that are removed into the crystallizing apatite are also
the cations found to be enriched in natural apatites, including strontium,
uranium, the rare earths, thorium, manganese and barium (Deer, Howie and
Zussman, 1962). The effects of phosphate concentration, calcite replacement
90
rate and pH on removal of Sr from solution was given by Ames (1959). Tamura
90
(1962) showed the effects of contact time and phosphate concentration on Sr
removal by calcite replacement. Belot and Gailledreau (1963) studied the
90
kinetics of Sr retention during the calcite-phosphate replacement reaction.
One of the problems with the use of replacement type reactions for radio-
nuclide removal from waste solutions involves kinetics. As mentioned above,
the reaction is diffusion-controlled. As the thickness of solid product
becomes greater, the diffusion path becomes longer and more tortuous, leading
to increasingly slower reaction rates. If the solid product coating is too
thick, reaction ceases and the radionuclides pass through the soil without
removal by a replacement reaction. In addition, it is usually necessary to
add a substantial amount of reacting ions such as phosphate to the waste solu-
tion before ground disposal, which can be costly and lead to a precipitate in
the solution prior to disposal. Adjustment of solution pH also may be required,
as in the case of the calcite-phosphate reaction.
PHYSICAL TRANSPORT AND FILTRATION
Colloidal particles, due to their large size compared to solution species
and their changes in charge in response to pH, exhibit different migration and
2-17
-------
retention behavior from solution species (Van Olphen, 1963; Grebenshchikova
and Davydov, 1965; Sheidina and Kovarskaya, 1970). Colloidal radionuclide
particles that have been described in the literature are of two types. These
are true colloids and pseudocolloids. A true colloid is defined as composed
of a radionuclide only, while a pseudocolloid is a radionuclide adsorbed on
another colloidal size particle. Several rare earths, zirconium, plutonium,
curium, antimony and americium could occur partially as true colloids, preci-
pitates or particulate suspensions (Rhodes, 1957; Bensen, 1960; Andelman and
RozzelT, 1970; Fukai and Murray, 1974; Price and Ames, 1975; Ames, 1976). .The
downward movement of colloidal particles through soils would mainly depend
upon physical factors such as permeability, rate of water movement and
particle size. This same dispersion and translocation of soil colloids occurs
under natural soil conditions (Drew, 1967; Gerasimov and Glazovskaya, 1960) and
may lead to problems with internal soil drainage. Ground disposal of low ionic
strength condensate wastes was reported by Routson (1973) to result in partial
plugging of the soil column and reduced infiltration rates because of disper-.
sion and filtration of upper soil colloids. Generally, the downward movement
would be limited because the colloids are subjected to filtration by the soils
and sediments (Routson, 1973; Price and Ames, 1975; Ames, 1976). The lateral
or horizontal movement of colloidal particles depends upon their degree of
exposure to wind and surface water action (Hakonson et al., 1975).
A recent report concerned with the long term leaching of silicate glasses
containing high level waste (Scheffler et al., 1977) serves as an example of
physical transport of several actinides. The base glass composition used to
contain the radionuclides was Si02 = 50.5 wt%, Ti02 = 4.2 wt%, A1203 = 1.4 wt%,
B203 =13.6 wt%, CaO = 2.8 wt% and Na20 =27.5 wt%. Two actinide-containing
glasses were made using the base glass formula. One contained 78.43 wt% base
glass, 16.64 wt% fission product and corrosion product oxides, 4.90 wt% Am02
and 0.03 wt% Cn^. The other glass contained 80.85 wt% base glass, 17.15 wt%
fission product and corrosion product oxides, 0.005 wt% Am02 and 2.00 wt% Pu02-
Leaching experiments with distilled water, 1M NaCl and a saturated brine solu-
tion were performed. Actinide concentrations were on the order of 10" M in
the leachates after a period of initial leaching, with rates of 10" to 10" g/
2
cm /day common. An examination of the leachate to determine the form of the
actinides indicated that particles attached to the actinides actually dissolved
2-18
-------
from the glass rather than the actinides alone, and that the particles were
2
high silica polymers with an average size of 10 urn. The actinides (curium
americium, and plutonium) were bonded to the silica polymers in the glass
and remained so in the leachate.
SATURATION EFFECTS
In migration of radionuclides, hydraulic conductivity is of primary
importance. Hydraulic conduction is the ratio of the flux density (the volume
of water flowing through a cross-sectional area per unit time) to the hydraulic
gradient (the head drop per unit distance in the flow direction). The reason
for the importance of hydraulic conductivity to radionuclide migration is the
requirement that a liquid (water) move the radionuclides either physically as
dispersed colloids or on dispersed soil colloids, or as ions or neutral com-
plexes in solution. If the soil water does not move, then the radionuclides
in, or contacted by, the soil water do not move except by diffusion, a rela-
tively slow process (Hajek, 1965; Hajek, 1966; Rancon, 1973).
Coincidentally, the most important difference between saturated and
unsaturated flow is the same hydraulic conductivity (Hillel, 1971). When the
soil is saturated, nearly all pores are filled and conducting, and conductivity
is at a maximum. As the soil becomes unsaturated, some of the pores become
air-filled and the conductive cross-sectional area decreases. In addition,
the first pores to empty under tension are the largest and most conductive,
and tortuosity is increased by these empty pores. In unsorted soils and sedi-
ments, the large pores that resulted in high conductivity at saturation become
barriers to liquid flow between smaller pores during unsaturated flow. Hence,
the transition from saturated to unsaturated flow may result in a steep drop
in hydraulic conductivity of several orders of magnitude as the tension
increases from 0 to 1 bar. At higher tensions, conductivity may be so low
that steep negative pressure gradients are required for any appreciable soil
water flow to occur.
An interesting corollary of the pore size-conductivity relationship is
that at or near saturation, a sandy soil conducts water more rapidly than a
clay soil with many micropores. When the soils are unsaturated, however, many
of the micropores in the clay soil remain filled, and consequently, the hydraulic
2-19
-------
conductivity in the clay soil does not decrease nearly as sharply as it does
in sandy soil under the same tension. In fact, an occurrence of a sand layer
in a clay soil can constitute a barrier to water flow through the clay soil
under unsaturated flow conditions (Hillel, 1971; Winograd, 1974).
Most of the work on radionuclide adsorption by soils has been on nearly
saturated to saturated flow systems, mainly because saturation represents a
system parameter where hydraulic conductivity, and subsequently radionuclide
migration, is at a maximum value. Because most of the laboratory radionuclide
adsorption work on soil columns is conducted at saturated flow conditions,
and the soil-water system is largely unsaturated during waste disposal opera-
tions, Knoll (1960) examined the possibility that the cation exchange capacity
of the soil under unsaturated flow conditions was not equal to the soil capa-
city at saturated flow conditions. His interest was in applying laboratory.
results to field situations. Small columns of'a fairly uniform, very fine
sand at constant temperature, packed bed density and tension were used with
+2 +2 -2
an influent solution containing 1.0 g/1 Ca ,0.1 mg/1 Sr and 2 x 10
Qfi
uCi/ml Sr. The results of the experiments are given in Table 2-6. There
was apparently no difference in strontium capacity between saturated and
unsaturated flow conditions.
TABLE 2-6. SOIL CAPACITIES FOR SATURATED AND UNSATURATED
FLOW CONDITIONS (KNOLL, 1960)
% Saturation Average Flow Rate, ml/hr Strontium Capacity, meg/100 g
100 1.3 4.0 x 10"4
100 2.5 4.4 x l.O"4
100 10.3 5.2 x TO"4
100 11.0 5.2 x 10"4
100 10.2 4.6 x 10"4 '
65 2.1 3.6 x 10"4
65 1.42 6.0 x 10"4
58 0.22 5.2 x 10"4
48 0.3 5.4 x 10"4
36 0.2 5.4 x 10"4
2-20
-------
However, it is worth noting that column flow rate differences between
100% saturation and 36% saturation are less than a factor of 100. With use
of an aggregate of less uniform pore size, flow rate (hydraulic conductivity)
differences could have been much greater, the volume of soil contacted by the
aqueous phase under unsaturated flow conditions much smaller, and hence the
cation exchange capacity of soil might have differed considerably between
saturated and unsaturated flow conditions. On the other hand, the clay frac-
tion of the soil constitutes the largest exchange capacity and smallest pore
sizes. Since the smaller pores are involved in unsaturated flow, there may
be little obvious effect on the exchange capacity of the soil in desaturating
it. If the columns are normalized by referring to a column pore volume rather
than a simple effluent volume, the difference in ion exchange capacities of
saturated and unsaturated flow columns tends to disappear.
Nielsen and Biggar (1961, 1962a, 1962b) have shown that the drier the
soil, the greater the effluent volume required to reach a tracer effluent/
influent ratio (C/Co) of 1.0. The desaturation eliminates the larger flow
channels and increases the stagnant water volumes which are difficult to dis-
place. The relatively unmoving water volumes caused by desaturation are shown
as C/Co versus volume of effluent (ml)/volume of water in the column (ml).
The latter coordinate represented the effluent volume divided by the pore
volume used by the soil at that saturation. Desaturating the soil progres-
sively shifted the initial breakthrough to the left and increased skewness of
the breakthrough curves. Aqueous diffusion is a relatively unimportant migra-
tion mechanism near saturation, but can become important as solution velocity
through the soil decreases at unsaturated flow conditions for low Kd
radionuclides.
Nelson et al. (1962) conducted model field tests in disposal sites, one
24 in.2 and the other 6 In. . The influent was a solution of 600 ppm Ca*
QC
and Sr. The water table at the field site was 12 ft below the ground surface,
overlain by a fairly homogeneous fine sand and silt soil with 31% porosity.
Only a small part of the soil beneath the model disposal site was found to be
water saturated as measured by neutron moisture probe. A section through the
disposal site was monitored to show moisture distribution to the groundwater
table. Unsaturated flow conditions were typical of most of the section.
2-21
-------
Baetsle et al. (1965) examined the breakthrough curves of Co, added
to the influent as K^ Co (CN)g, through Mol white sand. Percolation velocity
was about 3 cm/hr at soil saturation of 100%, 76%, 58% and 46%. The data were
plotted as effluent volume in milliliters versus fraction of the influent Co
concentration in the effluent at a given throughput volume. The curves were
somewhat different in shape than those of Nielsen and Biggar (1961) because
percolation velocity was ten times faster through the Co column, and the
columns were only about one-fifteenth as large as those used by Nielsen and
Biggar. The differences resulted in a flattening of the terminal ends of the
Co column breakthrough curves, and a very rapid appearance of the Co tracer
in'the initial column effluents. Columns of yellow eolian sand were used to
85
obtain Sr breakthrough curves at 100%, 45% and 30% saturation. These columns
tended to load in successive steps, indicating that the solution in the non-
saturated columns percolated through preferred channels. The normalized
breakthrough curves also were given for the strontium work. The normalization
consisted of plotting C/Co versus the volume of effluent in cm divided by the
3
volume of water in the soil column in cm. The unsaturated pore volume was
equal to the fraction of soil saturation times the soil pore volume at 100%
134
saturation.' Normalized curves also were given for Cs loading on the yellow
eolian sand. The stepwise loading was less pronounced but dispersion became
increasingly important as the soil became less saturated.
Several hydrodynamic dispersion coefficients in the unsaturated zone were
determined and these value's are given in Table 2-7. The same method was used
to obtain dispersion coefficients for saturated flow. After the pore volume
for a given saturation was determined on the column, a thin layer of I con-
taminated sand was added and the elution curve of the radionuclide obtained at
the same saturation. The elution curves counted at time t were then plotted
P /Prt
on a cumulatively linear probability scale as S;c/Co versus cm of column
length. The curve on probability paper became nearly a straight line.
Between the abscissa values 0.159 and 0.841, the standard deviation, a, was
determined, and using the relationship a = /2Dt, the dispersion coefficient,
D, was calculated.. Strontium adsorption had a smaller dispersion coefficient
in the less saturated columns. The velocity of percolation through the soil
averaged 4 cm/day, or about 5 x 10" cm/sec. The authors concluded that
2-22
-------
molecular diffusion played a major role in radionuclide movement as soil
saturation became a very small value. There is considerable controversy in
the literature at present regarding the extension of the dispersion concept
to a field scale.
TABLE 2-7. DISPERSION COEFFICIENTS IN THE UNSATURATED ZONE OF
A YELLOW EOLIAN SOIL (BAETSLE ET AL., 1966)
Radionuclide
131r
60
Co(CN)
3.13x10
-4
Soil Saturation, %
100
-
75 55 45
Dispersion, cm^/sec
2.65xlO"4 l.lxlO"3 1.9xlO"3
2.20xlO"4 9xlO"4
35
2.5x10"3
30
-
6.10x10
-4
7.4x10
-4
9.5x10
3.3x10
-4
-6
1.2x10
-3
6x10
-6
Schwille et al. (1967) reported some model experiments on fluid flow in
the transition zone from unsaturated to saturated soils. They stress the
difficulties involved in analytically treating liquid flows, even in homogene-
ous media,-and suggest--that whenever possible field tests should be used to
investigate the migration of radioactive solutions.
Routson and Serne (1972), in a study involving experimental verification
of their one-dimensional PERCOL model, determined the breakthrough curves and
85Sr capacity for unsaturated Tank Farm sandy loam soil columns at 96%, 75%,
62% and 45% of saturation. The column volumes used were adjusted for pore
volume at that saturation. The parameters are given for each column includ-
ing influent solution composition, soluble salts in the initially dry soil
column, CaC03 content, temperature, bulk density, flow rate, soil volume, soil
weight and column dimensions. At 45% saturation, the flow rate was reduced to
0.055 cm/hr. There was no apparent reduction of capacity as a function of
saturation in these systems. Obtaining curves and capacities at lower satura-
tions would have required nearly 1 year of experimental time. At the higher
tensions and lower saturation, only a very limited amount of solution would
be transmitted per unit area. The fraction of the total radionuclides adsorbed
during ground disposal of liquid wastes in the very low saturation zones was
considered to be negligible due to the small volumes of solution per volume of
soil.
2-23
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Rancon (1969, 1973) performed several field studies of water behavior in
unsaturated soils. Simple methods were developed to measure in situ the move-
ments of the soil wetting front (deffned as 2% FLO by volume), the radionuclide
front and the .kinetics of soil drying. Rancon stated that soil contamination
did not move when there was no transfer of liquid water by infiltration or
redistribution, at least over the time period of the study (3 months). The
wetted front always constituted the maximum envelope of the contamination, and
all movement and redistribution of the radionculides took place within this
limit. Iodine and strontium did not move in the soil column during soil drying.
by evapotranspiration. However, if the water table is shallow, a steady flow
of liquid can occur from-a water source below to the evaporation sink above
(Hillel, 1971) as a function of the micrometeorological conditions (evaporation
rate) and water transmitting properties of the soil profile. Such capillary
soil moisture transmission could transport radionuclides to the surface, just
as it has been known to transport salts to the surface (salinization). Without
131
a carrier, part of the I used became fixed in the soil during the drying out
phase of the evapotranspiration studies and did not move during a following
elution. Tritium was'not retained in the soil under similar circumstances.
The protection afforded the groundwater by unsaturated soils was emphasized.
Three breakthrough curves which might be encountered with the use of soil
columns are shown in Figure 2-1 (Nielsen and Biggar, 1962b). In the first, no
interactions between solid, solvent and solute have occurred. Piston flow
rarely, if ever, occurs in soil columns because pores of constant radii would
be required to attain it. Ion exchange resin breakthrough curves approach
piston flow (Samuelson, 1963). The second curve showing longitudinal disper-
sion effects can be seen with soil columns under the proper operating conditions
(where the solid is selective for the tracer ion and the column flow rate is very
slow, reducing hydrodynamic dispersion to a minimum). This symmetrical break-
through curve (area A = area B) can be used to derive a Kd value where C/CQ =
0.5. The third type of breakthrough curve in Figure 2-1 shows that the curve
is not symmetrical (area |A 7s area B). This curve is typical of a soil composed
of equal-sized aggregates with intra- and inter-aggregate pore systems of
differing size, which gives an extremely wide range in pore velocity distribu-
tion that is bimodal at saturation. In addition, the breakthrough curve can
2-24
-------
1.0
y 0.5
PISTONaOW
LO
0.5
0
1.0.
y 0.5
u
LONGITUDINAL
DISPERSION
WIDE RANGING
VELOCITY
DISTRIBUTION
"01 2 3 4
PORE VOLUMES
Figure 2-1. Types of column breakthrough curves
(Nielsen and Biggar, 1962b)
be displaced to the right by any chemical or physical process that retains the
solute within the column. Incomplete mixing throughout the entire soil solu-
tion moves the breakthrough curve to the left of a piston flow curve. In
reality, the actual radionuclide breakthrough curve usually represents some
combination of the five types of curves discussed above. Hence, column break-
through data can seldom be used to accurately determine Kd values.
SPECIFIC RETENTION
Specific retention, as defined by Meinzer (1933),. is the ratio between
the volume of water that a soil or sediment will retain against the pull of
gravity and its packed volume from an initially saturated condition. It is
2-25
-------
usually given as a percentage and can be expressed as R = 100 (r/V), where R is
specific retention, r the volume of water retained against"the pull of gravity
and V the packed volume of the soil.or sediment.
Disposal on a specific retention basis of selected liquid waste solutions
was practiced at Hanford from 1944 to about 1962. As used at Hanford, specific
retention in practice represented the volume of liquid that could be discharged
to a disposal pit of known dimensions without leakage to the groundwater. The
specific retention was expressed as a percent of the soil column volume as mea-
sured by the cross-sectional area of the disposal pit bottom and the height of
the soil column between water table and pit bottom. Lateral spreading was
ignored to introduce a conservative element into the specific retention
calculation.
A centrifugal technique for estimating soil specific retention properties
was reported by Bierschenk (1959). Data were obtained indicating that the
natural moisture content of Hanford soils could be reproduced by draining a
1-cm thick soil sample from the same horizon at 1000 gravities for 1 hour. The
average ratio of the natural field moisture content to. that of the centrifuged
samples ranged from 0.82 to 1.06. The~speciTic""Tetention capacity of Hanford
soils varied from less than 1 to 2 vol%. Usually the only reason for resorting
to specific retention was because normal ground disposal methods were inadequate
(Haney, 1957).
The problems associated with specific retention as a radionuclide waste
disposal technique include the fact that the forces holding the radionuclide
waste solution are relatively weak and that the radionuclides which are dis-
posed by this technique ofte.n remain in the retained solution. Any water added
at the top of the soil column, from natural precipitation or irrigation, would
lead to the migration of the retained waste solution into the groundwater.
Hence, specific retention as a waste disposal technique would be limited to
arid areas where there is little downward migration of precipitation to the
groundwater and no chance of other water additions to the top of the soil
column. In addition, large soil volumes are required, for specific retention
disposal because very little of the soil column ion exchange capacity is
utilized.
2-26
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THE DISTRIBUTION COEFFICIENT OR Kd
One of the most useful concepts for dealing with radionuclide migration
and retention processes on geologic materials is that of the distribution coef-
ficient or Kd. The application of the Kd concept to ion exchange reactions was
given above. However, the distribution coefficient can be directly derived
from the same thermodynamic principles underlying ion exchange and other physi-
cochenvical reactions. The more general derivation allows the application of
the Kd concept to a broad selection of radionuclide reactions with geologic
media.
The distribution of a dilute solute (trace ion) between two immiscible
solvents (aqueous solution and solid medium), can be derived from y and y1,
the chemical potentials of the trace ion in- the solution and solid medium,
respectively. At equilibrium, y = y1. If both are ideal dilute solutions,
then y = y* + RT£nx and u1 = u1* + RHnx', and by definition, u* + RT2,nx =
y1* + RTfcnx1, where u* and y1* are standard free energies and x and x' are the
concentration of dilute solute in the solution and solid medium, respectively
(Denbigh, 1971). Rearranging, RTJln ^-= - (u1* - u*). Then ^- = Kd because •
A A
both y'* and y* are independent of composition. The distribution coefficient,
(Kd), also is independent of the concentration of trace ion in the solution
»
and on the soil. The quantity y1* - y* is the standard free energy change,
AG*, for the reaction:
Trace ion in solution -»• Trace ion on solid medium (soil, rock),
and RTAnK = -AG*,
which is the usual relationship for a chemical reaction. If the solute (trace
ion) is very dilute, then the equilibrium mole fractions are proportional to
the concentration, and Kd equals the concentration of trace ion on the solid
medium divided by concentration of trace ion in solution (times a correction
factor to yield a result in solution volume per weight solid medium). The Kd
was originally proposed as a distribution law by Nernst in 1891 (Denbigh, 1971).
A consideration of mechanisms is not involved in the derivation of the Kd
concept in the above generalized manner. Removal onto the solid medium can be
due to a number of mechanisms. However, it should be remembered that the trace
ion was considered to be in an ideal dilute solution both in the solution and
2-27
-------
on the soil. The trace ion could conceivably be considered in an ideal dilute
solution if it is exchanged onto the soil and occupies only a very small frac-
tion of the soil ion exchange capacity. The requirements'that the Kd value
remain constant over a concentration range and that the soil-trace ion behave
ideally are not compatible with precipitation in solution as a trace ion removal
mechanism. With precipitation as a removal mechanism during Kd determinations,
the soil becomes an inert system constituent except insomuch as it may affect
the system pH and redox environment.
The soil can be considered a passive and incidental part of the.removal
process unless the trace ion can be shown experimentally to be adsorbed on the
solid medium. This can be accomplished by "blank" experiments in which the
solution containing the tracer and no soil or rock solids is shaken along with
the same solution containing the soil or rock solids. If there is evidence for
trace radionuclide removal without the presence of soil or rock solids in the
system, the Kd concept is not applicable. • ,
Other practical requirements of a Kd experiment include attainment of
.equilibrium between the solid medium adsorbed trace ion-and the trace ion con-
tained in the solution. Prior to contact of the solid medium and trace ion,
the solid medium should be in equilibrium with the none-trace solution macroions.
This procedure should be followed to ensure that only the trace ion is adsorbed,
and that only in trace concentrations. Trace concentration is shown experi-
mentally by demonstrating that Kd values remain constant over successive trace
ion dilutions. This procedure also demonstrates that equilibrium was attained
using several solution-solid contact times. The Kd value should refer to a
single ionic species, not a mixture of fission products. The basic concept
does not allow more than one ionic species to have a constant distribution
between two immiscible media.
r
The Kd is an experimentally derived value that is associated with the cir-
cumstances of its determination. It cannot be extended to another chemical sys-
tem with the same solid exchange medium. Likewise, the Kd from one exchange
medium in a given chemical system cannot be used for another solid exchange
medium in the same chemical system.
An obstacle to Kd comparisons is the common failure to adequately charac-
terize the solid exchange medium. Even with comparable solution compositions
2-28
-------
and derived Kd values, the solid exchange materials can be completely different
mineralogically, physically and chemically. A bulk chemical analysis, at least
a semiquantitative mineralogical and particle size analysis should be performed.
Very little of the literature Kd data include such characterizations of the
solid phase. With rocks, the characterization work also is very important.
Adsorption of radionuclides on rock fissures often is higher than on fresh
surfaces because the fissure surfaces have been altered by earlier percolating
solutions to minerals such as zeolites or clays that are radionuclide selective
and do not occur on fresh rock surfaces. Rancon (1972) suggested calculating
Kd based upon surface area rather than the mass of the solid in order to study
radionuclide adsorption by impermeable materials such as rocks. This Kd is
defined as:
i _ surface concentration ml solution
residual concentration in solution Cm2 Qf S(jrface area
Using a rock with a known surface area allows better comparisons than a known
weight because ion exchange is a surface area function for impermeable rocks.
However, once again, good characterization of the rock is required to allow
comparisons between rock types, or even within rock types.
The radionuclides which can be expected to react with rocks and soils in
other ways than ion exchange include americium, cerium, curium, europium, pro-
methium, plutonium, tritium and zirconium, and sometimes antimony, uranium and
ruthenium. Special care should be taken with the above radionuclides to assure
that the Kd values obtained are from trace concentrations and that the trace
ion is adsorbed on the solid medium. Where the oxidation potential (Eh) is
important, such as in plutonium systems, means will have to be found for Kd
value determinations in controlled Eh-pH environments.
Many techniques are used or under study for the determination of Kd values
including batch tests, low pressure column loading, axial- filtration and chan-
nel chromatography. The last two techniques are in a development and proving
stage, and may or may not be comparable to Kd values determined by a batch
technique. The standard Kd determination technique to which all other tech-
nique results are compared is the batch technique. It is theoretically possible
to derive column Kd values comparable to batch Kd values. However, difficulties
2-29
-------
are usually encountered in practice because the soil or rock unique hydrologic
properties are superimposed on the radionuclide adsorption effects in the result-
ing breakthrough curve (Baetsle et al., 1966). Consequently, the breakthrough
curve is not symmetrical about the 50% breakthrough point and the Kd value as
taken at the 50% point is in error. A batch Kd value is determined by agitat-
ing a weighed soil or rock sample with a known volume of chemically-characterized
solution until the radionuclide attains an equilibrium between solution and soil
or rock. To obtain reproducible and comparable radionuclide distribution coef-
ficients, the following.experimental requirements should be followed:
1. mineralogically and physically characterize representative rock or soil
and chemically characterize the solution before use in Kd measurements,
2. all rocks should be compared on a surface area rather than a weight basis,
3. dry a weighed sample of soil or rock at 105°C to obtain weight corrections
for moisture content,
4. pre-equilibrate the solids with nonradioactive waste constituents and
discard solution,
5. add the radionuclide at several concentration levels and equilibrate versus
time until no significant changes are observed (precipitation of solid
phases should be absent),
6. use triplicate samples for each equilibration,
7. a blank solution (containing no soil or rock) should be run along with
the samples at each radionuclide concentration level,
8. the equilibrium solution composition should be determined including major
cations, anions and organics,
9. the pH should be measured before and after equilibrium, and
10. a controlled redox potential is necessary where it has an effect on the
radionuclide oxidation state to assure that only one species of radionu-
clide is present in the equilibrating solution.
Kd Relationship to Migration
The importance of an accurately determined Kd value cannot be overempha-
sized when it is used in calculating the velocity of radionuclide movement in
geologic materials. If K is equal to the groundwater velocity divided by the
\(A rt
radionuclide migration velocity, then K = 1 + —*-, where:
2-30
-------
Kd = the distribution ccefficient = (concentration of radionuclide on the
solid/weight of solid)/(concentration in the solution/solution volume
in the ml),
p = bulk density of medium = weight of solid/volume of solid,, and
e = void fraction = groundwater volume/volume of solid (Burkholder and
Cloninger, 1977; Hiester and Vermeulen, 1953; Jackson et a!., in press).
It may be seen that as the Kd value increases, the radionuclide migra-
tion velocity must decrease a proportional amount, if the bulk density
of the medium, the void fraction and groundwater velocity are assumed
to be constant.
A doubling of the Kd value halves the radionuclide migration velocity in rela-
tion to the groundwater velocity. Lester et al. (1975) with their computer
model of radionuclide migration have shown schematically the effect of Kd value
on radionuclide discharge rate at the outlet of a soil column for a three-
membered decay chain. Impulse release was assumed and decay rates of all three
radionuclides were neglected. Differences in Kd'values for chain members sig-
nificantly reduced radionuclide discharge rates in comparison to the case
when the Kd values for all three radionuclides were the same, including zero.
Field Determination of Kd Values
Using the above equation, a groundwater velocity in the field can be
determined between two wells, or between a disposal site and a river bank
spring with the use of tritium for a tracer. Another radionuclide can be
added to the groundwater and the travel time over the same distance determined
for it. If the bulk density and void fraction of the soil can be determined
as well, a field Kd can be derived. Eliason (1967) has used the above method
to determine the travel times of groundwater and I between a ground disposal
site and the Columbia River near Hanford N-Reactor. The travel time for I
was 1.28 times that of the waste water (79 days travel time for H, 101 days
travel time for 131I). An 131I Kd value of 0.06 ml/g was estimated by Hajek
(1968) from these data. An advantage of this field method of Kd measurement
is that a resultant Kd value is obtained for soils and rocks which are seldom
homogeneous. The laboratory Kd is useful for field applications only if the
field soil or rock are assumed to be homogeneous and represented by the
2-31
-------
laboratory soil sample. However, the laboratory Kd measurement is relatively
rapidly accomplished and is inexpensive. The field Kd measurement has the dis-
advantage of being time-consuming and expensive. Another disadvantage of the
field. Kd measurements is the difficulty of accurately determining a composite
void fraction and bulk density for the field soil, sediment or rock.
An additional method of determining a Kd value for a waste disposal site
in the field involves a knowledge of the volume of waste disposal to the ground,
the radionuclide concentration per unit volume, the volume and bulk density of
the soil and the average radionuclide concentration per unit weight of soil.
Essentially this field method is a much larger version of a laboratory Kd
determination. A major problem with this type of field Kd determination is
that the saturated flow directly beneath the disposal site soon grades into
unsaturated flow., with much lower volumes of waste transmitted through the
unsaturated flow area. Usually an arbitrary decision is required to select a
wetted zone for the average soil radionuclide loading value. An exception
would be a radionuclide that is adsorbed very close to the disposal site. Thus,
the radionuclides with high Kd values present the least problems with this type
of field Kd determination.
An example of the latter type of field Kd is given by Schmalz (1972) who
obtained auger samples down to 53 cm in TRA disposal pond bottom sediments at
INEL, Idaho. These samples were analyzed for Cs, Sr, Co and Ce in
three intervals of 0 to 15.3 cm, 15.3 to 30.5 cm, and 30.5 to 45.9 cm. The
analytical results, were used to estimate the quantities of each of the above
radionuclides in the three measured zones and these gradations extrapolated to
15.3 m to.obtain a total radionuclide concentration in that volume of sediments.
Knowing the volume of sediment and its estimated average bulk density, the radio-
nuclide contents of the sediments, and the volume and radionuclide content of
the waste disposed to the TRA pond, Kd values were estimated. The laboratory
Kd value for Cs was 285 ml/g, for example, while the field Kd value was
600 ml/g. Considering the assumptions as to homogeneity of sediments underly-
137
ing the TRA pond and their Cs content, perhaps the field and laboratory Kd
values are not as far apart as the numbers would seem to indicate.
One of the problems associated with the second type of field Kd determina-
tion is that usually the influent to the. waste disposal site is not constant in
2-32
-------
either composition or volume, and disposal records are often too generalized to
determine waste solution compositional and volumetric variations. Redox poten-
tials and final pH values are seldom known and are difficult to measure in field
systems. The experimental requirements for a reproducible laboratory Kd value
can seldom be followed for a field Kd determination.
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Mattson, S. 1931. The Laws of Soil Colloidal Behavior. IV. Amphoteric
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2-36
-------
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2-37
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2-39
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SECTION 3
ELEMENT CHEMISTRY AND GEOCHEMISTRY
The chemistry and geochemistry of the 19 elements chosen for review are
examined in the following section from the standpoint of radionuclide
migration and retention by geologic materials. Some of the important radio-
isotopes normally found as fission products or neutron activation products
are enumerated. Adequate chemical data were available only for the elements
with naturally-occurring stable isotopes. The radiochemistry of all of the
isotopes of an element is nearly identical in respect to their reactions
with geologic materials, with possible minor exceptions in H and H. A
summary of what is presently known (1977) relative to the radionuclide-
geologic media interactions for the selected elements is given. It will be
noted that the data presented for cesium and strontium adsorption on soils
and rocks are inordinately large. This circumstance is due more to historical
accident than to an unusually active experimental interest in cesium and
strontium adsorption. In the past, many of the low to intermediate level
wastes were disposed to ground. Two of the most hazardous radionuclides
contained in these wastes were cesium and strontium, and hence the early
interest in examining their migration and retention in soils.
AMERICIUM
Natural Soil and Rock Distributions
The element americium has not been reported to occur naturally in soils
or rocks.
Brief Chemistry
Americium has not been reported in nature bat is formed in high alpha
238 241 8- 241
densities from uranium by the reaction U(a,n) -»• Pu * Am, or from
3-1
-------
high burnup piutorn urn having a high 241Pu content by the reaction 239Pu
(n,Y) - 240Pu(n,Y) -241Pu 5- 241Am (Keller, 1971).
The americium radionuclide data of importance to radioactive wastes are
given in Table 3-1. It is found in the following four oxidation states:
Am(III), Am(IV), Am(V) and Am(VI). The most stable state of americium is the
trivalent. The higher oxidation states are strong oxidizing agents and are
therefore stable only in media that contain no oxidizable substances (Coleman,
1963). The ionic radius of Am+3 is 1.07 A (Ahrens, 1952).
TABLE 3-1. AMERICIUM RADIONUCLIDE DATA
(KELLER, 1971)
Isotope
241Am
242mAm
242Am -
243Am
Ha If -Life
458
152
16
7370
years
years
hours
years
Decay Moi
a
IT, a
0-, EC
a
The stability constants and-complex formation of Am(III) are probably
more well known and studied than any of the other trivalent actinides. The
tendency to complex ion formation is a function of factors such as ionic size
and charge. In order of decreasing complexing tendency, the actinide series
is generally M >M02 >M >M02 . For univalent anions, complexing ability
with Am(III) follows the order for F~>N03~>C1~>C104~, and for bivalent anions,
CO,~2>oxalate~2>SO,,~2 (Keller, 1971).
0 *
Solid Phase and Solution Equilibria
Information regarding the possible solid phases of americium in soil and
rock environments is lacking. However, the relative solubilities of several
americium solids based upon estimated thermodynamic data (Latimer, 1952;
Keller, 1971) are given in Figure 3-1 as a function of pH. These are oxides
and hydrous oxides of americium contained in an oxidizing environment (pO« -
0.68 atm). The solubility of the solids decreases rapidly in the direction
of increasing pH. As the environment becomes less oxidizing, the activity
3-2
-------
of Am associated with each of the solid phases [except Am(OH)3] increases
rapidly. At PC^g) = 3-58 atm' for example, the curve for Am activity
from Am(OH)4(s) increases by two log units.
-4
e
<
81
-10
-12
-14
"""AnHOHL
AmO-lOH..
V-. (pO, 8.68)
4567
9 10
Figure 3-1. The relative stability of various americium
solids in an oxidizing soil environment
[p02(gj = 0.68 atm]
The activity of various americium species in solution in equilibrium with
AmCL(s) in an oxidizing environment (pO« = 0.68 atm) is given in Figure 3-2.
• 2+ ?+
The thermodynamic data for all the species except AmOH and AmHgPO^" were
selected from Si 11 en and Martell (1964). The AmOH2+ and AmHgPO^2* data were
selected from Keller (1971). In the normal soil pH range, the complex
2+ +
species Am(OH) and AmSO,, would control the total activity of americium,
' 3+
with Am of somewhat lesser importance. In the environmental conditions
chosen, chloride, nitrate and phosphorus species were relatively unimportant
in terms of the total americium present in solution. However it should be
2+
kept in mind that in a relatively high chloride environment, the AmCl , for
example, could become the principal americium species in solution. With an
3-3
-------
increase in .pCL (more reducing environment) and the concentration of anions:
the curves given in Figure 3-2 would move upwards. However, the relative
position of most of the isotherms would not change. The americium species
all of cationic in the normal soil pH range of from 4 to 8. It is to be
expected, therefore, that ion exchange could be an important americium
removal mechanism.
£ -12
Figure 3-2. The activity of various americium species in equilibrium
with Am02(s) in an oxidizing soil environment Cp02(g) -
- = " = =
0.68 atm
Experimental Adsorption Results
pCT = 3.0 and
= 5.0.
The adsorption of americium by soils has been studied by several investv
gators. Hajek and Knoll (1966) and Knoll (1965, 1969) described several
americium adsorption experiments on Hanford wastes with emphasis on the
effect of organics. An acidic, high salt (5.4M N03) waste spiked with
organics typically found in waste streams (lard oil, carbon tetrachloride,
DBBP and TBP) that contained americium showed very rapid breakthrough when
>,
percolated through a Hanford sand.^i The americium Kd was less than 1 ml/g at
a soil-solution pH as high as 3. Plutonium in these solutions had a Kd of
3-4
-------
2 to 3 ml/g. If the wastes were neutralized and the resultant supernatant was
mixed with 20% by volume organics, the Kd increased to 41 ml/g. A water leach
of the neutralized sludge mobilized some americium which, upon contact with
soil, yielded an americium Kd of 500 ml/g. In all cases, the americium Kd was
less than the plutonium Kd. In other experiments Knoll loaded columns of soil
with tap water spiked with americium. Various organic solutions were then per-
colated through the columns and the quantity of americium leached recorded.
One hundred and thirty column volumes'of TBP(20%)-CC14 (80%) leached 10% of
the adsorbed americium versus 5% of the plutonium. DBBP (30%)-CC14(70%) after
160 column volumes .lea.ched 15% of the americium and 40% of the plutonium.
D2EHPA-TBP in normal paraffin hydrocarbons after 30 column volumes leached 80%
of the americium and 30% of the plutonium. Hydroxyacetic acid at 0.25M released
100% of the americium and 55% of the plutonium after 10 column volumes. At a
lower concentration, 0.025M, hydroxyacetic acid removed 50% of the americium
and 40% of the plutonium after 70 column volumes. When the americium and plu-
tonium were added to the. organic solution and then percolated through soil
columns, the Kd values listed in Table 3-2 were observed, assuming a column bulk
density of 1.5 g/ml. There were significant differences between the two cases;
americium contacting soil then being leached with organic solutions and americium
added directly to organic solutions and contacting soils.
TABLE 3-2. AMERICIUM Kd VALUES FROM SEVERAL ORGANIC SOLUTIONS
TBP = TRIBUTYL PHOSPHATE, DBBP = DIBUTYLBUTYLPHOS-
PHORATE, D2EHPA = Di-(2ETHYLHEXYL) PHOSPHORIC ACID
(KNOLL, 1965)
Americium Kd,
Organic Solution ml/q
TBP(20%)-CC14(80%) 1.6
DBBP(30%)-CC14(70%) 0.6
D2EHPA-TBP in Hydrocarbons 0
Routson et al. (1975) used batch adsorption experiments (10 g soil to
25 ml solution) and 24-hour equilibrations to determine the Kd for americium with
desert sand and an Eastern sandy clay. The Kd was determined as a function
of solution calcium and sodium concentration at an initial pH of 2.5 to 3.1,
3-5
-------
and for the sandy clay, decreased as the calcium concentration increased.
At 0.002M Ca the americium Kd was 67 ml/g; at 0.20M Ca the Kd was 1 ml/c
For the Burbank sand the Kd was greater than 1200 for all concentrations of
calcium. The americium Kd for the sandy clay also was a function of the
sodium concentration in solution; 280 at 0.015M Na and 1.6 at 3.00M Na.
Although a final pH was not given, the Burbank sand is known to be alkaline,
suggesting that americium adsorption may be sensitive to pH.
Van Dalen et al. (1975) determined the americium Kd for Dutch subsoils '
from 90% saturated NaCl solutions at pH 7 to 8. Samples consisting of mainly
4
illite and kaolinite had a Kd of J5_x 10 ; for a river sand the Kd was 400.
Gypsum bearing and clay bearing sandstones had intermediate Kd values. The
pH dependence on sorption processes for americium between pH 5-8 was minor.
Fried et al. (1974a, 1974b) determined the adsorption of americium onto
a basalt. Addition of salts to the solution lowered the adsorption of
+3 +4
americium. La and Zr cations lowered the americium adsorption much more
+? +
than Sr ~ and Na suggesting ion exchange as the mechanism controlling
adsorption of americium. '
In Fried et al. (1977) the distribution of americium at a 32-year-old
disposal site in tuff, a very porous, soft volcanic rock, was investigated
by coring the tuff at the site to 14 ft. Counting of wafers at different
depths down the tuff core showed the americium to be concentrated mainly at
a peak at about 9 ft from the disposal site bottom. A core of the same tuff,
but uncontaminated, was used to duplicate the field results without much suc-
cess. A near-surface americium peak was obtained rather than the peak at
9 ft found in the field.
Fried (in press) also studied the movement of Am through Los Alamos
tuff by using water to elute americium already adsorbed on the tuff. The rela-
tive migration ratio for americium was 100 urn/meter of water flow.- Upper limit
for the relative migration ratio was 500 urn/meter of water flow. Over 95%
of the americium was found very close to the point of deposition. Surface
adsorption coefficients were derived for americium on National Reactor Testing
2
Station (now INEL) basalt in which the surface area in cm was used rather
than the basalt weight in the Kd calculation. In these experiments, rock
discs of known surface area were immersed in a solution of 1 x 10" M
3-6
-------
Small aliquots of the solution were removed and counted at 12-hr intervals
until a constant count was reached. The Kd value for americium was 24_._4_±
12 ml/cm for INEL basalt. The values for Kd decreased rapidly as Na+, Ca ,
Sr , La and Zr 4 ions were added to the americium equilibrating solution
in increasing quantities.
Sheppard et al. (1976) reported distribution between 12 characterized
241
soils from various locations and a solution containing Am(III). The loca-
tions included Muscatine, Illinois; Rich!and, Washington; Sarnwell, South
Carolina.; and Idaho Falls, Idaho. Equilibrations were for up to 4 months
with sampling of the solution at intervals. The distribution results are
given in Table 3-3. The distribution values ranged from 45,500 ml/g, for an
Idaho soil to 125 ml/g for a Washington soil. The distribution values were
not constant with time, but steadily increased 1% to 86%/month. With Burbank
soil, there was a 13% decrease/month. No change was detected for another
Washington soil (Hanford). It was postulated that such adsorption behavior
was consistent with the Am(III) being in the form of charged colloids or
hydrolysis products including Am , Am(OH)+", Am(OH)2+ and Amn(CH)op and
removal' by cation exchange and adsorption. It was reported that the removal
of soil organics and raising the system temperature did not appreciably affect
the rate of approach to equilibrium or the distribution values. Analysis of
the soil cation exchange capacity data indicated a high correlation (r =
0.96) between soil cation exchange capacity and clay content. Correlation
does not appear as marked between soil clay content and americium distribution.
Glover et al. (1977) also studied americium sorption by soils. Seventeen
characterized soils from around the United States were equilibrated for 48 hr
-10 -8
with distilled water solutions containing 10" M Am and 10" M Am. Eight
replicates were run per soiVat each americium concentration. Results were
expressed as Kd values. Kd values ranged from 82 to 10,000 on the different
1Q 1?
soils. Generally, the Kd values for 10 M and 10 M Am were the same, or
nearly so, with some exceptions. Regression analyses were performed on the
americium sorption data to determine whether or not significant relationships
existed between americium sorption and soil physical or chemical characteris-
tics. Initial results indicated that a direct relationship existed between
soil cation exchange capacity, clay content and americium Kd value.
3-7
-------
TABLE 3-3. AMERICIUM 50-OAY DISTRIBUTION. (COMPLETED FROM
THE DATA OF SHEPPARD ET AL., 1976)
Am Distribution, Monthly Change
Soil Identity ml XT in Distribution Ratio
Muscatine 4,830
silt loam
Burbank • 714
loamy sand
Ritzville 971 +1"
silt loam
Fuquay sand, 476 +32:;
0-5 cm depth
Fuquay sand, 417 4.32?;
5-15 cm depth
Fuquay .sand, '249 ' • +27"
15-50 cm depth
Hanford A 125 0"
Hanford B 833 4-34:.
Idaho A 3,920 +30?:.
Idaho B 43,500 +36r,
Idaho C 37,000 +81S
Idaho 0 10,900 +53?;
Polzer and Miner (1977) relate americium adsorption data on the same
soils to differences in the effective positive charge of americium (Korotkin,
1972) as influenced by system pH. Kd values for americium with several soils
from around the United States are given in basic and acidic systems.
Migration Results
Field Studies-
Fowler and Essington (1974) detected a possible difference in solubility
between americium and plutontum fallout in soils. Americium may be more solu-
ble than plutonium and may become the radionuclide of prime concern because
3-8
-------
it has a faster migration rate in soils. Several instances of decreasing .
Pu/Am ratios with depth in the soil profile could be explained by differential
Plutonium and americium solubilities.
Laboratory Studies--
Hajek (1966) used contaminated soil samples from the Hanford 216-Z-9
covered trench upper 60 cm to determine the leach rate of americium. Small
columns of 5 and 7.5 g of soil were leached with groundwater (composition
o
not given) and IN NaN03 at leach rates of less than 4 ml/cm /day. Effluents
were collected at intervals and plotted as percent americium eluted versus
column volumes of leaching solution throughput. About 7.5% of the americium
was leached with 8 column volumes of groundwater, after which the leach
rate became zero. In the other column, about 6.5% of the americium was
eluted with 20 column volumes of.groundwater, and americium was still being
slowly leached. With the IN NaNO-, leaching, the americium leaching curve
level off at 33% americium removed from the column. The author considered
this to indicate that at least part of the americium was exchangeable.
Cline (1968) demonstrated that soil pH was directly related to amenciurn
migration. Americium nitrate was applied to two soil columns and 254 cm of
irrigation apqlied as an influent. In the acid soil (pH 4.5) the americium
was retained in the top 5 cm with 98% in the top 1 cm. Americium penetrated
to 20 cm in the basic soil and only 76% remained in the top one centimeter
of the column. A larger fraction of the americium was Am at 4.5 than at
7.5.
Knoll (1969) studied the problem of disposal of organic wastes to the
ground and the affect of this disposal on the adsorption or stripping of
radiorrucTides from the soil column. The organic wastes studied included
dibutyl butylphosphonate (D88P), di-(2-ethylexyl) phosphoric acid (D2EHPA), a
mixture of Lard Oil and CC1, (Fab Oil), methylisobutylkatone (Hexone), a com-
mercial mixture of straight'chain hydrocarbons including decane, undecane,
dodecane, tridecane and tetradecane (NPS) and tributyl phosphate (TBP). The
effect of organics on soil ion exchange properties was studied for an alka-
line, Burbank sandy loam soil type by saturating the soil in a 10 cm column
with each of the above organics and loading with Cs traced water. The
3-9
-------
cation exchange properties of the soil were not affected by the prior con-
tact with organics, at least in terms of the Cs capacity. In other columns,
the americium was loaded on the column and then contacted with the organic
waste. The americium in the effluent was monitored. The results of leaching
with organic wastes is shown in Table 3-4. The soil adsorption of americium
carried by organics was investigated by passing the americiurn-contaminated
organics through the same soil. The 50% column breakthrough of americium
occurred at 9 column volumes of TBP-CCl^ (20 to 80%), at 3..5 column volumes
of DBBP-CCl^ (20 to 80%) and the soil .adsorbed none of the americium contained
in 0.4M D2EHPA-0.2M TBP in NPH. .
TABLE 3-4. LEACHING OF AMERICIUM ADSORBED Of! THE
SOIL BY ORGANIC WASTES (KNOLL, 1969)
Americium Column
Organic Haste Removal, % Volumes
D8BP-CC14 (30 to 70%) 15 80
Fab Oil 1 30
0.4M D2EHPA-0.2M TBP in NPH 80 50
0.25M Hydroxyacetic Acid, pH 3.7 100 5
0.125M Hydroxyacetic Acid, pH 3.8 55 10
0.025M Hydroxyacetic Acid, pH 3.8 50 , 70
Summary
Americium adsorption (Kd values) on various soils and rocks have been
correlated with 1) cation exchange capacity of soils (Sheppard, Kittrick and
Hart, 1976; Glover and Miner, 1977), and 2) the concentration of the competing
cations (Na*, Ca , Sr , La and Zr ) so that the adsorption of americium
decreases with an increase in concentration of competing cations (Fried,
Friedman and Quarter-man, 1974; Fried, Friedman and Weeber, 1974; Routson,
Jansen and Robinson, 1975; Sheppard, Kittrick and Hart, 1976; Fried, 1977).
Although none of the workers have tried to distinguish the contributions of
polymer adsorption, cation exchange, hydrolysis and precipitation to the Kd
value, the laboratory data indicate that americium removal from solution may
be predominantly via an ion exchange mechanism. This conclusion is supported
3-10
-------
by the predictions based upon thermodynamic data. The americium species
are all cationic in the normal soil pH range from pH 4 to pH 8 (Figure 3r2).
References
Ahrens, L. H. 1952. The Use of lonization Potentials., Part I. Ionic Radii
of the Elements. Geochim. et Cosniochim. Acta. 2:155.
Cline, J. F. 1968. Uptake of ~ Am and L^?u by Plants. BNWL-714,
pp. 8.24-8.25..
Coleman, J. C. 1963. The Kinetics of the Disproportionate of Americium (V).
Inorg. Chem. 2:53.
Fowler, E. B. and E. H. Essington. 1974. Soils Element Activities October.
1972 - September, 1973. IN: The Dynamics of Plutonium in Desert Environ-
ments. P. B. Dunnaway and M. G. White (eds). NVO-142, pp. 7-16.
Fried, S. (in press) The Retention of Plutonium and Americium by Rock. To
appear in Science.
Fried, S., A. M. Friedman, J. Mines, G. Schmitz, and M. Wheeler. 1977.
Distribution of Plutonium and Americium at an Old Los Alamos Waste Disposal
Site. IN: Proc. Sym. on Dynamics of Transuranics in Terrestrial and Aquatic
Environment. Gatlinburg, TN. October 1975.
Fried, S., A. M. Friedman, and L. A. Quarterman. 1974a. Annual Report on
Project AN0115A, Fiscal Year 1974. ANL-8115.
Fried", S., A. M. Friedman, and R. Weeber. 1974b. Studies of the Behavior
of Plutonium and Americium in the Lithosphere. ANL-8096. pp. 10-11.
Glover. P. A., F. J. Miner, and W. L. Polzer. 1977. Plutonium and Americium
Behavior in the Soil/Water Environment. I. Sorption of Plutonium and Ameri-
cium by Soils.. BNWL-2117, pp. 235-246.
Hajek, B. F. 1966. Plutonium and Americium Mobility in Soils. BNWL-CC-925.
Hajek, B. F. and K. C. Knoll. 1966. Disposal Characteristics of Plutonium
and Americium in a High Salt Acid Waste. BNWL-CC-649.
Keller, C. 1971. The Chemistry of the Transuranium Elements. Vol. 3.
Kernchemie in Einzeldarstellungen. Verlag Chemie/GmbH.
Knoll, K. C. 1965. Reaction of High Salt Aqueous Plus Orqanic Waste With
Soil. BNWL-CC-313.
Knoll, K. C. 1969. Reactions of Organic Wastes in Soil. BNWL-860.
Korotkin, Yu.S. 1972. Study of Transplutonium Element Hydrolysis. II.
Hydrolysis of Americium (III) in the Presence of Ions with Positive and
Negative Hydrat-ion Energy. JIWR-P-6403.
-------
Latimer, W. M. 1952. Oxidation Potentials. Second Edition. Prentice-Hall,
Englewood Cliffs, NJ.
Polzer, W. L. and F. J. Miner. 1977. Plutonium and Americium Behavior in
the Soil/Water Environment. II. The Effect of Selected Chemical and Physical
Characteristics of Aqueous Plutonium and Americium on Their Sorption by Soils.
BNWL-2117.
go
Routson, R. C., G. Jansen, and A. V. Robinson. 1975. Sorption of JTc,
237Np and 24lAm on Two Subsoils from Differing Weathering Intensity Areas.
BNWL-1889.
Sheppard, J. C., J. A. Kittrick, and T. L. Hart- 1976. Determination of
Distribution Ratios and Diffusion Coefficients of Neptunium, Americium, and
Curium in Soil-Aquatic Environments. RLOT2221-T-12-2.
Sillen, L. G. and A. E. Marten. 1964. Stability Constants of Metal-Ion
Complexes. Special Publication No. 17. The Chemical Society, London.
Van Dalen, A., F. DeWitte, and J. Wiskstra. 1975. Distribution Coefficients
for Some Radionuclides Between Saline Water and Clays, Sandstones and Other
Samples from the Dutch Subsoil. Reactor Centrum Nederland. 75-109.
ANTIMONY
Natural Soil and Rock Distributions
The average concentrations of antimony found in igneous and sedimentary
rocks are given in Table 3-5. Hawkes (1954) reported on the antimony content
of the soils of the Nyeba lead-zinc district of Nigeria. Soils that were
30 to 120 meters from ore bodies contained from 1 to 5 ppm Sb. Ward and
Lakin (1954) reported 2.3 to 9.5 ppm Sb in Idaho and North Carolina soils.
3-12
-------
Boyle (1965) reported that soils from the Galena Hill area, Yukon, Canada,
contained on the average 1.0 ppm Sb. Little is known about the behavior of
antimony during rock weathering. Antimony may accumulate in shales by adsorp-
tion on clay minerals and hydrous oxides such as Mn and Fe precipitates.
Antimony may be concentrated in soils relative to igneous rocks, also as an
adsorbate.
TABLE 3-5. ANTIMONY CONCENTRATION IN IGNEOUS AND SEDIMENTARY ROCKS
(TUREKIAN AND WEDEPOHL, 1961)
ppm
Igneous
Ultramafic
0.1
Basaltic
0.2'
Granitic
0.2
Syenite
0.1
Shale
1.5
Sedimentary
Sandstone
0.01
Carbonate
0.2
Brief Chemistry
1 pi
The stable isotopes of antimony found in nature are Sb (57.25%) and
123
Sb (42.75%). The radionuclides of interest in waste disposal are given
in Table 3-6.
TABLE 3-6. ANTIMONY RADIONUCLIDE DATA
(WEAST,
Isotope
124Sb
125Sb
126msb
126Sb
Half-Life
60.3 days
2.7 years
19 minutes
12.5 days
Decay Mode
3"
8-
IT, B-
6-
Antimony may exist in (-III), (0), (III) and (V) oxidation states, readily
losing its 2(s) and 3(p) electron's. Ionic radii are Sb"3 2.45 A, Sb+3 0.76 A
+5 °
a'nd Sb 0.62 A (Ahrens, 1952). The important antimony minerals include native
antimony (Sb), stibnite (Sb2S3), jamesonite (2PbS'Sb2S3), boulangerite
(5PbS'2Sb2S3), kermesite (Sb^O) and senarmontite (Sb203). Stibnite is
the most common antimony mineral.
Antimony, in the form of Sb , forms complexes such as Sb(QH)2+,
Sb(OH) ~ SbCl.OH", SbCU"2 and SbS "3 (Sillen and Martell, 1964). The exis-
H J +3 ° +5 J . J.-3
tence of simple Sb and Sb is unlikely. Further, Sb is amphoteric. Sb ,
3-13
-------
for example, reacts as Sb20, + 3H20 + 20H - 2Sb(OH)4" in bases, and in acids,
Sb00^ •*• H90 + 2H* -»• 2Sb(OH)?+. Sb*" dissolves in strong acids forming com-
C. •$ d. £•
plexes such as SbCl,~, and in strong bases forms complex antimonides such as
r o
Sb(OH)c" (Latimer, 1952). SbCU dissolves in a limited amount of water to
o • j
give a clear solution which, upon dilution, precipitates insoluble oxochlo-
rides such as SbOCl and SbALCl- (Cotton and Wilkinson, 1962). Izrael et al.
125
(1969) mentioned that Sb was amphoteric, with a solubility minimum at
pH 7.0.
Solid Phase and Solution Equilibria
Thermodynamic data for Sb203 (Sillen and Martell, 1964) and Sb^Og,
Sb(OH)3(s), Sb204(s), Sb205, Sb2S3, SbCl3, and SbF3 (Wagman et al., 1968) com-
pounds were selected and plotted in Figure 3-3 to determine the relative sta-
bility of the antimony solid phases. This figure represents a highly oxidiz-
0 68
ing environment (00 = 10 ). Under these conditions SbCl3, SbF3, and Sb^
are very soluble and therefore are not shown.
Figure 3-3. The activity of SbO in equilibrium with various antimony
solids in an oxidizing soil solution environment Cp02(g)
= 0.68 a tan]. _
-------
0 68
Under these oxidizing conditions (02 = 10 atm) the minerals in
increasing order of stability throughout the pH range can be arranged as
follows: orthorhombic Sb20v cubic Sb^O-j, orthorhombic Sb/L, cubic Sb^L,
Sb(OH)3, Sb204, and Sb^Oj.. Solids containing antimony in a +5 state
(Sb20- and Sb20j and the Sb^S., would change their relative positions with
changes in the oxidation-reduction conditions, whereas the other solids will
retain their relative positions. Sb2Qc and Sb2Ch will become less stable
with an increase in reducing conditions, while Sb2S., will become the most
stable solid phase in extremely reducing conditions. Sb^S- is very soluble
in an oxidizing environment because of the oxidation of sulfur.to sulfate,
which is the stable solution species (Garrels and Christ, 1965). Under
extremely reducing conditions and total sulfur levels of approximately 10" M
?_
where the dominant sulfur species is no longer the SO." , the most stable
mineral will be Sb2S3.
In oxidation-reduction conditions corresponding to oxygen pressures of
<10" " atmospheres, Sb(OH)3 would be the most stable solid.
The relative activity of various solution species in equilibrium with
Sb203(s) is given in Figure 3-4. If equilibrium with Sb(OH)3(s) is assumed,
the activity of all the species would be approximately 2.9 log units lower
with the lines still parallel to the lines given in Figure 3-4. Thermodynamic
data for Sb(OH)0 were obtained from Baes and Mesmer (1976) and all the other
data were selected from Wagman et al. (1968). In the environmental pH range
of interest (4 to 8), it can be seen that HSb02° and Sb(QH)3° are the species
that would govern the total activity of antimony present in solution. The
other species [SbOF°, Sb(OH)/, SbO+, Sb(OH) ~ or Sb09~ and NH.SbC,,0] do not
Cm *T £ ^ £
contribute significantly to the total antimony present in the solution and
hence can be ignored. Since the predominant solution species [HSb02°,
Sb(OH)3°] are neutral, ion exchange as an antimony adsorption mechanism
is not expected to be important.
Experimental Adsorption Results
Essington and Nork (1969) used two different models to predict radio-
activity migration through basalt and Rainier Mesa tuffs. Many Kd measure-
125
ments were made, and a Kd for Sb of 1.4 ml/g was used in the modeling work.
-------
Figure 3-4. The activity of various antimony species in equilibrium
with Sb203(s) (cubic) with pNH4+ = 3.0 and pF- = 4.5.
Oxidation-reduction changes do not affect the concentra-
tion of species reported in this figure.
125Sb was reported concentrated in Mediterranean bottom sediments by
Giacoletto and Triulzi (1970), and 124Sb was suspected by Osterberg et al.
(1966) to be present in precipitates collected from treated Columbia River
water. 124Sb is not a fission product, but an activation product resulting
from the former operation of once-through- cooling of the Hanford reactors.
Izrael and Rovinskii (1970) measured the chemical states of fission
IOC
products in crater water and reported the Sb was 85% anionic, 8.5%
cationic and 6.5% colloidal. Arnold and Grouse (1971) have determined the
radioactive contamination of copper recovered from ore fractured with
nuclear explosives. Recovery was simulated by leaching test shot rubble of
the radionuclide content given in Table 3-7. One-hundred gram rubble samples
were leached for 6 days with 200 ml of synthetic copper ore leach liquor
(2g Cu+2, Ig Fe+3, 3g Fe"2/liter as sulfates, with H2S04 to maintain desired
pH of 1.0, 2.0 and 3.6). The 125Sb leached from the rubble was about 12% at
pH 1.0, 4% at pH 2.0 and too low to permit radionanalysis at pH 3.6. The
gross composition of the rubble has not been identified.
3-16
-------
TABLE 3-7. RADIOCHEMICAL ANALYSIS OF TEST SHOT DEBRIS
SAMPLED ABOUT 20Q-FT ABOVE THE SHOT POINT
AT ABOUT 2 YEARS IN AGE (ARNOLD AND GROUSE,
1971)
Radlonuclide Concentration, dpm/g x 104
144Ce <0.03
9°Sr Q.2
106Ru 1.4
60Co Absent
Gross Gamma 1.0
125
Adsorption of radionuclides, including Sb, by low-grade copper ora
in Safford, Arizona, was also determined. The chemical analysis of this
ore is given in Table 3-8.
TABLE 3-8. COMPOSITION Of SAFFORD COPPER ORE
(ARNOLD AND CROUSE, 1971)
Concentration, Concentration,
Element wt% _ Element wt% _
Cu 1.27 Ti , 0.56
A1 7.31 K 4.03
Ca 0.006 Na 1.98
Mg 0.73 F 0.10
Fe 3.71 C03 0.28
Mn 0.002 IS 0.62
Ni 0.011 SiO, . 58.5
Five ml of copper ore leach liquor, of the same composition as listed
above, were'adjusted in pH with H2S04 and the 25Sb added. The solution
was contacted 16 hr with 1 g of minus 20 mesh Safford ore. At pH 1.5, the
distribution coefficient obtained was about 40 m'!/g, at pH 2.1, about 50 ml/g
and at pH 2.7, about 45 ml/g. Arnold and Crouse (1971) state, however, that
3-17
-------
with other adsorption tests on Safford ere, and with solutions containing
125
Sb prepared by leaching radioactive debris, antimony was not as strongly
adsorbed as in the above test. These results suggest that the chemical forms
of antimony in the two tests were not the same.
R. J. Serne of Battelle, Pacific Northwest Laboratories (?ML), 1973,
equilibrated several sediments with simulated neutralized high-level waste
solutions containing 3Sb. The sediments and solutions are described in
Tables 3-9 and 3-10 and equilibrium distribution values (Kd's) are given in
Table 3-11. Some Ca(OH)2 sludge developed in solutions III and IV during
makeup. This occurred because of the relatively high calcium nitrate content
IOC
and final pH (12.0). This sludge was removed before the Sb was added,
125
and the solutions were used for Kd measurements. Sb levels were measured
immediately before use in the Kd measurements so that sludge removal did not
affect Kd values.
TABLE 3-9. SOLUTIONS USED IN 125Sb EQUILIBRATIONS WITH GLACIOFLUVIATILE
SEDIMENTS FROM WELLS. SEDIMENTS WERE PRE-EQUILI3RATED WITH
THESE SOLUTIONS MINUS THE 125sb ADDED LATER AS A
MICROCONSTITUENT
Molarity
Constituent
Solution
1.0
0.002
0.65
0.025
4.0
0.002
0.65
0.025
1.0
3.5
0.65
0.025
4.0
3.5
0.65
O.C25
0.6
1.18
0.79
0.002
0.15
0.25
O.C1
SjSG3
ca(:;G3),
:m,:;o3 "
pH=12
NaNO.
Ca(N03)2
'
-------
TABLE 3-10.
SEDIMENTS USED IN OBTAINING EQUILIBRIUM
DISTRIBUTION COEFFICIENTS WITH 125$b AND
THE SOLUTIONS GIVEN IN TABLE 3-9. CAPAC-
ITIES WERE DETERMINED ON SEDIMENTS MINUS '
THE >2mm MATERIALS
wt%
Cation Exchange
CaCOs g/IOOg
Sediment >2mm Removed Capacity, mea/lGQq Sediment
1,
2,
3,
4,
5,
5,
7,
silt
gravelly sand
sand
silty sand
caliche
silty sand
gravel
TABLE 3-11.
Sediments
1
2
3
4
5
6
7
0.0 6.96 13.2
39.1 6.73 1.24
20.1 6.16 1.30
5.0 5.95 2.60
7.88 37.00
5.6 5.28 7.65 '
58.9 6.73 0.95
EQUILIBRIUM DISTRIBUTION COEFFICIENT VALUES
(ml/g) BETWEEN SEVEN GLACIOFLUVIATILE SEDIMENTS
AND THE SOLUTIONS OF TABLE 3-9 CONTAINING ^25sb.
EACH VALUE IS AN AVERAGE OF THREE EQUILIBRATIONS
Solution
I II III IV V
0.036 0.548 79.7 122.8 0.081
0.0 0.483 81.4 71.1 0.045
0.0 0.491 55.4 109.0 0.055
1.27 2.09 83.7 43.6 0.228
0.73 0.56 15.8 14.7 0.140
0.65 0.40 15.3 24.1 .0.140
1.28 1.19 17.6 24.0 0.180
The relatively large 125Sb Kd values obtained on the sediments with
solutions III and IV are generally up to three orders of magnitude higher
than the Kd values measured with the other solutions. The high calcium
solutions also are solutions III and IV, suggesting that precipitation of
calcite has scavenged or coprecipitated a portion of the antimony. The
solutions were all quite basic in pH(12), enabling them to absorb C02 from
the air and precipitate Ca^.
3-19
-------
• ' 125
Movement of Sb from trenches and seepage pits in the Melton and
Bethel Valleys, Oak Ridge, Tennessee, to White Oak Creek was investigated
by Ouguid (1975, 1976). Typical results are given in Table 3-12. Rela-
tively minor amounts of the Sb are found in the groundwater or seepage
125
water compared to the amounts located on the soil. Similar Sb values
_2
for burial ground 4 in Melton Valley showed a range of from 1.0 to 1.3 x 10
dpm/ml in the groundwater and seepage water. This also represents a rela-
tively minor antimony movement from the soil into the groundwater. Computed
antimony Kd values ranged from 10 to 10 ml/g.
TABLE 3-12. 125Sb ANALYSES OF WATER AND SOIL SAMPLES
FROM A SEEP NEAR TRENCH 7 (DUGUID, 1975)
Water Samples
Sample
Date Count Rate, dpm/ml
3-5-73 £2.0
3-19-73 £4.0
5-11-73 6.7
Soil Samples
Count Rate, dpm/q
0-3
3-6
6-9
9-12
12-15
15-18
18-21
£690
£240
£120
£41
£18
£13
<9.7
Several equilibrium adsorption experiments between a radioactive melt-
glass, a low-grade copper ore and a copper leaching solution were conducted
by Jackson (1976) under studies for the nuclear chemical mining of copper.
144
The radionuclides that could be measured in.the neltglass included Ce,
3-20
-------
ELI, b Zu and Eu. Tracers added to the system to obtain distribution
coefficients included 85Sr, 88Y, 103Ru, 124Sb, 134Cs, 141Ce and 152Eu.
The chemical compositions of the glass and chalcopyrite ore are given
in Table 3-13. The ore was mostly composed of quartz and sericite with
minor feldspar, chlorite, biotite, calcite, chalcopyrite, pyrite and
hematite.
TABLE 3-13. ' CHEMICAL COMPOSITIONS OF
MELT GLASS AND PRIMARY
COPPER ORE (JACKSON, 19716)
Constituent
Cu
Fe
S
Si0
CaO
MgO
A1203
co2
TiO,
Ba "
Sr
Glass, wtf;
0.33
0.05
0.004
Ore. vit%
-
1.5
-
72.3
4.3
4.3
1.9
0.47
16.1
_
0.7
3.2
2.2
66.9
1.5
2.5
0.9
0.5
15.5
0.4
The chemical composition of the leaching solution is listed in Table 3-14.
Distilled water was used instead of leaching solution in a single experiment.
Autoclaves of 300 ml and 3.785 liters capacity were charged with various
combinations of leaching solution or water, ore and glass, and then run for
lengths of time varying from 47 days to 256 days. Oxygen flow (4 cm3/sec) was
maintained through the system at a temperature of 353 K (90°C). Except for the
first experiment, glass and ore leaching was accomplished separately. Glass
3-21
-------
TABLE 3-14. COMPOSITION OF LEACHING SOLUTION INFLUENT
(JACKSON, 1976)
Species Concentration, mg/1 Species Concentration, mg/1
Cu+2 4000 Zn+2 50
Mg+2 2000 Ni+2 30
K+ 900 . Na+ 20
Al+3 800 Cc+2 20
Ca+2 500 ' S04"2 25,000
Mn*2 200 NOj 200
Fe*3 100 '" Cl" 20
pH adjusted to 1.7 with H2S04
124
dissolving rates are listed on a surface area basis. Sb distribution was
computed as a surface area dependent function (Kds) from the measured Kd
values for both the ore and the glass. Antimony was one of the most strongly
sorbed alements. The ore phases adsorbed antimony activity much more strongly
than leaching glass surfaces. In the experiment using water for leaching,
the pH of the solution soon fell as a result of reactions with the sulfides
124
in the ore. The ' 'Sb Kd increased in all cases with reaction time, suggest-
ing that antimony was being removed by the decomposition products forming in
the system, such as gypsum, anhydrite or jarosite. In one case, 53.2% of
124
the original Sb was found in the solid products. The jarosite. composition
was nominally KFe^SO-MOHL, but varied from this as a function of run tem-
perature and growth location. The enrichment of antimony in ion-rich solids
could be expected from the similar occurrence in soils and shales.
Migration Results
Field Studies—
/
Haney and Linderoth (1959)-and Haney (1967) studied the disposition of
radionuclides beneath several ground disposal waste faci'lities at Hanford.
125
For example, Sb was detected to a depth of 24.4 m in soil sampling wells
drilled beneath the 216-BY covered trenches. These trenches were in service
from December 1954 to December 1955 and received 3.4 x 10 liters of U-plant,
high salt, scavenged waste (total dissolved soTids = 350 g/1, chiefly NaNO.,;
p'H = 9.5) whose °Sb content averaged 2 yCi/ml. Sb was the major
3-22
-------
subsurface soil contaminant from the bottom of the covered trenches to
24.4 m. Soil contamination averaged 1.5 x 10" uCi/g to 18.3 m and fell off
rapidly to 1.5 x 10""" yCi/g at 24.4 m. The 216-BC disposal trenches received
4x10 liters (0.9 column volumes) of the same type of waste on a specific
retention basis, which relies on long-term storage of the waste in the pore
space of the soil column above the water table. Approximately 1740 Ci of
~ Sb were discharged to the BC-3 covered trench of the 8C facility. The
Sb soil contamination was 1.8 x 10~^-uCi/g at 6.1 m, 4 ;< 10 yCi/g at
9.1 m and 10" uCi/g at 12.2 m. For comparative purposes, the Cs and
90
Sr soil contamination maxima beneath the BY trenches fell by 5 orders of
T ^ C
magnitude in the same vertical distance as the L Sb maxima fell 2 orders of
magnitude. The antimony, present.in a larger concentration (2 uCi/ml) than
Cs or Sr (0.5 to 0.005 uCi/ml each), does not adsorb on the soil with
nearly the Kd value of cesium or strontium, and travels much further down
the soil column as a consequence.
The leakage of 435,000 liters of Hanford high salt, high pH (1M NaOH,
5M NaNO-J high level waste from the 241-T-106 underground waste storage tank
in-1973 (Anon., 1973) afforded an opportunity to study the migration of the
various radionuclides in this waste through the soil column underlying the
tank. A series of wells were drilled to obtain contaminated soil samples for
radioanalysis. The results showed that antimony, cobalt and ruthenium moved
the most rapidly through the soil column. All of these radionuclides tend
to form complex species which are neutral or negatively charged, or hydrolyzed
species that also are neutral to negatively charged. In either case, the
mobility of the antimony can be expected to be relatively high. The caustic
content of this waste type is high enough to sclubilize a portion of the
soil column (Shade, 1974).
Magno et al. (1970) investigated the fate of the radionuclides from the
Nuclear Fuel Services plant in western New York state. The plant lagoon sys-
tem waters and soils were sampled and analyzed. They estimated from these
I or
data that approximately 90% of the Sb in the plant effluent passed through
the lagoon system and into nearby surface streams. Ninety-eight percent of
IOC
the Sb was described as "dissolved" (in solution), exiting from the last
lagoon. The antimony is not, therefore, associated with suspended solids,
but is in solution, as Figure 3-4 postulates that it should be.
3-23
-------
Laboratory Studies--.
Laboratory studies on antimony migration through rocks or soils are lack-
125
ing. There are studies of Sb plant availability due to chelation (Hale and
121
Wallace, 1970; Wallace, 1969) and even the use of Sb as a groundwater tracer
(Jennings and Sch.roeder, 1968), but little of primary concern with antimony
migration.
124
.Saas and Grauby (1973) studied the transfer mechanisms of Sb found
in reactor cooling water and other wastes to river, irrigated soil and ground-
water. The river water contained organomineral pollution in the form of
124
industrial and municipal sewage wastes that interacted with the Sb to
form chelates, an exchangeable fraction and hydrosoluble components defined
as the totality of the organic and mineral components of a soil that are
104
soluble in water. The percentage of the above h-ydrosoluble Sb fraction
was plotted versus depth in cm for an alluvial, calcareous soil column
124 124
receiving Sb in river water. Ten to 30% of the Sb was in the hydro-
soluble form. The point of this work was demonstrating the interactions
124
between pollutants in the river water and Sb that results in increased
antimony migration. These pollution interactions with radionuclides should
be addressed in nuclear reactor siting studies.
Summary
The principal aspect of antimony chemistry that governs its adsorption
reactions with soils and rocks over the pH range 4 to 8 is the prevalence of
neutral and complexed species (Figure 3-4). Low Kd values are found in high
salt, high pH solutions as well (Serne, 1973), with relatively rapid migration
through soils (Haney and Linderoth, 1959; Haney, 1967; Anonymous, 1973) and
surface waters (Magno et al., 1970). There are indications that antimony can
be precipitated or coprecipitated in solids (Serne, 1973), even at low pH
values (Jackson, 1976). In low redox conditions (02 < 10" atmospheres),
antimony trihydroxide may precipitate and may be a stable compound that con-
trols antimony concentrations (Figure 3-3). Many of the antimony species
associated with organopollutants are water soluble (Saas and Grauby, 1973)
and migrate readily with the water. Thermodynamic data on the antimony com-
plexes formed with soil organic materials do not exist at present because the
complexes are chemically poorly defined, and in many cases, unknown.
3-24
-------
References
Ahrens, 1. H. 1952. The Use of lonization Potentials. Part I. Ionic Radii
. of the Elements. Geochim. et Cosmochim. Acta. 2:155.
Anonymous. 1973. 241-T-106 Tank Leak Investigation. ARH-2874.
Arnold, W. D. and D. J. Grouse. 1971. Radioactive Contamination of Copper
Recovered from Ore Fractured with Nuclear Explosions. ORNL-4677.
Baes, C. F. and R. E. Mesmer. 1976. The Hydrolysis of Cations. John Wiley
and Sons, New York.
Boyle, R. W.. 1965. Geology, Geochemistry, and Origin of the Lead-Zinc-Silver
Deposits of the Keno Hill-Galena Hill Area, Yukon Territory. Geol. Surv.
Canada Bulletin III.
Cotton, F. A. and G. Wilkinson. 1962. Advanced Inorganic Chemistry. Inter-
science Publishers.
Duguid, 0. 0. 1975. Status Report on Radioactivity Movement from Burial
Ground in Melton and Bethel Valleys. ORNL-5017.
Duguid, J. 0. 1976. Annual Progress Report of Burial Ground Studies at
Oak Ridge National Laboratory: Period Ending September 30, 1975. GRNL-5141.
Essington, E. H. and W. E. Nork. 1969. Radionuclide'Contamination Evalua-
tion-Milrow Event. NVO-1229-117.
Garrels, R. M. and C. L. Christ. 1965. Solutions, Minerals and Equilibria.
Harper and Row.. New York. pp. 213-216.
Giacoletto, C. and C.. Triulzi. 1970. Radiochemical Studies for Radioactivity .
Determinations on Marine Sediment Samples: Cs-137, Sb-125, and Co-60.
CISE-N-118.
Hale, V. 0. and A. Wallace. 1970. Effect of Chelates on Uptake of Some
Heavy Metal Radionuclides from Soil by Bush Beans. Soil Sci. 109:262-2.53.
Haney, W. A. and C. E. Linderoth. 1959. Exploratory Field Study of a Ground
Waste Disposal Facility. HW-60115.
Haney, W. A. 1967. Final Report on the Effects of Ben Franklin Dam. BNWL-41.2.
Hawkes, H. E. 1954. Geochemical Prospecting Investigations in the Nyeba
Lead-Zinc District, Nigeria. U.S.G.S. Bulletin 1000-8.
Izrael, U. A., V. N. Petrov, A. A. Pressman, F. A. Rovinsky, E. D. Stukin,
and A. A. Ter-Saakov. 1969. Radioactive Contamination of the Environment
by Underground Nuclear Explosions and Methods of Forecasting It.
Izrael, Yu.A. and F. Ya. Rovinskii. 1970. Hydrological Uses of Isotopes
Produced in Underground Nuclear Explosions for Peaceful Purposes. UCRL-
Trans-10458.
w ™ £ 3
-------
Jackson, D. D. 1976. Radiochemical Studies for. the Nuclear Chemical Mining
of Copper. UCRL-52025.
Jennings, A. R. and M. C. Schroeder. 1968. Evaluation of Selected Isotopes
as Ground Water Tracers. Watar Resour. Res. 4:329-838.
Latimer, W. M. 1952. The Oxidation States of the Elements and Their Poten-
tials in Aqueous Solutions. -Prentice-Hall, Inc.
Magno, P., T. Reavey, and J. Apidianakis. 1970. Liquid Waste Effluents from
a Nuclear Fuel Reprocessing Plant. 6RH-NERHL-7Q-2. .
Osterberg, C., N. Cutshall, V. Johnson, J. Cronin, D. Jennings, and
L. Frederick. 1966. Some Non-Siological Aspects of Columbia River Radio-
activity. IN: Disposal of Radioactive Wastes into Seas, Oceans and Surface
Waters. IAEA-SM-72/T8A, pp. 321-335.
Saas, A. and A. Grauby. 1973. Mechanisms for the Transfer to Cultivated
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River-Irrigated Soil—Ground Water. IN: ' Environmental Behavior of Radio-
nuclides Released in the Nuclear Industry. IAEA-SM-172/57, pp. 255-269.
Shade, J. W. 1974. Reaction of Hanford Sediments with Synthetic Waste, A
Reconnaissance Study. ARH-CD-176.
Sillen, L. G. and A. E. Martell. 1964. Stability Constants of. Metal-Ion
Complexes. The Chemical Society.
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Some Major Units of the Eartn's Crust. Bull. Geol. Soc. America. 72:175.
Wagman, D. D., W. H. Evans, V. 8. Parker, I. Halon, S. M. Bailey, and
R. H. Schumm. 1968. Selected Volumes of Chemical Thermodynamic Properties.
U.S. Department of Commerce.
Wallace, A. 1969. Behavior of Certain Synthetic Chelating Agents in Soil
and Biological Systems. UCLA-34-P-51-26, 92 p.
Ward, F. N. and H. W. Lakin. 1954. Determination of Traces of Antimony in
Soils and Rocks. Analytical Chemistry. 26:1168-1173.
Weast, R. C., Editor. 1976. Handbook of Chemistry and Physics. The Chemical
Rubber Co., Cleveland, Ohio, pp. B301-B302.
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CERIUM
Natural Soil and Rock Distributions
The abundance of cerium in rocks is given in Table 3-15. Cerium generally
increases in concentration from basaltic to granitic rocks. In a study of
Russian platform soils, Balashov et al. (1954) found that the soils with
highest cerium contents were alkaline, suggesting precipitation as the
hydroxide. The acid soils were lower in cerium because cerium had been
removed and had migrated. Soils derived from granitic rocks could be expected
to show the highest cerium contents if the weathering resulted in an alkaline
environment. Vinogradov (1959) gave the average cerium content of soils as
50 ppm. Half, or more, of the cerium in rocks is found in accessory minerals
such as apatite, while the rema-inder also substitutes for calcium but only in
the main- stage minerals such as plagioclase.
TABLE 3-15. CERIUM ABUNDANCE IN ROCKS
Rock Type • Ce, ppm
Continental basalts " 59
Kimberlite 1T9
UHramafic 0.1
Basaltic 48
Granitic, high Ca 81
Granitic, low Ca 92
Syenites 160
Shales 59
8616 Russian platform shales 67
6051 Russian platform sandstones 33
11205 Russian platform limestones 6.5
average of above 57
Sandstones 92
Dolomite and limestones 11.5
Reference
Frey et al., 1968
Burkov and Podporina, 1966
Turekian and Wedepohl, 1961
Turekian and Wedepohl
Turekian and Wedepohl
Turekian and Wedepchl
Turekian and Wedepohl
Turekian and Wedepohl
Ronov et al., 1967
Ronov et al., 1967
Ronov et al., 1967 .
Ronov et al., 1967
Turekian and Wedepohl
1961
1961
1961
1961
1961
1961
Turekian and Wedepohl, 1961
3-27
-------
Brief Chemistry
Cerium occurs naturally in four stable isotopes shown in Table 3-15.
141 144
The radioisotopes Ce and Ce are both present in radioactive wastes as
fission products with half-lives of 33 days and 284 days, respectively. After
100 years, little fission product cerium remains.
TABLE 3-16. STABLE ISOTOPES OF CERIUM
(INGHRAM ET AL,, 1947)
Isotope Natural Abundance, '•
136Ce 0.15
• 138C2 0.25
14°Ce 88.5
142Ce 11.1
The usual oxidation state of cerium is tripositive with .ionic radius
of 1.03 A (Ahrens, 1952).
The solubility product of Ce(III) hydroxide was given by Vickery (1953)
20
as 1.5 x 10" . For comparison, the solubility products of Al(OH)., and
1 c ^"1Q \ J
Fe(OH)3 are 3.7 x 10 and 3.8 x 10 , respectively (D'Ans and Lax, 1967).
In geologic environments, cerium forms weak complexes. Cerium complexes
of interest to soil environments will be reviewed in a subsequent section.
The most important and the most common cerium minerals include monazite (light
rare earths, Th) P04 and bastnaesite (light rare earths, Th) FCO-j. Many
other cerium-containing minerals exist but most are relatively rare.
Solid Phase and'Solution Equilibria
The relative stability of various cerium solid phases is shown in Fig-
ure 3-5. The thermodynamic data of the compounds used in Figure 3-5 were
obtained from Schumm et al. (1973), for Ce203 and CeO^, Baes and Mesmer
(1976) for Ce(OH)3 and Sillen and Martell (1964) for CePO^. Among the solids
depicted in Figure 3-5, Ce02 is the only one that changes its position with
a change in oxidation-reduction conditions. In an oxidizing environment
(p02 < 23.68 atm), the solids in increasing order of stability are ^2^3'
3-28
-------
-10
u
S1
-12
-14
-16
-13
AT pCj-0.68
Ce(OHL:Ce02ATp02 -77.22
ATp02 "0.68
9
PH
10
12
Figure 3-5. The activity of Ce in equilibrium with phosphate
levels from Variscite and Gibbsite (V and G), Dical:
cium Phosphate Dihydrate (DCPD) and OctacaTcium
Phosphate (OCP).
Ce(OH)3, CeP04, and CeOg. Ce2S3 is too soluble to plot in Figure 3-3 in an
oxidation-reduction environment where S04 will be the most dominant sulfur
species. In oxidation-reduction conditions corresponding to pG^ of >E3.68,
CePO, would be the most stable solid phase.
The activity of various solution species of cerium in equilibrium with
CePO,(solid) and under such concentrations of F", Cl~, S0^~, NO^, and PO^"
as commonly found in the environment is shown in Figure 3-6. The thermody-
namic data for all of the hydrolysis species were obtained from Baes and
2+
Mesmer (1976). The data for CeNO^ were obtained from Si 11 en and Martell
(1964). All the other data were obtained from Shumrn et al. (1973). Since
Figure 3-6 is based on solution complexes in equilibrium with Ce(III) as
CePO,, Ce(IV) ions and complexes would decrease in activity with increase in
reducing conditions. In the environmental pH range of interest the Ce(IV)
ions and complexes do not contribute significantly to the total cerium con-
centration in solution.
3-29
-------
Figure 3-6. The activity of various cerium species in equilibrium
with CeP04fs,) in an oxidizing soil environment Cp02(g).
= 0.68 atmj, pF" = 4.5, pCl" = pS04- = 2.5, pNOs' =3.0
and pH9PO,,~ = 5.0.
Trivalent cerium species in increasing order of importance are CeCl ,
CeiNO,"1", CeF , Ce , and CeSOt. Thus, CeSO. would be the most dominant
3 44
solution species in the environmental pH range of interest (pri 4 to 8).
Experimental Adsorption Results
Nishita et al. (1956) studied five soils and two clay minerals in systems
with 144Ce. Pretreatment of bentonite with dilute acid gave.much less cerium
uptake by the clay than the untreated bentonite. The hydrogen ion was more
144
strongly adsorbed than Ce. The pH of the leaching solution had a profound
effect on release of cerium by the soils. At pH values less than 5, 99.5%
144
recoveries of Ce were reported.
144
Rhodes (1957) studied the uptake of Ce on a Hanford soil as a function
of pH. The Kd values rose steadily from 3.0 at pH 1.6 to >1980 at pH 6.1. .
The Kd then decreased to a minimum of about 100 at pH 10 and increased again to
>1980 at pH 12 and above. The reduced uptake by soil of cerium between pH
7 to 12 did not occur in the presence of macroconcentrations (3M) of sodium
ions. Thus, the Ce was present either as a radiocolloid in this region,
which was flocculated by the high concentration of sodium ions, or as an ion
3-30
-------
species that was altered by the sodium ions to increase uptake by the soil.
_p
The complete removal of 4.2 x 10 M cerium at pH 5.5 to 6 was reported, which
at this concentration exceeded the capacity of the soil by ten times. The
precipitation of cerium hydroxide was indicated.
Bensen (1960) reported that cerium adsorbed on the soil could not be
exchanged above pH 7.4 and was difficult to exchange at pH 5.5 to 7.4. The
fraction of cerium exchanged by barium at pH 5.5 was 10% and virtually none
at pH 7.4. All of the cerium could be removed by ion exchange at pH 2.2,
with less removal as a function of pH to 7.4 when none of the cerium was
exchangeable from the soil. Above pH 7.0, the presence of 0.5M alkali metal
cations and 0.25M alkaline earth metal cations did not affect cerium removal
from solution by soil. Below pH 3.0, the adsorption of cerium was depressed
similarly by accompanying salts, implying that the cerium is ionic below
pH 3.0 and removed by ion exchange on the soil.
Bensen (1960) reported that the soil uptake of cerium added as Ce(III)
was identical to that added as Ce(IV). However, when an oxidizing agent
(Na3i03) was added, uptake of Ce(IV) increased to 99% at pH 1, while only
58% of the Ce(III) was. adsorbed at this pH. Bensen concluded that the
absence of sufficient holding oxidant resulted in an immediate reduction of
Ce(IV) to ee(III) by contact with the soil.
Kokotov et al. (1962) investigated Ce adsorption by soils and found
that cerium adsorption was much less below pH 2.0 due to increased hydrogen
ion competition.
. Bochkarev et al. (1964) reported that 85% of 144Ce applied to soil sur-
faces was accounted for 2 years following the application in the top 5 cm of
medium and coarse clay loams.
Kampbe-11 (1964) investigated Ce uptake by colloidal suspensions of
kaolinite, illite and montmorillonite type clays, all of which had a high
144
affinity for Ce. Only under very acid conditions did the cerium remain
in solution. DTPA also was effective in keeping cerium in solution at
trace levels.
3-31
-------
Dlouhy (1967) studied the distribution coefficients of cerium on
Casaccia soil and..tuff both in the laboratory and in field experiments. The
laboratory distribution coefficients ranged from 1050 to 1300 cm /g for the
3
cerium on the soil and 3000 to 3800 cm /g for cerium on the tuff. The solu-
tion conditions and soil and tuff were not characterized. Coring of seepage
pits showed the cerium to be essentially contained in the upper 6 cm of the
pit soil.
Molchanova (1968) studied the adsorption of Ce by soils and found
batter uptake at acidic pH values above 2.0 than at neutral to alkaline values.
Cerium was more easily removed from soils by iron salts than aluminum or cop-
per salts. The presence of iron colloids in the soil reduced the sorption of
144
'Ce more sharply than did iron salts.
Brown et al. (1969) studied the formation of Ce colloids and the
adsorption of'cerium by humic acid and bentonite suspensions. As ionic
strength and cerium concentration decreased and pH increased, equilibrium
shifted toward radiocolloidal cerium. At pH 6.5 and ionic strength of 10" ,
about 90% of the cerium was in the radiocolloidal form. A multiple regression
equation is given that shows the effects of pH, ionic strength and cerium
concentration on radiocolloidal cerium in solution.
Migration Results
Field Studies—
Magno et al. (1970) determined the forms and relative mobilities of
radionucl'ides from the lagoon system of the Nuclear Fuel Services plant in
144
Western New York state. They reported that 0.02 Ci of Ce were discharged
144
from the last lagoon from May to October 1969. None of the-discharged Ce
144
activity was dissolved in the water. All of the Ce activity discharged
144
was found in the suspended solids. Approximately 8.5 Ci of Ce were dis-
charged from the plant to the lagoon system over the same time period (May-
October 1969). Thus, although all of the cerium was in particulate form,
about 3% migrated through the lagoon system in the form of suspended solids.
The migration of 435,000 liters of liquid high-level waste (^5M NaN03,
1M NaOH) from 241-T-106 underground tank through Hanford soils was followed
by the drilling of several sample wells -after the event (Anon., 1973). The
3-32
-------
contaminated soil zone was determined on the basis of these analyses and
the relative migration distances of the radionuclides evaluated. The rare
earths, including cerium, were more mobile than plutonium and cesium but
less mobile than antimony, cobalt and ruthenium. Cerium could be expected
to be at least partially complexed in such a high salt, high pH solution.
Brookins (1976) reported that the rare earth elements including cerium
have been retained at the site of a 1.8 billion year old fossil nuclear
reactor at Oklo, Gabon. The rock surrounding the reactor is shale infilling
a fracture system in organo-argillaceous sandstone.
Laboratory Studies—
Nishita et al. (1956) studied the extractability of Ce from clays and
soils with distilled water and 1M ammonium acetate. The fraction remaining
on the soil or clay mineral (bentonite and kaolinite) after the acetate
elution or leaching solution, was passed through the column was called the
144 144
nonexchangeable Ce fraction. For both the soils and clays, Ce was
hardly moved by the water leaching or exchanged from soils or bentonite.
144
Kaolinite showed a large portion of exchangeable Ce and a relatively small
144 144
-amount-of—nonexchangeable Ce. Most of the Ce on the soils and on the
bentcnite were in the nonexchangeable fraction. The somewhat more acid
nature of the kaolinite probably accounts for the larger percentage of cerium
associated with it as exchangeable cerium.
Schulz (1965) included cerium and the other rare earths under the cate-
gory of immobile radioactive elements. These elements were precipitated as
hydroxides or carbonates or were very strongly bound by soil clays. Perhaps
144
this view should be somewhat tempered by the previous cases of Ce migration
as a particulate or in complexed species.
Eichholz et al. (1967) studied the fractionation of several radioactive
elements between several natural waters and the suspended solids contained
in them. The results gave an indication of the ability of cerium to
migrate as a suspended solid or adsorbed on the suspended solids already
present in most natural waters. The suspended solids are characterized in
Table 3-17. The most ob.vious relationship is a direct one between the sus-
pended solids content and cerium adsorption. However, there also is a
3-33
-------
TABLE 3-17. ADSORPTION OF 144Ce ON SUSPENDED
SOLIDS (EICHHOLZ ET AL., 1967)
Suspended Dissolved
Water Source
Colorado River
Camp McCoy
Bayou Anacoco
Lodgepole Creek
Chattahoochee River
1 Ce Removed, %
72.0
44.6
36.9
96.9
26.1
Solids, mg/1
299
12
24
965
131
Solids, mg/1
350
60
63
200
31
£H_
7.5
6.9
6.2
6.8
7.3
suggestion of an inverse relationship between cerium adsorption and solution
pH. At the lower pH, more of the cerium should be present as cationic species
and hence better able to adsorb on the suspended solids. It should also be
pointed out that the Lodgepole Creek solids are largely montmorillonite, a
high exchange capacity clay mineral.
Summary
Mo information is available on solid compounds of cerium that may be
presant in soils and sediments. It appears thatCePO* (Figures 3-5 and 3-6),
especially in alkaline conditions, may form in soils and sediments and may
control cerium concentrations.
In all terrestrial environments Ce(IV) and its complexes would have
insignificant effects on cerium concentration (Figure 3-6). Thus, cerium
would be expected to be present as Ce(III) (Figure 3-6 and Ahrens, 1952;
Bensen, 1960). In the absence of strong complexing 1igands,~and especially
in acidic environments, cerium would be expected to be present as Ce (Fig-
ure 3-6; Bensen, 1960) and ion exchangeable. The increase in pH has been
shown to increase the cerium adsorption (Rhodes, 1957; Bensen, 1960;
Nishita, 1956; Kokotov et al, 1962; Kampbell, 1964). As expected, the
presence of complexing ligands such as DTPA increases cerium concentration
in solution and reduces adsorption (Kampbell, 1964). In slightly acidic
(pH 6.5) to alkaline conditions, cerium is thought to be present as radio-
colloid which has been shown to increase adsorption (Rhodes, 1957; Brown
et al., 1969). Kampbell (1964) reported that kaolinite, montmorillonite,
and ill He have a strong affinity for cerium. Cerium adsorbed on suspended
3-34
-------
solids has been shown to migrate in field and laboratory studies (Magno.
et al., 1970; Eichholz et al., 1967). However, this type of movement would
be expected to be dependent upon the particle size of the suspended solids
i
and the pore size distribution of soils and rocks.
References
Ahrens, L. H. 1952. The Use of lonization Potentials. Part I. Ionic Radii
of the Elements. Geochim. et Cosmochim. Acta. 2:155.
Anonymous. 1973. 241-T-106 Tank Leak Investigation. ARH-2874.
Baes, C. F., Jr. and R. E. Mesmer. 1976. The Hydrolysis of Cations. John
Wiley and Sons, New York.
Balashov, Y. A., A. B. Ronov, A. A. Migdisov, and N. V. Turanskaya. 1964.
The Effects of Climate and Facies Environments on the Fractionation of the
Rare Earths During Sedimentation. Geochemistry International. 10:951-969.
Bensen, D. W. 1960. Review of Soil Chemistry Research at Hanford. HW-6720!.
Bochkarev, V. M., Z. G. Antropova, and E. I. Belova. 1954. Migration of
Strontium-90 and Cerium-144 in Soils of Different Texture. Soviet Soil
Science. 9:936-938.
Brookins, D. G. 1976. Shale as a Repository for .Radioactive_Was_te:__ The
Evidence from Oklo. Environmental Geology. 1:255-259.
Brown, R. E., R. E. Franklin, and R. H. Miller. 1969. Reactions of Cerium-144
in Solution and Suspensions of Soil Humic Acid and Bentonite. Soil Science
Society of America, Proceedings. 33:677-681.
Burkcv, V. V. and Ye. K. Podporina. 1966. First Data on Rare Earths in
Kimberl.ite. Dole1.. Acad. Sci. USSR. Earth Science Section. 171:215-219.
D'Ans, J. and E. Lax. 1967, Taschenbuch fiir Chemiker and Physiker. 3 Auf1.,
Bd. I. Springer.
Dlouhy, A. 1967. Movement of Radionuclides in the Aerated Zone. _IN_: Dis-
posal of Radioactive Wastes into the Ground. IAEA-SM-93/18, pp. 241-249.
Eichholz, G. G., T. F. Craft, and A. N. Galli. 1967. Trace Element Frac-
tionation by Suspended Matter in Water. Geochim. et Cosmochim. Acta. 31:
737-745.
Frey, F. A., M. A. Haskin, J. A. Poetz, and L. A. Haskin. 1968. Rare Earth
Abundances in Some Basic Rocks. J. Geophys. Res. 73:6085-6098.
Inghram, M. G., R. J. Hayden, and D. C. Hess, Jr. 1947. The Isotopic Consti- '
tution of Lanthanum and Cerium. Phys. Rev. 72:967-970.
3-35
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Kampbell, D. H. 1964. Cerium, Iron, and- Manganese Sorption by Soil Colloids
and Uptake by Plants. Thesis. Univ. of Missouri.
Kokotov, Yu. A., R. F. Popova, I. Tsing-Chih, and M. Shih-tsi. 1962. Scrp-
tion of Long Lived Fission Products by Soils and Argillaceous Minerals. II.
Sorption of Cerium-144 by Soils. Radiokhimya. 4:227-228.
Magno, P., T. Reavey, and J. Apidianakis. 1970. Liquid Waste Effluents
from a Nuclear Fuel Reprocessing Plant.. BRH-NERHL-70-2.
Molchanova, I. V. 1968. Behavior of Cerium-144. in Various Tvpes of Soil.
Tr. Inst. Ekol. Rast. Zhivotn. 61:4-11.
Nishita, H., B. W. Kowalewsky, A. J. Steen, and K. H. Larson. 1956. Fixation
and Extractability of Fission Products Contaminating Various Soils and Clays:
I. Strontium-90, Ruthenium-106, Cesium-137, and Cerium-144. Soil Science.
81:317-326.
Rhodes, P. 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.
Ronov, A. B., Yu. A. Balashov, and A. A. Migdisov. 1967. Geochemistry of
the Rare Earths in the Sedimentary Cycle. Geochemistry International. 4:1-17.
Schulz, R. K. 1965. Soil Chemistry of Radionuclides. Health Physics. 11:
1317-1324.
Schumrn, R. H., D. D. Wagman, S. M. Bailey, W. H. Evans,"and V. B. Parker. 1973.
Selected Values of Chemical Thermodynamic Properties. Tables for the Lanthanide
(Rare Earth) Elements (Elements 62 through 76) in the Standard Order of Arrange-
ment. NBS Technical Note 270-7.
Sillen, L. fr. and A. E. Martell. 1964. Stability Constants of Metal-Ion
Complexes. Special Publication No. 17. The Chemical Society, London.
Turekian, K. K. and K. H. Wedepohl. 1961. Distribution of the Elements in
Some Major Units of the Earth's Crust. Bull. Geol. Soc. America. 72:175.
Vickery, R. C. 1953. Chemistry of the Lanthanons. Butterworths Science
Publications.
Vinogradov, A. P. 1959. The Geochemistry of Rare and Dispersed Chemical
Elements in Soils. Consultants Bureau, Inc.
3-36
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CESIUM
Natural Soil and Rock Distribution
The cesium content of several rock types is given in Table 3-18. Bertrand
and Bertrand (1949) studied the soils of France and Italy and found the largest
concentration of cesium in soils over igneous rocks and in alkaline soils. .A
range of 0.3 to 25.7 ppm Cs was reported. Approximately 5 ppm Cs was reported
by Vinogradov (1959) in soils of the Russian Plains and 1 ppm in Japanese
soils. An average of 1.1 ppm Cs was found in suspended material and 2.6 ppm
Cs in bottom sediments of seven western United States streams (Sreekumaran
et al., 1968). Hirst (1962) studied the cesium content of modern marine sands
and clays, and reported them to contain 1.0 to 3.2 ppm Cs and 7.9 to 13 ppm
Cs, respectively. In general, cesium is enriched in the potassium minerals,
although enrichment may only consist of a few hundred ppm cesium. The large
size of the cesium cation makes it difficult for it to substitute in crystal
positions normally occupied by potassium, its nearest geochemically abundant,
alkali metal relative. Hence, cesium often becomes concentrated enough to
form its own mineral phase (pollucite).
TABLE 3-18. CESIUM CONTENT OF ROCKS
Rock Type Cs Content, ppm
1.1
Reference
Basalt (estimated average)
Basalt (estimated average)
Basalt (estimated average)
Basalt (estimated average)
Granodiorite
Granite
Granite
Granites, Russia
Marine shales
Sandstone
Limestone
Modern sediments
1.7
1.0
1.0
3.0
6.8
5.0
5.5
5.0
1.0
1.0
6.0
Turekian and Wedepohl,
1961
Heier and Adams, 1964
Taylor, 1964
Taylor and White, 1966
Heier and Adams, 1964
Heier and Adams, 1964
Taylor, 1964
Beus and Fabrikova,
1961
Horstman, 1957
Horstman, 1957
Horstman, 1957
Horstman, 1957
3-37
-------
Brief Chemistry
The only stable isotope of cesium is Cs. Fission product cesium
radioisotopes include four main cesium isotopes, with only Cs, Cs and
Cs (half-lives of 2.05 years, 3 x 10 years and 30.23 years, respectively)
of significant concentrations 10 years after separation from PWR (Pressurized
Water Reactor) fuels (Schneider and Platt, 1974). Pollucite, (Cs, Na)
(AlSipOg) x HLO containing 22 to 36 wt% cesium, is the only independent cesium
mineral of any importance.
In all its natural compounds, cesium occurs as a monovalent cation with
o
a radius of 1.67A (Ahrens, 1952). There is, little, if any, tendency for
cesium to form complexes in natural environments, and the solubility of most
cesium compounds in water is very high.
Solid Phase and Solution Equilibria
Thermodynamic data was collected for the following solid phases of cesium:
Cs(OH), Cs20, CsCl, CsC104, Cs2S04, Cs2C03, CsHCOp CsNO,, CsF. All are highly
soluble and, therefore, solid phase diagrams for cesium are not presented.
Baes and Mesmer (1976). report that cesium may be associated with OH" ions
in solution and'that the extent of this association cannot be estimated accu-
rately. Chariot (1957) reported that cesium rarely forms solution complexes.
Therefore cesium would be expected to appear in solutions predominantly as
Cs . Only 1.5% of the cesium from underground nuclear explosions was found
(Izrael et a!., 1970) to be colloidal,' which may have been adsorbed on other
colloids. The main soil and rock reaction with cesium is expected to be ion
exchange.
Experimental Adsorption Results
The effect of cesium concentration and pH on cesium adsorption by a
Hanford calcareous soil was studied by McHenry (1954). The data indicated
that trace cesium concentrations are essentially completely adsorbed above
pH 4.0. When accompanied by 4M NaCl, however, only up to 75% of the trace
cesium was adsorbed and the adsorption was nearly independent of pH over a
wide range. When the concentration of cesium and competing ion was at 50%
of the soil ion exchange capacity, competing cation depressed, cesium adsorp-
tion by soil in the order: H > La £ Sr > Na > NH^.
3-38
-------
At 1% of the soil capacity and lower cesium concentration, competing
cation effects on cesium adsorption were slight. 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 in.
greater concentrations. The presence of small amounts of mineral phases in
the soil with high cesium selectivities may account for this behavior (Ames,
1959; Jacobs and Tamura, 1960; Tamura and Jacobs, 1950; Schultz et al., 1960;
Sawhney, 1964; Gaudette et al., 1966). Mica-like minerals such as illite
o
tend to "fix" cesium. When expanded layer silicates were collapsed to a 10 A
mica spacing by potassium saturation and heating, increased cesium selectivity
resulted (Tamura and Jacobs, 1961; Tamura, 1963; Coleman et al., 1963). Some
o
researchers have considered the exchange of trace cesium on 10 A. spaced
minerals to be irreversible (Klechkovski and Gulyakin, 1958; Spitsyn et al.,
1963), but the reverse rate may only be much slower than the adsorption rate
(Routson, 1973).
Several authors have determined cesium distribution coefficients between
various solutions and rocks. These are listed in Table 3-19. Others have
determined adsorption isotherms on several natural materials, and from the
isotherms, derived thermodynamic functions for the specific cation exchange
reactions. Some of the thermodynamic studies, include Thomas (1967), Tamers
and Thomas (1960), Merriam and Thomas (1956), Gaines and Thomas (1953),
»
Lewis and Thomas (1963) and Eliason (1966) on cesium ion exchange on systems
of cations with clays. Ames (1959), Frysinger (1962) and Howery and Thomas
(1965) studied ion exchange and cesium thermodynamics on the natural zeolite
clinoptilolite.
Prout (1958) studied the effects of system variables on Savannah River
soil (80%. sand, 20% clay) uptake (Kd values) for trace cesium, strontium and
Plutonium. The pH and adsprbed ion concentration were some of the variables
examined. Cesium adsorption fell off rapidly with pH at less than 6 and
decreased as more than 10% of the soil cation exchange capacity was occupied
by cesium.
Berak (1963) has contributed much of the cesium adsorption work on dif-
ferent rock types and minerals. The standard equilibrating solution used is
-5
given in Table 3-20. Carrier cesium 1 x 10 M as CsCl was used along with the
3-39
-------
TABLE 3-19. CESIUM DISTRIBUTION COEFFICIENTS FOR SEVERAL ROCKS,
MINERALS AND SOIL TYPES
Medium
CsKd, ml/q
Condition
Reference
Al luvium
Central Nevada
Desert Alluvium
Hot Creek Valley,
Nevada
Tuff, Rainier Mesa,
NTS, Nevada
Tuff, Rainier Mesa,
NTS, Nevada
Carbonate, Yucca
Flat, NTS,. Nevada
Granodiorite,
Climax Stock,
NTS, Nevada
Alluvium, INEL,
Idaho
Granite,
Central Nevada
SasaH, NTS,
Nevada
Basalt
Amchitka
Island
Shaley siltstone
New Mexico
Sandstone,
New Mexico
SUPERSATURATED ROCKS
Rhyolites,
Czechoslovakia
121-3165
70-2640
1020
12,100-17,800
13.5
8-9
1030-1810
285-360
450-950
34.3
792-9520
280
39
6.5(seawater)
309
102
(containing quartz)
100,000
33
15 "
67
20
354
4
72
50,000
107
57
500-4000 -jm
500-4000 urn
>400 urn
100-200 mesh
>4000 urn
100-200 mesh
0.5-1.0 mm
(Laboratory soil )
(Field soil) •
>4000 urn
32-80 mesh
500-4000 pm
500-4000 urn
500-4000 urn
>4000 pm
>4000 um
0.1-0.2 mm
c
Nork et al . , 1971
Nork and Fenske, 1970
Nork and Fenske, 1970
Goldberg, et al., 1962
Nork and Fenske, 1970
Beetem et al., 1962
Schmalz, 1972
Nork and Fenske, 1970
Angelo et al., 1962
Essington and Nork,
1969
Nork and Fenske, 1970
Nork and Fenske, 1970
Berak, 1963
3-40
-------
TABLE 3-19. (continued)
Medium
Central Slovakia
Volcanic Glasses,
East Slovakia
Central Slovakia
Iceland
CsKd. Ml/q
40
31
32
11
9
50,000
23
12,400
42
25
26
45
41
42
39
23
18
:23 ; -
13
31
10
34
13
34
16
33
72
12
Condition
Reference
0.1-0.2 mm
Berak, '963
Rhyodacites,
East Slovakia
Dacites, Slovakia
45
49
85
96
41
0.1-0.2 mm
Berak, 1963
3-41
-------
TABLE 3-19. (continued)
Medium
CsKd, ml/q
Condition
Reference
Quartz Porphyries
Central Bohemia
Bohemia
North Bohemia
SATURATED ROCKS (contain
Trachytes,
North Bohemia
Slovakia
Andesites,
Central Slovakia
Pyroxenic, East
Slovakia
Garnetic, East
Slovakia
Altered, East
Slovakia
Pyroxenic,
Central Slovakia
Propyllitized
Granul. pyrox.,
vitreous
Spillites,
albitized
Diabases,
Bohemia
North Moravia
52 0.1-0.2 mm
32
730
104
32
no quartz or feldspathoids)
0.1-0.2 mm
89
16
54 0.1-0.2 mm
52
77
34
58
' 73
27
39
160
277
56
317
10 0.1-0.2 mm
61 0.1-0.2 mm "
14
37
93
300
640
540
640
19
Berak, 1963
Berak, 1963
Berak, 1963
Berak, 1963
Berak, 1963
3-42
-------
TABLE 3-19. ('continued)
Medium
CsKd, ml/g
Condition
Reference
Melaphyres,
basalt
North Bohemia
Slovakia
Basalts, North
Bohemia,
Vitreous
UNDERSATURATED ROCKS
Picrites, North
Moravia
Teschenites with
olivine, North
Moravia
Peridotite,
North Moravia
Pyroxenites, NW
Bohemia
0.1-0.2 mm
1330
108
37
122
900 0.1-0.2 mm
246
130
910
122
167
(nonfeldspathoidal )
456 0.1-0.2 mm
257
526
89
185
257
526
1330
456
30
318 0.1-0.2 mm
199
56
234
258
70
96 0.1-0.2 mm
2 0.1-0.2 mm
0
Berak, 1963
Berak, 1963
' Berak, 1963
Berak, 1963
Serak, 1963
Berak, 1963
3-43
-------
TABLE 3-19. (continued)
Medium
CsKd. ml/g
Condition
Reference
Monchi quite,
with olivine,
North Moravia
Ouachitite,
with apatite
North Maravia
Olivine basalts,
North Bohemia
East Slovakia
Central Slovakia
North Bohemia
Andesitic, Central
Slovakia
UNDERSAT'JRATED ROCKS
Phono lites, North
Bohemia
Tephrites, North
Bohemia
Chabazite
Leuci te
Chabazite
Nephel ine
Nepheline-leucite
Leuci te
Basanites, North
Bohemia, Nepheline-
leucite
Nephel ine
95 0.1-0.2 mm Berak, 1963
28 0.1-0.2 mm Berak, 1963
52
1215 . 0.1-0.2 mm Berak, 1963
1370
1685
75
1023
170
185
(feldspathoidal)
55 0.1-0.2 mm Berak, 1963
34
1150
1600
594 0.1-0.2 mm Berak, 1963
140
589
517
675
1134
1062
397 0.1-0.2 mm Berak, 1963
98
44
3-44
-------
TABLE 3-19. (continued)
Medium
Leucite
Nepheline
Nephelinites, NW
Bohemia, Metilite
Leucite
Glivine-nepheline
basalt
Hauyne-nepheline
basalt
Leucitites, NW Bohemia
Nepheline
Melilitites, olive-
hauyne basalt
MINERALS
Quartz
Agate
Chalcedony
Opal
01ivine
Humi te
Chondrodite
Thaumasite
Leucophanite
Zircon
Topaz
Kyanite
Sillimanite
Staurolite
CsKd, ml/q
1550
3240
52
34
329
117
571
700
3746
1328
398
109
145
0
2
11
8-61
5
8
0
16
2
3
4
43
3
1
Condition
Reference
0.1-0.2 mm
Berak, 1963
0.1-0.2 mm
Berak, 1963
0.1-0.2 mm
0.1-0.2 mm
Berak, 1963
Berak, 1963
3-45
-------
TABLE 3-19. (continued)
Medium
Garnet, alamandine
Garnet, grossular
Garnet, andradite
Garnet, uvarovite
Hydrogarnet
Vesuvianite
Sphene
Rinkita
Axinite
Hemimorphite
Beryl
Oioptase
Tourmaline
Wollastonite
Rhodonite
Hypersthene
Aegirine
Augite
Diopside
Enstatite
Jeffersonite
Tremolite
Chrysoti1e
Sepiolite
Attapulgite
Polygorskite
Zoisite
Talc
Pyrophyllite
Biotite
Muscovite
Penninite
CsKd. ml/q
0
3
0 •
5
54
16
3
9
0
6
7
8
3 .
0
3
18
50
14
6
1
4
8
0
900
89
150
33
0
8
21
15
2
Condition
Reference
3-46
-------
TABLE 3-19. (continued)
Medium
CsKd. ml/q
Condition
Reference
Delessite
Sen' cite
mite
Glauconite
Celadonite
Serpentine
Kaolinite
Halloysite
Allophane
Hisingerite
Smectites
Nontronite
Apopnyllite
Melilite
Feldspars
Ortnoclase
Sanadine
Albite
Oligoclase
Andes ine
Laboradorite
Scapolite
Leucite
Analcite
Pollucite
Nepheline
Soda lite
4070
21
400
87
84
17
45
81
276
257
4900
2400
355
425
545
377
345
388 0.1-0.2 m Berak, 1963
8. -
68
0.1-0.2 mm . Berak, 1963
9
58
6
9
12
28
24
26
56
54
5
20
8
3-47
-------
TABLE 3-19. (continued)
Medium
Cancrinite
Zeolites:
Mordenite
Stilbite
Heulandite
Faujasite
Hannotome
Phillipsite
Chabazite
Natrolite
Scolecite
Thomsonite
Clinch River, TN
bottom sediments
Silty Clay, ID
Sandstone,fine
light gray
Shaley siltstone,
carbonaceous,
black
Sandstone, fine
light gray
Sandstone, very
fine, silty, dark
gray
CsKd, ml/q
67
25,000
19,900
193
1230
230
416
9900
4900
20
9
7050
2326 - 1 hr
50,152 - 3 days
88,048 - 7 days
3169-1 hr
50,152 - 3 days
82,769 - 7. days
3000
139
166
389
309
472
541
102
141
298
346
456
630
Condition
0.1-0.2 mm
pH fi
pH 9
Quartz, ill ite,
apatite
4000 urn
500 pm
<62 um
4000 um
500 um
<62 um
4000 um
500 ym
<62 um
4000 um
500 um
<62 um
Reference
Berak, 1963
Carrigan, et al,
1967
Wilding and Rhodes,
1963
Sokol, 1970
Sokol, 1970
Sokol, 1970
Sokol, 1970
3-48
-------
tracer cesium. A 1% solid suspension was equilibrated for 1 day at 20°C.
Results indicated that the secondary silicates (clays, zeolites) removed
cesium from solution much better than the primary silicates (feldspars, quartz,
etc.).
TABLE 3-20. THE EQUILIBRATING SOLUTION COMPOSITION
PLUS TRACE 137Cs THAT WAS UTILIZED IN
ALL OF BERAK'S (1963)Kd WORK WITH 24
HOUR EQUILIBRATIONS AT 20°C
Ions
Ca2+
Mg2+
Na"1"
SO2"
NOj
Cl"
HCO;
mg/1
80.2
12.2
23.0
96.0
62.0
35.4
61.0
0
0.0001
0.001
0.01
0.02
0.03
0.04
0.05
0.08
0.10
154
147
137
357
483
695
727
634
373
265
Jacobs (1962) used vermiculite to illustrate the favorable 10 A spacing
caused by potassium added to the cesium in the influent, solution. Cesium
selectivity (Kd) increases through.a maximum while the cation exchange
capacity of the vermiculite decreases as shown in Table 3-21.
TABLE 3-21. EFFECT OF ADDITION OF POTASSIUM TO THE
INFLUENT ON THE SORPTION OF CESIUM BY
VERMICULITE FROM 0.5M NaNOa CONTAINING
THE MASS EQUIVALENT OF 2 uc 137Cs/ml
(JACOBS, 1962)
Potassium CsKd, Final CEC,
Concentration. M ml/g meg/100 q
75.4
74.9
70.2
' 60.0
38.1
26.2
17.9
15.3
9.1
8.6
3-49
-------
The cation exchange capacities of five standard clays and seven soil
samples were determined using cesium, ammonium and manganese ions by Beetem
et al. (1962). These authors preferred the use of cesium chloride- Cs to
accurately determine clay and soil capacities for samples as small as 30 mg.
Rhodes and Nelson (1957) studied the effects of pH and sodium nitrate
concentrations on cesium distribution coefficients using a glaciofluviatile
silty sand composite recovered from Hanford wells. The results are given in
Tables 3-22 and 3-23. In high salt solutions cesium Kd values were much less
affected by pH variations than cesium tracer in water only.
TABLE 3-22. EFFECT OF pH ON 137Cs ADSORPTION. SOIL
WAS A COMPOSITIE GLACIOFLUVIATILE SILTY
SAND FROM HANFORD, WASHINGTON. (RHODES
AND NELSON, 1957)
CsKd, ml/g
H2° NaCl, 4M
36.5
138.0
200
>200 •
>200
>200
>200
>200
14.0
12.2
17.4
18.0
18.0
12.0
16.0
18.8
0.4
1.8
3.7
4.9
7.0
7.8
8.5
10.1
TABLE 3-23. EFFECT OF NaNOa CONCENTRATION ON 137Cs ADSORPTION.
SAME COMPOSITE QLACIOFLUVIATILE SOIL WAS USED AS
IN TABLE 3-20 (RHODES AND NELSON, 1957)
NaN03,M
0.25
0.50
1.0
2.2
3.6
4.8
6.0
CsKd, n
245
195
100
50
30
20
18
3-50
-------
Wahlberg and Fishman (1962) determined cesium Kd values with kaolinite,
illite, montmorillonite and halloysite over a large cesium concentration
range and in competition with sodium, potassium, calcium or magnesium. The
type and amount of competing ions and the cesium concentration determined
the cesium Kd variation with potassium having the most depressing effect.
Parsons (1962) determined a CsKd value of 100 in sand, sandy loam,
silt and humus at a groundwater pH of 5.4 to 6.6.
Janzer et al. (1962) determined cesium Kd values on a number of 100 to
200 mesh rock samples. The samples were divided into two parts. One part
of the sample was equilibrated with the solution without cesium tracer, and
finally with a cesium tracer plus the same solution. The other sample half
was directly equilibrated with the cesium traced solution. There were minor
differences between the Kd ranges of the two sample groups. The five
equilibrating solutions varied from 1,000 to 45,000 mg/1 in dissolved solids.
Equilibration times were up to 30 days and each experiment was run in tripli-
cate. The rock type descriptions and cesium Kd values are given in Table
3-24 as minimums for the highest dissolved solids solutions to maximums for
lowest dissolved sol.ids solutions. A membrane technique was used to.keep
solid and cesium traced solution phases separated.
TABLE 3-24. SUMMARY OF CESIUM Kd VALUES (JANZER ET AL., 1962)
Kd. ml/q
>le Number
45
95
320
331
350
365
380
479
510
526
552
620
653
680
Minimum
613
1,034
9
841
27
36
445
419
51
1,176
536
661
302
396
Maximum
10,300
13,400
48
15,530
170
1,140
6,657
3,600
2,824
8,990
7,150
8,250
3,766
5,280
Rock Description
Sandstone, calcite cement
Siltstone, sandy
Gypsum rocks and siltstone
Dolomite
Gypsum rock
Dolomite
Siltstone, silty sandstone,
dolomite cement
Gypsum rock, siltstone,
sandstone
3-51
-------
Baetsle et al. (1964) reported several cesium Kd values as a function of
pH and water composition for Mol , Belgium, soil. The Kd values are given in
Table 3-25 and may be directly compared with strontium Kd values on the same
soils. The causes of the low deionized water Kd value relative to the cesium
Kd values with mains (tap?) water are not known. Mercer (1967) gave cesium Kd
values as a function of sodium and calcium normality for clinoptilolite, a
natural zeolite. The Kd values varied from about 5 x 10 ml/g at 0.1N NaCl to
about 3 x 105 ml/g at 0.001N NaCl . At 0.001N CaCl2 the clinoptilolite cesium
Kd value was about 3 x 10 ml/g and 3 x 104 ml/g at 0.1N CaCl. Certain of
the secondary silicates, such as some of the zeolites, efficiently remove
cesium from solution even when the concentrations of competing ions are rela-
tively large. Rancon (1967) reported cesium Kd values of 21,000 on illite,
300 on kaolinite and 2000 ml/g on montmorillonite.
TABLE 3-25. VARIATION OF TRACE CESIUM Kd VALUES WITH pH (BAETSLE
ET AL., 1964)
Soil Type
Eollan Sand
Horizon A
Horizon B
Horizon C
Mol White Sand
Mol Lignitic Sand
Ames and Hajek (1966) statistically analyzed cesium adsorption data.
Hajek and Ames (1968) also showed the problems associated with determining
cesium distribution coefficients on the A horizon of Burbank loamy fine sand
without first equilibrating the soil with the solution minus the cesium tracer.
The resulting cesium.Kd values depend on the solution to soil ratio used in
the equilibration, a's seen in Tables 3-26 and 3-27.
A curve for trace cesium loading on a 1 g Burbank soil is also given.
Even these small columns, when used for Kd determinations, should be pre-
equilibrated with the solution minus the radioactive tracer.
3-52
Kd, ,nl/q
Deionized
Water, pH4
25-38
'31-32
226
50
-
Mains Water,
PH7.7 pH3
51 22.4
145 . 68
89 79
30
3.9-7.1 -
Groundwater,
pH4
10.1 '
-
-
1.3-3.2
3.4-4.3
-------
TABLE 3-26.
Solution
3N NaN03
0.5N Nad
0.5N CaCl
SOLUTION:SOIL RATIO EFFECT ON TRACE
CESIUM EQUILIBRIUM DISTRIBUTION
COEFFICIENTS (HAJEK AND AMES, 1968)
CsKd, ml/q
So1ution:Soi1
TOO
50
TO
466
1091
5211
380
1445
5362
271
1354
2420
206
1173
1260
TABLE 3-27. CESIUM EQUILIBRIUM DISTRIBUTION COEFFICIENTS
FOR SAND, SILT AND CLAY FRACTIONS OF BURBANK
LOAMY FINE SAND IN 0.5N NaCl (HAJEK AND AMES,
1968)
CsKd, ml/g
Clay
Solution: Soil
500 200 100
6557 6008 5364
Silt
Solution: Soil
50 10 5
548 472 408
Sand
Solution:Soil
50 10 5
1657 1014 893
Three soils representative of the Hanford Site, the Ritzville silt loam,
Burbank loamy sand and Ephrata sandy loam, were studied in detail for their
ion exchange properties with cesium (Routson, 1973). The three soils were
thoroughly characterized in terms of their physical properties (size fractiona-
tion), chemical properties (exchangeable Ca, Mg, Na and K in each soil horizon
and the total cation exchange capacity of each) as well as their mineralogical
properties (semi quantitative mineralogy of each size fraction, plus organic
carbon and calcite content of each fraction). Cesium equilibrium distribu-
tion coefficients from 0.2N NaCl were determined on the "as received" soil
horizon samples, sodium-based horizon samples and horizon samples with organic
carbon removed, calcite removed and both removed. The results are given in
Table 3-28. Soil samples were pre-equilibrated with a 0.2N NaCl except for
3-53
-------
the "as received" cases. The Ritzville soil averaged about 12 meq/100 g, the
Ephrata soil about 7 meq/100 g and Burbank about 5 meq/100 g cation exchange
capacities. The Kd values were directly correlated with the soil cation
exchange capacities.
TABLE 3-28. TRACE CESIUM ADSORPTION BY THREE HANFORD
SOILS FROM A 0.2M NaCl SOLUTION (ROUTSON,
1973)
Ritzville
A12
81
Cca
C
Burbank
A12
AC
AC2
1C
Ephrata
A12
81
IB2
IIC
As Received
4869
4332
3111
3529
2775
2680
2573
1588
3941
4696
3491
1919
Na-based
4869
4930
5455
. 6169
1964
2571
3533
1633
2385
2410
1950
1553
Organic C
Removed
2750
2154
2608
3510
2467
2682
2513
1155
3027
2835
2402
2012
Calcite
Removed
5980
6486
4060
5174
2540
2448
2070
1473
5230
4071
2316
2423
Organic C+
Calcite
Removed
4069
3152
2592
2578
3024
2338
2036
1010
2671
2901
1850
836
R. J. Serne of PNL, 1973, also determined cesium distribution coefficients
from high salt, high pH solutions. The solution compositions and sediments
are the same as those described in Table 3-9 and 3-10. The results of the high
salt work are provided in Table 3-29. Each cesium Kd was determined in tripli-
cate. The values given in Table 3-29 represent the average of the three Kd
values.
3-54
-------
Solution
I
13.41
9.37
9.06
8.11
10.93
7.88
8.74
II
6.00
3.47
3.42
3.08
3.70
2.04
2.79
III
1.80
1.37
1.27
1.37
1.83
0.93
1.14
IV
0.85
0.61
0.58
0.73
2.92
1.20
1.47
V
12.02
7.14
7.21
7.68
11.30
7.97
9.49
TABLE 3-29. CESIUM Kd VALUES BETWEEN SEVEN GLACIOFLUVIATILE
SEDIMENTS AND THE SOLUTIONS OF TABLES 3-9 AND
3-10
Sediments
1
2
3
4
5
6
7
Migration Results
Field Studies--
Mawson (1956) and Evans (1958) reported the results of studies on the
movement of fission products disposed to ground in a sand at Chalk River,
Ontario. The rate-of movement was followed with a series of wells from which
groundwater samples were collected and analyzed. Cesium moved the least and
at the lowest velocity of several radionuclides including strontium, cerium
and ruthenium in dilute, neutral wastes. Cesium was also the slowest radio-
nuclide to move in acid wastes as well because neutralization was soon accom-
plished due to soil buffering.
The drilling of core wells into abandoned covered trench disposal sites
at Hanford was used to determine cesium distribution in the field. Raymond
(1965) found that the bulk of the 137Cs was contained in the upper 3 ft
beneath the disposal site.
Brown (1967) reviewed the hydrology and geology of the Hanford areas and
gave well profiles of Cs at two disposal sites. Anomalies in the radionu-
clide distribution were correlated with stratigraphic features such as a
caliche (CaC03) bed and a loess lens. The average rate of downward moisture
movement decreased from 1.5 m/yr in 1958 to 1959 to 0.5 m/yr from 1963 to. 1966,
as a result of disposal site retirement in 1956. A study to determine Cs
3-55
-------
leaching by groundwater of contaminated soils from core wells showed that
137
50 column volumes were required to leach 11% of the Cs. Another 450
column volumes removed an additional 4% of the Cs. Trace Cs was added
to the groundwater. Distribution coefficients on sediments from below the
groundwater table were 300 ml/g for Cs. Eight relatively long-lived radio-
nuclides were detected in the groundwater including Cs. Cs was adsorbed
Q
in the upper 10 m of 216-S-l and -2 covered trenches. About 1.5 x 10 liters
of low salt wastes containing about 2000 Ci of Cs were disposed to 216-S-l
and -2 prior to shutdown in 1956.
Haney (1967) reported the disposition of 137Cs in the 216-BY covered
trenches that had received 3.4 x 10 liters of U-Plant high salt scavenged
waste that contained 4.1 x 10 Ci of undifferentiated 6 emitting radionuclides
including 3300 Ci of Cs. The 8 BY Covered trenches were in service from
December 1954 to December 1955 at Hanford. Monitoring and sampling wells
drilled into and near the ground disposal facility showed that the cesium peak
occurred at 28.7 uCi/g soil at about 3 meters below the bottom of the trench
90
and rapidly decreased with depth. Approximately 12,900 Ci of Sr also were
90
contained in the waste. The Sr peak soil loading occurred at about 5 meters
below the bottom of the trench.' Groundwater at the site was 67 m below
ground surface. Soil samples below 23 m contained less than 2 x 10" uCi/g
of either cesium or strontium.
137
The migration of Cs from Oak Ridge seepage pits was studied by
137
Lomenick et al. (1967). It was found that most of the Cs in the pits was
tied up in an illitic clay sludge. Water movement was found to be 0.15 m/day
while Cs movement was 0.127 cm/yr.
Magno et al. (1970) determined radionuclide migration through the efflu-
ent lagooning system of the Nuclear Fuel Services plant in Western New York.
From their analytical data, they estimated that 75% of the 134Cs and 137Cs
discharged from the plant were deposited in the sediments in the lagoon system.
From 40 to 95% of the cesium at the lagoon system exit was associated with
suspended solids.
In 1973, the 241-T-106 underground liquid high level waste storage tank
at Hanford leaked about 435,000 liters of waste into the surrounding sediments
(Anon., 1973). A series of wells were drilled to ascertain the locations and
3-56
-------
movements of the various radionuclides in the sediment column. Monitoring
and core sample analysis showed the relative mobilities of the radionuclides
in the 1M NaOH, 5M NaNO., solution. Plutonium was the least mobile, but
cesium mobility was only slightly more rapid. None of the radionuclides
migrated below 27 m. The water table was at 62 m depth.
Brookins et al. (1975) and Brookins (1976) have reported that a study
of the 135Ba and T37Ba daughters of 135Cs and 137Cs found at the prehistoric
natural reactor site at Oklo in Gabon indicate the probable retention of fis-
sion product cesium at the reactor site.
Laboratory Studies—
Nishita et al. (1960) studied the effects of potassium and cesium on the
137
release of Cs from three soils to plants. Addition of potassium was found
to be inefficient in reducing Cs uptake by plants. The addition of non-
137
radioactive cesium always increased plant uptake of Cs. even when the
cesium level was injurious to the plant.
Davis (1961) reported that the soil exchange capacity was essentially
constant above pH 3.5. Cesium was most effectively desorbed from soils by
neutral salts, and more readily by potassium than by sodium or calcium.
Trace amounts of cesium were held more strongly by soils' than trace amounts
of strontium.
Nishita et al. (1962) used an equilibrium method in a followup study to
determine the effects of stable cesium and potassium on the movement of
137 42
tracer quantities of Cs and K. The concentration of stable cesium and
137
potassium was an important factor. When Cs was only a negligible fraction
of the cesium in solution, only a negligible fraction of Cs was adsorbed.
In equimolar mixtures of cesium and potassium, the adsorption of cesium
relative to potassium decreased as the ionic strength increased. In low con-
137
centrations, cesium was more effective than potassium in releasing Cs from
soils. In high concentrations, potassium can be as effective as cesium in
releasing Cs.
The adsorption of cesium by soils and its displacement by salt solutions
were investigated by Coleman et al. (1963). Montmorillonite, illite and kao-
linite loaded with cesium were readily leached by IN KC1 or IN CaCl2.
3-57
-------
Vernriculite and heated (600°C for 12 hr) potassium montmorillonite bound
cesium tightly while contacting solutions of CaCU and AlCU but not KC1 or
NH.C1 solutions. Prolonged leaching with IN KC1 removed 97.5% of the cesium
adsorbed on heated potassium montmorillonite while IN CaC12 replaced 88%.
The exchange-displacement behavior of cesium suggests that its adsorption in
interlayer spaces allows admission of potassium and ammonium ions but
restricts the entry of calcium. Apparent specific adsorption of cesium
against exchange with CaCl- occurred at large cesium concentrations.
Miller and Reitemeier (1963) published some of their earlier work on
strontium leaching adding some additional cesium leaching work. Leaching
treatments consisted of applications of 30 and 300 in. of deionized water,
0.005N NaCl and 0.005N CaCl9. There was little downward movement of cesium.
134
Assays showed that 96.6 to 100% of the Cs was in the surface 1.4 in. of
the soil columns (Norfolk loamy sand, Hagerstown silt loam, Fort Collins loam,
Miami silt loam, Huntley clay) after 300 in. of leaching by the above
solutions.
134
Cesium K values for Cs in the saturated subsoil at Mol, Belgium, also
were given by Baetsle et al. (1964). The K values are listed in Table 3-30
89
and may be compared with similar values obtained on the same soils for Sr
152
and Eu. Mol sand is nearly pure quartz. The effect of pH on cesium migra-
tion in relation to water can be seen in the K values. In all cases-, except
for the perturbed profile, K values decreased (cesium mobility increased in
relation to water velocity) as the pH decreased. K values are markedly dif-
ferent even with single pH unit.
Cesium distribution coefficients at pH 3 and the rate of cesium movement
in relation to the groundwater movement in Houthulst clays are given in
89
Table 3-31. These values may be compared for the same soil samples with Sr
and 152Eu. Note that the K values for cesium decreased in the Houthulst
clays by an order of magnitude as compared to the Mol sands. It is unfortu-
nate that the minerals constituting the Houthulst clays were not specified.
Carlile and Hajek (1967) reported an example of physical transport
through a soil column. About 0.1% of the Cs was found in 2 cm diameter
by 40 cm length N-Area Hanford soil columns. Until 5 column volumes of
3-58
-------
TABLE 3-30. K VALUES OF CESIUM IN THE SATURATED SUBSOILS
AT MOL, BELGIUM (BAETSLE ET AL., 1964)
K IS DEFINED ON PAGE 2-30.
Soil Types
Mol sand
Perturbed profile
Mol sand + eolian sand
Mol sand + lignite
Mol sand
Mol sand
Mol sand
Mol sand
Mol sand
Mol sand
pH4
9
125
-
76
38
48
55
38
31
25
pH3
6
431
24
20
26
32
19
14
15
13
TABLE 3-31. CESIUM DISTRIBUTION COEFFICIENTS AND RELATIVE
MIGRATION RATES IN HOUTHULST CLAYS AT pH 3
(BAETSLE ET AL., 1964)
Depth, m Kd. ml/g
0-1
1-2
2-2.5
2.5-3
3-4
4-5
5-6
6-7
7-8
3-9
9-10
0-1
1-2
2-3
*3-4
4-5
5-6
6-7
7-8
8-9
9-10
111.5
77.0
83.0
47.7
-
145.2
133.5
107.0
-
103
163
-
31
29
165
342
570
693
494
462
441
Core Well
I
_£ Depth, m
372
257
278
160
-
485
466
357
-
344
543
Core Well
-
104
98
550
1140
1899
2309
1646
1536
1470
10-11
11-12
12-13
13-14
14-15
15-16
16-17
17-18
18-19
18-19
II
10-11
11-12
12-13
13-14
14-15
15-16
16-17
17-18
18-19
19-20
Kd, ml/a
198
295
386
475
427
742
765
484
470
250
536
527
664
374
247
291
158
220
193
241
1C
560-
933
1286
1583
1423
2472
2548
1613
1566
834
1786
1756
2212
1246
824
970
527
733
644
804
3-59
-------
Columbia River water (dissolved solids =119 mg/1) had passed through the
soil, leakage did not fall below detection limits (10 % leakage). Effluent
samples were centrifuged, treated with hydrogen peroxide to destroy organic
material, and recentrifuged. The peroxide treatment did not affect cesium
removal by centrifugation, which removed 50 to 75% of the cesium in either
case. It was found that the low ionic strength river water dispersed the
soil colloids on which the Cs was adsorbed. The breakthrough of cesium
continued until the colloids were flushed from the column. Preleaching with
river water rid the soil columns of colloids and resulted in no initial
cesium leakage.
Eichholz et al. (1967) investigated the partitioning of dissolved radio-
nuclides between suspended sediment particles and aqueous solutions. Their
work does not necessarily have a direct application to particulate migration
through soils, but it does give a feeling for the types and amounts of radio-
nuclides that can be transported on suspended solids. Several bodies of
water were sampled and utilized by the authors in the adsorption study as
sources of suspended solids. Some of the properties of the natural water.
samples are listed in Table 3-32. A fission product mixture was added; the
system allowed to come to equilibrium; and the water was recovered and passed
through a mixed bed ion exchange resin. The resin and sediment were then
counted with the radionuclides extraction results shown in Table 3-33. It
»
is noteworthy that a large portion (80 to 90%) of the radionuclides was asso-
ciated and traveled with the particulate matter in a highly concentrated form.
There is a direct correlation between cesium adsorption and the suspended
solids concentration, with some effect of dissolved solids in the water.
Levi and Miekeley (1967) reported that experiments with vermiculite and
the cesium fixation phenomenon showed that a lattice contraction was not
responsible for cesium uptake as previously suggested by others. Cesium fixed
with respect to isotopic exchange was shown to become mobile by exchange with
many other ions. Two different "defixation" mechanisms were indicated, one
caused by strongly hydrated ions and the other by small nonhydrated ions.
Cesium ions can be exchanged for potassium, but not potassium ions for cesium.
Strontium isotopic exchange curves were obtained with mixed Cs-Sr vermiculite.
Strontium ions can be exchanged in the system but at a reduced rate in their
3-60
-------
TABLE 3-32. PROPERTIES OF NATURAL WATER SAMPLES (EICHHOLZ ET AL., 1967)
Suspended Dissolved Conductivity
Source Solids, ppm Solids, ppm pH umhos
Colorado River,
Utah
.Camp McCoy,
Wisconsin
Bayou Anacoco,
Lousiana
Lodgepole Creek
Nebraska
Chattahoochee
River, Georgia
Billy's Lake
Georgia (swamp)
229
12
24
350
60
63
68
7.5
6.9
6.2
4.2
540
80
60
965
131
200
31
6.8
7.3
300
45
45
Solids
Composition
95% quartz,
5% calcite,
feldspar, 11 lite,
kaolinite
30% quartz +
feldspar,
6% kaolinite, 24%
illite
30% quartz,
2% kaolinite,
6% smectite
20% illite,
80% smectite
33% quartz,
44% kaolinite
23% illite
2% quartz,
balance was
amorphous
and/or organic
TABLE 3-33. ADSORPTION OF RADIONUCLIDES ON SUSPENDED SOLIDS
(EICHHOLZ ET AL., 1967)
Source
137
Cs,
Colorado River
Camp McCoy
Bayou Anacoco
Lodgepole Creek
Chattahoochee
River
Billy's Lake
83.9
3.3
16.2
96.3
26.1
0.3
mobility. The authors suggest that it may be possible to slow migration of
various radionuclides in soils by mixing them with cations fixed like cesium
in vermiculite.
Nishita and Essington (1967) continued work on the affects of chelating
agents on 89Sr, 91Y,
106Ru, 137Cs and 144Ce migration in different types of
soils. Movement by water leaching, irrespective of soils, was in the order
3-61
-------
137Cs, 91Y, 144Ce < 89Sr « 106Ru. Except in the Hanford sandy loam (a
California soil), practically no movement in Cs, Y or Ce occurred by
water leaching. EDDHA was generally the least effective, while the effective-
ness in promoting leaching of DTPA and EDTA varied with the soil type and
radionuclide. Soil pH probably had a large effect on complexing agent
effectiveness, but the causes were not determined.
Knoll (1969) reported that two waste organic solutions had detrimental
effects on retention of radionuclides by the soil. One was a mixture of 0.4M
di-(2-ethythexyl) phosphoric acid and 0.2M tributyl phosphate in paraffin
hydrocarbons (C-JQ to C^), and the other was hydroxyacetic acid. With the
0.4M D2EHPA-0.2TBP solution, strontium removal from a 5.0 cm soil column
was 70% at 10 column volumes. Only about 12% of the cesium was removed in the
same throughput, as was 80% of the adsorbed americium and 30% of the plutonium.
With a 0.25M hydroxyacetic acid solution, pH 3.7, plutonium, strontium and
americium removal was essentially complete at 10 column volumes. As the con-
centration of hyd'roxyacetic acid decreased, the migration rate of the strontium,
plutonium and americium also decreased. None of the cesium on the column was
eluted at any of the hydroxyacetic acid concentrations.
Salo and Saxen (1974) investigated the role of humic substances in the
transport of radionuclides. Atomic absorption and X-ray fluorescence were
used to determine the stable elements bound to the humic materials fractionated
into different molecular size groups. Surface waters from Bothnian Bay,
Finland, were used to study the molecular size distributions of humic sub-
stances, and Kd values were determined by gel chromatography. The distribu-
tion of strontium and cesium into different molecular size groups of the
humic substances were discussed. Cesium concentration was directly correlated
with the colored colloids, unlike the strontium, manganese and zinc. It was
postulated that the Cesium was adsorbed on the colored colloids. A wavelength
of 350 mm was used to determine color intensity.
Summary
Cesium mainly exists in solutions as Cs . Therefore, the principal
reaction mechanism of cesium adsorption on soils and rocks is expected to be
cation exchange. In some instances cesium shows a direct Kd-cation exchange
capacity relationship (Routson, 1973), and in other instances, no relation
3-62
-------
between Kd and exchange capacity (Jacobs, 1962). At trace cesium concentra-
tions competing cation effects are minor (McHenry, 1954), while at 10% or
more of the exchange capacity occupied by cesium, the Kd value beings to fall
off (Prout, 1958). Potassium is better than sodium or calcium for desorbing
cesium (Davis, 1961). There is a tendency for cesium to become "fixed" in
o
10A mica-like minerals (Tamura and Jacobs, 1961; Jacobs, 1962; Tamura, 1963;
Coleman et al., 1963; Lomenick, 1967), which may be due to a slow desorption
rate rather than fixation (Levi and Miekeley, 1967; Routson, 1973). The
secondary silicate minerals such as zeolites and clay minerals generally show
much higher cesium Kd values than the primary silicate minerals found in
igneous rocks (Berak, 1963).
The migration rate of cesium in the groundwater is faster at lower pH
values but slower through clays than through sands due to larger clay Kd
values (Baetsle et al., 1964). Much of the cesium that migrates in surface
waters does so adsorbed on suspended solids (Eichholz et al., 1967; Magno
et al., 1970).
Occasionally cesium can migrate through the soil column adsorbed on sus-
pended solids (Carlile and Hajek, 1967). . Chelating agents and organic soil
materials have little affect on cesium migration through soils (Mishita and '
Essington, 1967; Knoll, 1969; Salo and Saxen, 1974).
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Spitsyn, V., V. C; Balukova, and T. A. Ermanova. 1963. Studies on Sorption
and Migration of Radioactive Elements in Soil. IN: Treatment and Storage
of High-Level Radioactive Wastes. IAEA, Vienna, pp. 569-577 (in Russian).
Sreekumaran, C., K. t. Pillai, and T. R. Folsom. 1968. The Concentrations
of Lithium, Potassium, Rubidium and Cesium in Some Western American Rivers
and Marine Sediments. Geochim. et Cosmochim. Acta. 32:1229-1234.
Tamers, M. A. and H. C. Thomas. 1960. Ion-Exchange Properties of Kaolinite
Slurries. J. Chem. Phy. 64:29-32.
Tamura, T." 1963. Cesium Sorption Reactions as Indicator of Clay Mineral
Structures. IN.: Proc. Natl. Conf. Clays, Clay Minerals. 10:389-398.
Tamujra, T. 1963. Selective Ion Exchange Reactions for Cesium and Strontium
by Soil Minerals. Colog. Intern. Retention Migration Ions Radioactifs Sols,
Saclay, p. 95-104,
Tamura, T. and D. G. Jacobs. 1960. Structural Implications in Cesium
Sorption. Health Physics. 2:391-398.
Tamura, T. and D. G. Jacobs. 1961. Improving Cesium Selectivity of Bentonites
by Heat Treatment. Health Physics. 5:149-154.
3-67
-------
Taylor, S. R. 1964. Abundance of Chemical Elements in the Continental
Crust: A New Table. Geochim. et Cosmochim. Acta. 28:1273-1285.
Taylor, S. R. 1964. Trace Element Abundances and the Chondritic Earth
Model. Geochim. et Cosmochim. Acta. 28:1989-1998.
Taylor, S. R. and A. J. R. White. 1966. Trace Element Abundances in
Andesites. Bull. Volcano!. 29:177-194.
Thomas, H. C. 1967. The Thermodynamics of Ion Exchange on Colloidal Mate-
rials with Applications to Silicate Minerals. Thesis. University of North
Carolina.
Turekian, K. K. and K. H. Wedepohl. 1961. Distribution of the Elements in
Some Major Units of the Earth's Crust. Bull. Geol. Soc. America. 72:175.
Vinogradov, A. P. 1959. The Geochemistry of Rare and Dispersed Chemical
Elements in Soils. Consultants Bureau, Inc.
Wahlberg, J. S. and M. J. Fishman. 1962. Adsorption of Cesium on Clay
Minerals. U.S.G.S. Bulletin 1140-A.
Wilding, M. W. and D. W. Rhodes. 1963. Removal of Radioisotopes from .
Solution by Earth Materials from Eastern Idaho. IDO-14624.
3-68
-------
COBALT
Natural Soil and Rock Distributions
The average cobalt concentrations found in igneous and sedimentary rocks
are listed in Table 3-34. As can be seen, the cobalt content of rocks steadily
decreases from the dark ultramafic rocks to the light-colored granitic rocks.
The ultrabasic rocks usually contain the bulk of the cobalt in olivines and
pyroxenes, but an occasional sulfide body in these rocks provides a workable
ore deposit (Young 1957). Cobalt normally is enriched in areas of low redox
potential in the environment.
TABLE 3-34. COBALT CONCENTRATIONS IN IGNEOUS AND SEDIMENTARY
ROCKS IN ppm (YOUNG, 1957)
Igneous Sedimentary
Granitic
Ultramafic Basaltic High Ca Low Ca Syenite Shale Sandstone Limestone
150 48 71 1 19 0.3 0.1
During weathering processes, cobalt is separated from manganese and
iron because its (III) oxidation state is normally unstable unless cotn-
plexed. However, certain bacteria are known to liberate chelated Co(III),
as stable amino acid complexes from the soil. Little is known, as yet, of
these reactions; but the-technique of predicting mineral solubilities for
cobalt based on thermodynamic data may not be entirely applicable if the
bacterial reactions are ignored (Moore, 1964). -Vinogradov (1959) reported
the average soils content of cobalt as 8 ppm. Cobalt contents of up to
1.96 wt% CoO have been reported for calcite (Hacquaert, 1925).
Brief Chemistry
eg
Only one stable cobalt nuclide, Co, is known to exist in nature.
Cobalt-60 is the principle cobalt activation product radionuclide. Cobalt
oxidation states include Co(II) and Co(III). Cobalt is geochemically similar
to the other ferromagnetic and closely related transition metals, iron and
nickel.
Sulfides of nickel and cobalt constitute the chief cobalt ore minerals.
, « o
Co ionic radius in octahedral coordination is 0.74 A, intermediate between
3-69
-------
Ni+2 (0.70 A) and Fe+2 (0.77 A) (Shannon and Prewitt, 1969). The relative
stabilities of the Co(II) and Co(III) states are greatly affected by complex-
ing in aqueous solutions (Latimer, 1952). Although Co cannot exist in
aqueous solution, complexing can stabilize the trivalent state which would
normally decompose water.
Solid Phase and Solution Equilibria
The sources for thermodynamic data for various solid phases were: Sillen
and Martell (1964), Co(OH)3, CoC03, CoOOH; Wagman et al. (1969), CoO, CoHP04;
Robie and Waldbaum (1968), CO^POJ^; Chase et al. (1975), Co,0.. The sta-
bility of cobalt solid phases depends on the pH and the oxidation-reduction
environment. For example, in acidic (pH 6) and oxidizing (p02 0.68) environ-
ments, the solids in increasing order of stability are: Co(OH)3, Co(OH)2
pink, CoHP04, Co3(P04)2, CoC03, Co304, and CoOOH as shown in Figure 3-7.
With the change in oxidation-reduction environment, all the minerals except
Co(OH)3, Co30», and CoOOH would stay in their relative positions. The latter
minerals would decrease in stability. With the change in oxidation-reduction
conditions from the extremely oxidizing condition to moderately well oxidized
or reducing conditions, (p02 > 20), CoC03 would be the most stable mineral.
Under extremely reducing conditions such as pO^SO, CoS would be the most
stable mineral. Cobalt sulfide (CoS) is the chief cobalt ore mineral. In
general, the stability of all the cobalt minerals increases with an increase
in pH.
2+ 3+
Cobalt exists as- Co and Co and forms solution complexes with common
2
soil anions (OH , Cl , S0|~, N03). The relative concentration of these solu-
tion species in equilibrium with CoCO^ and in an oxidizing environment (09 =
n CQ * £
10 ) is depicted in Figure 3-8. If a solid phase other than CoC03 or a
constant concentration of Co is chosen for the diagram, the lines will shift
up and down, while the relative concentration of various solution species
would stay the same. The thermodynamic data for CoS04 were obtained from
Sillen and Martell (1964) the hydrolysis species data were obtained from
Baes and Mesmer (1976). Data for all the other species were obtained from
Wagman et al. (1968, 1969).
3-70
-------
3456
Figure 3-7. The relative stability of cobalt solids in an oxidizing
soil environment CpOz(g) = 0.68 atm], pCa2+ = 2.5 and
phosphate levels in equilibrium with Variscite and
Gibbsite (V and G), Dicalcium Phosphate Oihydrate (DCPD)
and Octacalcium Phosphate (OCP).
-24
r
Figure 3-8. The activity of various cobalt ions in equilibrium with
CoC03(s) in an oxidizing soil environment Cp02(g) =0.68
atm], pC02(q) = 1.52 atm, pCl" = pS042 = 2.5, pN03 = 3.0
and
- = 5.0.
3-71
-------
2+
It can be seen from the figure that cobaltinous ion (Co ) and its solu-
tion complexes are the most stable and dominant solution species in the oxidiz-
ing environment. The most dominant solution species up to a pH value of
2+
approximately 9.5 is Co .
2+
Beyond pH 9.5, Co(OH)° becomes dominant. The activity of Co ion in
solution decreases 100-fold with an increase of a pH unit, whereas the acti-
vity of Co(OH)°2 does not change with the pH. These predictions are similar
to Baes and Mesmer (1976). If CoCOg was the solid phase that controlled the
solubility, the total cobalt concentration in solution at pH 7 would be
-7 -1?
approximately 10" M and approximately 10 M at pH 9.5.
Experimental Adsorption Results
Spencer and Gieseking (1954) studied cobalt adsorption on and release
from Wyoming bentonite and Swygert clay. Cobalt was adsorbed more strongly
than calcium. They suggested that cobalt was adsorbed on the clays as
monovalent hydroxy-cobaltous ions and as divalent cobaltous ions.
Haney (1957) described a test disposal of TBP (tributyl phosphate)
scavenged wastes produced by uranium recovery from the Bismuth Phosphate
Process. The Bismuth Phosphate Process recovered only weapons grade plutonium
from the irradiated, solubilized fuel elements. Rhodes and Nelson (1957)
referred to the same waste as Uranium Recovery Plant Waste and gave a cobalt
content of from 4 x 10"5 to 4 x 10 yc/cm of Co. Glaciofluviatile soil
column studies were undertaken and little Co was removed by adsorption on
the soil. The Co was reported to form a complex ion, probably a cobalti-
cyanide complex, that did not undergo ion exchange reactions with the soil.
The presence of nonexchangeable Co became a limiting factor in the disposal
of some wastes. Sorathesn et al. (1960) reported that Co behaved like a
colloid at pH 6 to 9, and gave a Co Kd value of 23,624 on ill He.
Friend (1963) reported that suspended solids in a pond environment took
up Co rapidly, with suspended illite having the most selectivity for cobalt
removal of the clay minerals. Wilding and Rhodes (1963) studied the effects
of 100 ppm citrate and EDTA on cobalt equilibrium distribution coefficients
with the zeolite clinoptilolite and Wei 1-81 sediments. At pH 6.5, citrate
enhanced cobalt adsorption on clinoptilolite, the Kd value being 48 with no
3-72
-------
treatment and 102 with 100 ppm citrate added. With EDTA addition at pH 6.5,
the .cobalt Kd value fell from 100 to 1 ml/g. Well-81 sediments adsorbed
about the same amounts with and without citrate. At pH 7.7, the cobalt Kd
was 30 ml/g with no citrate and 26 ml/g with citrate. At 7.7, the cobalt Kd
was 1000 ml/g with no EDTA and 1.2 ml/g with EDTA. Tiller et al. (1963)
found no correlation between cobalt adsorption and soil mineralogy. Hawkins
(1964) described a process of cobalt removal from waste solutions by scaveng-
ing on manganese dioxide precipitate and removal of the remaining cobalt by
ion exchange on lignite, soil and the natural zeolite clinoptilolite. Dis-
tribution coefficients for cobalt were 56, 800 and 24 ml/g for Idaho soil,
lignite and clinoptilolite, respectively.
Gonzalez and Gomez (1961) found that most of the cobalt in Andalusian
soils was found in oxide and hydroxide gels associated with clays. There was
a positive correlation between cobalt content of the soil and Feo^3 content.
Basu and Mukherjee (1965) obtained cobalt isotherms on montmorillonite
clay. Cobalt exchange behavior was very similar to that of potassium, ammo-
nium and nickel cations.
Humic acid was isolated from the A horizon of the Emory silt loam and
equilibrated-wfth Co and Co chloride solutions by Dunigan and Francis
(1972). Subsequent salt solution washings of the humic acid removed 91 to
99% of the 60Co.
James and Healy (1972) showed the adsorption isotherm for 1.2 x 10" M
+2
Co on silica at 25°C, as a function of pH. Adsorption is 10% at pH 6 and
greater than 90% at pH 8. Computed hydrolysis data at this cobalt concen-
tration are also shown.
+2
Tewari et al. (1972) investigated the adsorption of Co by Fe^O^, A^O-j
and Mn02 as a function of concentration, solution pH and temperature in con-
nection with the transport of cobalt adsorbed on corrosion product oxides.
+2
Hydrolysis of Co was suggested as a possible mechanism for the marked
dependence of cobalt adsorption on pH and temperature.
3-73
-------
Migration Results
Field Studies—
Haney (1S67) reported the disposition of radioactivity beneath several
Hanford ground disposal sites, among them the 216-BY covered trenches. The
eight BY trenches were used from December 1954 to December 1955 and received
3.4 x 10 liters of radioactive U-plant, high-salt scavenged waste (O.IM^
POA3", total dissolved solids = 350 g/liter, mostly NaNO,, pH = 9.5, 60Co =
-3 -53 *
4 x 10" to 4 x 10~ yCi/cm ). The waste contained enough phosphorus (some
of it as degraded tributyl phosphate) to complex most of the cobalt in a neu-
tral to anionic species that essentially passed through the soil column
without reaction. The groundwater beneath the trenches was contaminated with
1 x 10 uCi/cm of Co. The Co on the soil column (68 m to the water
table) was uniformly distributed at concentrations of from 1 x 10" uCi/g of
-4
soil to 1 x 10 uCi/g of soil.
Magno et al. (1970) reported that Co constituted only <0.1% of the
beta activity in surface drainage from the Nuclear Fuel Services plant Lagoon
system in western New York State. Even that small amount of Co constituted
40% of the Co entering the system of lagoons because of the small amount of
Co disposed to the lagoon. The Co in the effluent was relatively small,
but was all in solution and thus able to travel with the surface water. The
implication is that the "soluble" Co was migrating as a complexed species.
Duguid (1976) studied the seepage waters from burial grounds at Oak Ridge
used for the disposal of intermediate-level liquid wastes from 1962 to 1965.
Trench 7 seepage contained Co combined in an organic complex that was not
readily adsorbed by the surrounding shale. Groundwater, as well as surface
, contained these ree
Laboratory Studies—
The chemical mechanism of Co transport in groundwater from intermediate-
level liquid waste Trench 7, Oak Ridge, Tennessee, was studied by Means et al.
(1976). A small volume groundwater seep 50 m east of Trench 7 contained Co
in soil concentrations of 10 to 10 dpm/g and water concentrations of
10 dpm/ml. The Co was transported as a complex with molecular weight less
than 700 according to Sephedex gel G-10 studies. The G-10 is designed to
3-74 .
water, contained these readily-transported Co complexes.
-------
retard the passage of organic solutes with molecular weights less than 700.
The implication is that the Co is complexed by low molecular weight natural
organics. Migration of Co would be negligible in the absence of organic
complexation because the soil adsorption capacity for inorganic forms of Co
was reported to be extremely high. Manganese oxides in the Conasauga shale
were the principal adsorbents of ionic and weakly complexed b Co.
Saas and Grauby (1973) studied the migration of cobalt and its transfer
mechanisms from nuclear reactor to river water, irrigation water and finally
groundwater. Three fractions of Co were differentiated including a chelated
fraction, an exchangeable fraction and a hydrosoluble fraction. The hydro-
soluble fraction was defined as the totality of the organic and mineral com-
ponents of the soil that are water soluble. The different fractions in a
brown, calcareous soil were given. The hydrosoluble fraction constituted
about 70%, the exchangeable fraction about 20% and about 10% was chelated
from river water containing Co. The hydrosoluble form on an alluvial
calcareous soil column varied from 10% at the point of input to 60% at 17 cm
depth. The water soluble material was only slightly biodegradeable, even-
tually migrating to the groundwater through application in irrigation water.
The pollutants in the' river water,, including industrial wastes and mun-icipal
sewage, are responsible for the rapidly migrating hydrosoluble fraction. The
authors suggest that the quality of surface drainage water will have to be
taken into account during nuclear reaction siting.
Summary
Among the various soil components, FegO^ (Gonzalez and Gomez, 1964) and
illite (Friend, 1963; Sorathesn et a!., 1960) have been shown to have a strong
affinity for cobalt. The adsorption of cobalt by soil minerals have been
attributed to mainly an ion exchange phenomenon (Hawkins, 1964; Basu and
Mukherjee, 1965). With the increase in pH (especially 6 to 9) the adsorption
of cobalt increases (James and Healy, 1972; Tewari, Campbell and Lee, 1972).
This increase in adsorption with pH has generally been explained as due to
the formation of cobalt colloids (Sorathesn et al., 1960) and hydrolysis of
?+
Co (Tewari et al., 1972). The presence of various complexing ions such as
EDTA, cyanide, and fulvic acids have been shown to reduce the adsorption by
soil minerals (Wilding and Rhodes, 1963; Rhodes and Nelson, 1957; Dunigan and
3-75
-------
Francis, 1972). These laboratory results are consistent with the theoretical
predictions (Figure 3-8) which indicate that in acidic pH values cobalt would
2+
predominantly exist as Co and with the increase in pH it hydrolyzes to
CoOH* and Co(OH)|.
Discrete cobalt minerals have not been identified in soils. High concen-
trations of cobalt are required to precipitate cobalt minerals presented in
Figure 3-4 under acidic pH values. It is not likely that any of the minerals
would exist in the acidic soil environment. However, in highly alkaline con-
ditions, CoCOg would maintain fairly low concentrations of Co in solution and
thus could exist in the soil environment.
Cobalt is apparently relatively easily complexed by natural organics
(Duguid, 1976; Means et al., 1976) and by synthetically-produced organics
(Haney, 1967; Saas and Grauby, 1973)', and as such is able to migrate with
relative ease through the soil column with which it would normally react
by adsorption.
References
Baes, C. F., Jr. and R. E. Mesmer. 1976. The Hydrolysis of Cations. John
Wiley and Sons, New York.
Basu, A. N. and S. K. Mukherjee. 1965. Interaction Between Montmorillonite
Clay and Trace Element Cations. II. Exchange Behavior of Cobalt, Nickel and
Chromium Ions in Clays. J. Indian Soc. Soil Sci. 13:251-256.
Chase, M. W., J. L. Curnutt, H. Prophet, R. A. McDonald and A. N. Syverud.
1975. JANAF Thermochemical Tables, 1975 Supplement. Journ. Phys. and Chem.
Ref. Data 1974. (1)4.
Duguid, J.. 0. 1976. Annual Progress Report of Burial Ground Studies at Oak
Ridge National Laboratory: Period Ending September 30, 1975. ORNL-5141.
Dunigan, E. P. and C. W. Francis. 1972. Adsorption and Desorption of
Cobalt 60, Strontium 85, and Cesium 137 on Soil Humic Acid. Soil Science.
114:494-496.
Friend, A. G. 1963. The Aqueous Behavior of 85Sr, 137Cs, 65Zn, and 60Co as
Determined by Laboratory Type Studies. .IN.: Transport of Radionuclides in
Fresh Water Systems. TID-7664.
Gonzalez, G. F. and A. M. Gomez. 1961. Geochemistry of Cobalt in Soils of
Western Andalusia. IV. Cobalt Clays and Correlation Between Cobalt and
Iron, Clay and Manganese Contents. An. Edafol. Agroboil. 23:563-572.
3-76
-------
Hacquaert, A. L. 1925. Pink Calcite from Tantara (Katanga). Natuurw. Tijds.
Antwerp, vol. 7, p. 102.
Haney, W. A. 1957. Disposal of High Cobalt-60 Scavenged Wastes. HW-48862.
Haney, W. A. 1967. Final Report on the Effects of Ben Franklin Dam on
Hanford. BNWL-412, pp. 8-13.
Hawkins, D. B. 1964. Removal of Cobalt and Chromium by Precipitation and
Ion Exchange on Soil, Lignite and Clinoptilolite from Citrate-Containing
Radioactive Liquid Waste. IDO-12036.
James, R. 0. and T. W. Healy. 1972. Adsorption of Hydrolyzable Metal Ions
at the Oxide-Water Interface. J. Colloid Interfac. Sci. 40:65-81.
Latimer, W. M. 1952. The Oxidation States of the Elements and Their Potentials
in Aqueous Solutions. Prentice Hall, Inc.
Magno, P., T. Reavey, and J. Apidianakis. 1970. Liquid Waste Effluents'from
a Nuclear Fuel Reprocessing Plant. BRH-NERHL-70-2.
Means, J. L., D. A. Crerar, and J. 0. Duguid. 1976. Chemical Mechanisms of
60co Transport in Ground Water from Intermediate - Level Liquid Waste Trench 7:
Progress Report for Period Ending June 30, 1975. ORNL/TM-5348.
Moore, L. R. 1964. The Microbiology, Mineralogy and Genesis of a Tonstein.
Prac. Yorkshire Gebl. Soc. 34:235-292.
Rhodes, D. W-. and J. L. Nelson. 1957. Disposal'of Radioactive L-iquid Wastes
from the Uranium Recovery Plant. HW-54721.
Robie, R. A., and D. R. Waldbaum. 1968. Thermodynamic Properties of Minerals
and Related Substances at 298.15°K (25°C) and One Atmosphere (1.013 bars)
Pressure and at Higher Temperatures. Geol. Surv. Bull. 1259.
Saas, A. and A. Grauby. 1973. Mechanisms for the Transfer to Cultivated Soils
of Radionuclides Discharged by Nuclear Power Stations into the System: River-
Irrigated Soil-Ground Water. IN; Environmental Behavior of Radionuclides
Released in the Nuclear Industry. IAEA-SM-172/57, pp. 255-269:
Shannon, R. D. and C. T. Prewitt. 1969. Effective Ionic Radii in Oxides
and Fluorides. Acta. Cryst., B. 25:925.
Sillen, L. G. and A. E. Martell. 1964. Stability Constants of Metal-Ion
Complexes. Special Publication No. 17. The Chemical Society, London.
Sorathesn, A. G. Bruscia, T. Tamura, and E. G. Struxness. 1960. Mineral and
Sediment Affinity for Radionuclides. CF-60-6-93.
Spencer, W. F., and J. E. Gieseking. 1954. Cobalt Adsorption and Release
in Cation Exchange Systems. Soil Sci. 78:267-276.
Tewari, P. H., A. B. Campbell, and W. Lee. 1972. Adsorption of Cobalt (+2)
by Oxides from Aqueous Solution. Can. J. Chem. 50:1642-1648.
3-77
-------
Tiller, K. G., J. F. Hodgson, and M. Peech. 1963. Specific Sorption of
Cobalt by Soil Clays. Soil Science. 95:392-399.
Vinogradov, A. P. 1959. The Geochemistry of Rare and Dispersed Chemical
Elements in Soils. Consultant Bureau.
Wagman, 0. D., W. H. Evans, V. B. Parker, I. Halow, S. M. Bailey, and
R. H. Schumm. 1968. Selected Values of Chemical Thermodynamic Properties.
Tables for the First Thirty-Four Elements in the Standard Order of Arrange-
ment. MBS Technical Note 270-3.
Wagman, D. D., U. H. Evans, V. B. Parker, I. Halow, S. M. Bailey and
R. H. Schumm. 1969. Selected Values of Thermodynamic Properties. Tables
for Elements 35 Through 53 in the Standard Order of Arrangement. MBS Tech-
nical Note 270-4.
Wilding, M. W. and 0. W. Rhodes. 1963. Removal of Radioisotopes from Solution
by Earth Materials from Eastern Idaho. IDO-14624.
Young, R. S. 1957. The Geochemistry of Cobalt. Geochim. et Cosmochim.
Acta. 13:28-41.
CURIUM
Natural Soil and Rock Distributions
The element curium has not been reported to occur naturally in soil or
rocks.
Brief Chemistry
238 250
Thirteen isotopes of curium are known from Cm to Cm, most with
relatively short half-lives. There are no naturally-occurring curium isotopes.
3-78
-------
239
The curium isotopes can be produced in reactors by n reactions on Pu or by
decay of californium isotopes (Keller, 1971). The LWR reprocessing waste iso-
topes are given in Table 3-35. Curium-250 also has a long half-life, but
yields are normally very low. Curium can exist in Cm(III) or Cm(IV) oxidation
states, but stabilization of Cm(IV) requires high concentrations of fluoride
_2
ions to form the anionic complexes CmFg and CmF ^ (Keller, 1971). Only tri-
valent curium is normally stable in aqueous solutions. The ionic radius of
+3 °
Cm in coordination number 6 is 0.986 A. Trivalent curium forms many com-
plexes with inorganic ligands found in wastes and in groundwaters.
TABLE 3-35. CURIUM RADIONUCLIDE DATA
(WEAST, 1976)
Decay Mode
a
a
a
a
a
a
Solid Phase and Solution Equilibria
Curium forms hydride, hydroxide, halides, oxides, and organometallic com-
pounds (Keller, 1971). Except for CmF^, a search for thermodynamic data for
these compounds was unsuccessful. Therefore, predictions regarding the sta-
bility of formation of various compounds of curium in soils cannot be made
at this time.
Curium is found in two oxidation states [Cm(III) and Cm(IV)] in aqueous
solution (Keller, 1971). Curium (IV) is not stable in solutions because of
self-rad.iation and rapid changes to Cm(III). Only trivalent curium is stable.
in aqueous solutions. Keller (1971) report's that the behavior of Cm in
solution is similar to lanthanide elements. Most of the lanthanides form
-19
trivalent hydroxides in natural waters with solubility products of <10
(Vickery, 1953). Thomas and Jacobs (1969) found that the activity of curium
in 0.1M NaCl solutions which contained approximately 10 M/l of curium
3-79
Isotope
242Cm
243Cm
.244Cm
245Cm
246Cm
247Cm '
Half-Life
163 days
32 years
17.6 years
9,300 years
5,500 years
470,000 years -
-------
decreased at pH values above 3. They attributed this decrease in activity in
pure solutions to either precipitation or adsorption on walls. If this is
assumed to be due to precipitation, the solubility product of curium hydroxide
would be at least as low as the lanthanides, and probably much lower.
Since thermodynamic data for the curium compounds is not available, it
was decided to plot the solution species in equilibrium with Cm(OH)3 with
an assumed log K° of Z for the dissolution of Cm(OH), [Cm(OH), + 3H+ £ Cm3+
+ 3H20] in order to show the relative activity of various solution species
of curium with the changes in pH. Keller (1971) reviewed the curium data
and stated that little work has been performed on the complex chemistry of
curium and since a wide variation occurs in stability constants determined
by various authors, these values should be regarded as relative rather than
absolute. The thermodynamic data used to develop Figure 3-9 were obtained
from the following sources: Trotman-Dickinson (1973), CmCl2"1", CmF2+, CmSO/;
Jones and Choppin (1969), CmN032+; Shalinets and Stepanov (1972), CmOH2+,
Cm(OH)2+ and Moskin (1969), CmH2P042+.
z-io
Z-12
z-w
Z-16
Z-18
Z-20
1-22
6 7
pH
10
Figure 3-9.
The activity of various curium ion species in equilibrium
with Cm(OH)3(s) with pF" = 4.5, pCl = pSO*2' = 2.5,
pNO^ =3.0 and pH0P07 = 5.0.
3-80
-------
It can be seen from Figure 3-9 that curium forms various solution com-
plexes with OH", F", Cl", NO^, SO*", and P0^~. The activity of all the com-
plexes decreases with the increase in pH. The solution complexes of P0?~,
Cl", NOg, do not contribute significantly to the total curium concentration
in solutions. The most dominant solution species in pH ranges of <3.3, 3.3-
5.7, 5.7 and up are Cm + , Cm(OH) , and Cm(OH)t, respectively.
Experimental Adsorption Results
242
The adsorption characteristics of Cm on several clays was measured in
distilled water and 0.1M NaCl by Thomas and Jacobs (1969). Adsorption was not
a function of ion exchange capacity because greater adsorption occurred in
0.1M NaCl solutions (Jacobs et al., 1966) than in distilled water solutions.
When adsorption studies were made as a function of pH, some curium was removed
from solution even without the addition of clay at pH values above 3. This
was due to precipitation or to adsorption on glass walls. Addition of clay
caused a considerable increase in adsorption possibly because increased surface
242
area was available for deposition. Column loading studies of Cm tagged
0.1M NaCl showed normal ion exchange chromatographic breakthrough at pH 1.
Above pH 1, the effluent deviated from normal chromatographic curves and
242
remained constant during the entire run. The effluent/influent Cm concen-
tration ratio at pH 3 was about 0.008; at pH 7 the ratio was 0.028; and at
242
pH 10 the ratio was 0.150. Although Cm precipitated from stock solution
at pH values above 3, the column leakage of curium increased with, increasing
pH. The leakage is probably caused by hydroxy and oxo-complexes forming
radiocolloids. Movement of curium through soil is probably restricted by
filtration and surface adsorption of radiocolloids rather than by ion exchange.
244
Sheppard et al. (1976) investigated the distribution values of Cm on
several Washington and South Carolina soils. The 50 day distribution results
for trace Cm in a distilled water solution are given in Table 3-36. The
distribution for curium between soil and solution increased with time. The
distribution values were computed in the same manner as a Kd value. However,
Kd values are defined as obtained at soil-solution equilibrium which these
values did not attain. The slow increase in distribution values were inter-
preted by the authors as due to the presence of colloids. It was also shown
that iron and manganese oxides were effective scavengers for curium.
3-81
-------
TABLE 3-36. CURIUM 50-DAY DISTRIBUTION COMPUTED
FROM SHEPPARD ET AL., (1976)
Distribution, Monthly Change in
Soil Identity ml/g Distribution
Muscatine 1330 +8%
Silt Loam
Burbank 106 +27%
Loamy Sand
Ritzville 704 +40%
Silt Loam
Fuquay Sand 1850 +26%
0-5 cm Depth
Fuquay Sand 1850 +28%
5-15 cm Depth
Fuquay Sand 1240 +18%
15-50 cm Depth
Migration Results
Field Studies—
Duguid (1976) reported that Bondietti had shown that the alpha contamina-
tion in-water from seep S-4 along the south'side of burial ground 5, Oak Ridge,
Tenr
238r
Tennessee, was due to 3.2 x 10"7 yCi/ml of 244Cm and 3.2 x 10~8 yCi/ml of
Pu. The burial trenches in burial ground 5 were overflowing because pre-
cipitation infiltrated the trenches, reached the less permeable trench bottom
244 238
and flowed out the lower end of the trench in a seep. The Cm and Pu
were presumably picked up by the water in its passage through the waste-filled
burial trench. Bondietti and Reynolds (1977) reported that seepage water,
244
presumably from burial ground 5, contained 700 ± 70 dpm/1 of Cm in the
filtrate from a 0.45 ym filter. •
Laboratory Studies—
Nishita et al. (1976) determined the extractability of plutonium and
curium from a contaminated soil as a function of pH and soil components. The
organic matter and pH influences hydrolysis, precipitation, coprecipitation
and adsorption of the radionuclides. In the work, 2 g of soil were suspended
in 25 ml of extracting water, in duplicate, and the pH adjusted with NaOH or
HN03. The Kd values computed from Nishita et al. (1976) are given in Table 3-37
3-82
-------
for the untreated Aiken clay loam, a kaolinitic soil, at several pH values.
The Aiken clay loam also was treated sequentially with progressive treatments
to remove salts, organic matter, carbonates, manganese oxides, ion oxides,
free silica and alumina and amorphous aluminosilicates. The organic matter
was shown to have the greatest ability to retain curium, but this was a func-
tion of pH as well. Below pH 4.5, curium was less strongly adsorbed whether
organic matter was present or not.
TABLE 3-37. Kd VALUES FOR 242Cm IN UNTREATED AIKEN
CLAY LOAM (NISHITA ET AL., 1976)
242,
Kd, ml/g pH fc^Cm Kd. ml/g
1.21
2.12
2.56
4.69
85.7 ± 12
2,457 ± 6
6,803 ± 46
100,000 ± 30,000
8.54
9.43
10.31
11.25
4608 ± 1635
1776 ± 142
690 ± 38
358 ± 2.7
5.55 100,000 ± 10,000 12.22 190 ± 3.3
7.08 71,429 ± 10,204 13.25 272 ± 23
Summary
No information is available on possible curium compounds in soils and
sediments. It appears that curium precipitation as Cm(OH)., in alkaline solu-
tions may control curium concentrations in those solutions. Theoretical
calculations (Figure 3-9) indicate that Cm hydrolyzes even in acidic solu-
tions (pH 3.3 and above). Very little reliable information is available on
interactions of curium with soils and sediments, and none on curium interac-
tions with rocks. What little data are available indicate that curium adsorp-
tion is not a function of ion exchange capacity of the soil (Thomas and
Jacobs, 1969; Jacobs et a!., 1966), and that precipitation and/or formation
of radiocolloids may predominately influence curium adsorption reactions
(Thomas and Jacobs, 1969; Sheppard et al., 1976).
3-83
-------
References
Bondietti, E. A. and S. A. Reynolds. 1977. Field and Laboratory Observations
on Plutonium Oxidation States. BNWL-2117, pp. 505-538.
Duguid, J. 0. 1976. Annual Progress Report of Burial Ground Studies at Oak
Ridge National Laboratory: Period Ending September 30, 1975. ORNL-5H1.
Jacobs, D. G., Y. E. Kim, and 0. M. Sealand. 1966. Application of Mineral
Exchange to Reactor Technology. ORNL-4007, pp. 27-33.
Jones, A. D. and G. R. Choppin. 1969. Complexes of Actinide Ions in Aqueous
Solutions. Actinide Reviews. 1:311-336.
Keller, C. 1971. The Chemistry of the Transuranium Elements. Vol. 3.
Kernchemie in Einzeldarstellungen. Verlag Chemie GmbH.
Moskin, A. I. 1969. Complex Formation of the Actinides with Anions of Acids
in Aqueous Solutions. Soviet Radiochemistry. 11:447.
Nishita, H., M. Hmailton and A. J. Steen. 1976. Extractability of 238Pu and
Cm from a Contaminated Soil as a Function of pH and Certain Soil Components.
HNOs - NaOH System. Presented at Annual Meeting of Soil Science Society of
America, Houston, Texas.
Shalinets, A. 8. and A. V. Stepanov. 1972. Investigation of Complex Forma-
tion of the Trivalent Actinide and Lanthanide Elements by the Method of Electro-
migration. XVII.- Hydrolysis. Radiokhimiya. 14:280-282.
•
Sheppard, J. C., J. A. Kittrick, and T. L. Hart. 1976. Determination of
Distribution Ratios and Diffusion Coefficients of Neptunium, Americium and
Curium in Soil-Aquatic Environments. RLO-2221-T-12-2.
Thomas, W. A. and D. G. Jacobs. 1969. Curium Behavior in Plants and Soil.
Soil Science. 108:305-307.
Trotman-Dickinson, A. F. (Executive Ed.). 1973. Comprehensive Inorganic
Chemistry. Pergamon Press, New York.
Vickery, R. C. 1953. Chemistry of the Lanthanons. Butterworths Sci. Pub!.,
London.
Weast, R. C., Editor. 1976. Handbook of Chemistry and Physics. The Chemi-
cal Rubber Co., Cleveland, Ohio, pp. B349-B350.
3-84
-------
EUROPIUM
Natural Soil and Rock Distributions
The europium content of several common rock types are given in Tables 3-38
and 3-39.
TABLE 3-38. EUROPIUM CONTENT OF IGNEOUS ROCKS IN ppm
(HASKIN AND SCHMITT, 1967)
Low K Hawaiian Nepheline
Peridotlte Tholeiite Tholeiite Andesite Granodiorite Leucogranite Syenite
0.2 1.9 1.3 1.0 1.2 0.17 10
TABLE 3-39. EUROPIUM CONTENT OF SEDIMENTARY ROCKS IN ppm
(HASKIN ET AL., 1966)
Shale Quartzite Limestone Subgraywacke
2.0 0.09 0.14 1.3
Haskin and Schmitt (1967) gave the value of 2.9 ppm Eu in a composite
of 40 North American shales, with a variation in sedimentary rocks of about
five times that value.
In a study of rare earths in Russian platform soils, Balashov et al.
(1964) reported high europium contents in alkaline soils, the europium precipi-
tating as hydroxide. Acidic soils prob.ably allow complexing and removal of
the europium and other rare earths.
Brief Chemistry
Europium occurs in two stable isotopes out of a total of 17 isotopes,
Eu (47.8%) and Eu (52.2%). The radionuclides of interest are given
in Table 3-40. The oxidation states of europium are trivalent with
o o
radius 0.95 A and divalent with radius 1.09 A (Ahrens, 1952). The
+2
divalent state is typical of europium occurrence, substituting for Ca
and K in main stage feldspars (plagioclases, microcline, orthoclase). The
potential of the Eu(III)/Eu(II) couple is -0.43 volts (Latimer, 1952).
TABLE 3-40. EUROPIUM RADIONUCLIDE DATA (WEAST, 1976)
Decay Mode
6*, EC, 8"
Isotope
152Eu
154Eu
155Eu
156Eu
Half-Life
13.6 years
3.6 years
1 .81 years
15 days
3-85
-------
Europous ions are stable in water solutions at low hydrogen ion concentrations
.(basic pH), and can be formed from Eu by zinc reduction. However, the
europium in soil solutions is more apt to be in the trivalent than divalent
state.
Solid Phase and Solution Equilibria
Europium forms oxides, hydroxides, and salts with chlorides and sulfates
(EuO, Eu203, Eu304, Eu(OH)3, EuCl3'6H20, and Eu2(S04)3'SH20). Thermodynamic
data for these solid phases were used to plot the stability of the minerals
(Figure 3-10). It can be seen that, except for the hydroxides [Eu(OH)3], all
the other minerals considered are too soluble in environmental pH ranges of
interest (pH 4-10). In Figure 3-10, Eu(OH)3> Eu203, Eu2(S04)38H20 are the
only solids shown. The thermodynamic data for Eu203 and Eu(OH)3 were obtained
from Shumm et al. (1973), and for Eu2(S04)3'8H20 were obtained from Latimer
(1952). The other solids were too soluble and fall beyond the graph bounda-
ries. Of all the solids considered in Figure 3-10, Eu(OH)3 is the most stable
in alkaline environments. The results of Serne and Rai (1976) indicate that
kinetics of Eu(OH), precipitation is very fast and that Eu(OH), is the most
3+
likely europium solid to maintain the activity of Eu in solution in slightly
acidic to alkaline conditions. However, where Eu(OH)3 does not precipitate
or the europium concentration is below the Eu(OH)3 solubility line, cation
exchange reactions are the probable europium adsorption mechanism.
2+ 3+ 3+
Europium can exist as Eu and Eu . However, Eu is the only state
which is stable in water (Baes and Mesmer, 1976). The activity of various
solution species of europium in equilibrium with Eu(OH)3 and at assumed con-
centrations of various anions commonly found in soils is given in Figure 3-11.
2+ 2+
The thermodynamic data of all the species except EuOH and EuF were obtained
2+ 2+
from Schumm et al. (1973). The data for EuOH and EuF was obtained from
Baes and Mesmer (1976) and Walker and Choppin (1967), respectively.
In a general way the solution complexes of europium in decreasing order
?+ + "4" ?+
of importance can be arranged as follows: Eu,P907 , EuSO,, Eu(SO.):,, EuF ,
2+ 2+
EuCl , and EuNOt . Under the conditions assumed in Figure 3-11, the pre-
dominant solution species in pH ranges of <4.5, 4.5-7.75, 7.5-8.85, and >8.35
will be EuSoJ, Eu^Oy"1", EuSOj, and EuOH , respectively. Thus, significant
3-86
-------
-4
-6
-8
-10
-12
6 7
PH
10
Figure 3-10. The relative stability of various europium
solids at pSQ2- =2.5.
-2
ff -6
-S
-ID
-12
EuOH
,2+
10 11
pH
Figure 3-11.
The activity of various europium species in equilibrium
with Eu(OH)3(s). Other soil solution conditions included
pS042- = pCl" = 2.5, pN03~ = 3.0, pF" = 4.5 and phosphate
levels were from Variscite and Gibbsite (V and G),
Dicalcium Phosphate Dihydrate (DCPD) and Octacalcium
Phosphate (OCP).
3-87
-------
quantities of uncomplexed europium (Eu ) can only be expected in soils with
very low levels of sulfate and phosphate. The solution complexes of europium
with fluoride, chloride, and nitrate do not contribute significantly to total
europium concentration in solution and hence can be ignored.
Experimental Adsorption Results
Bensen (1960) reported that the uptake of rare earths (cerium, promethium
and europium) was virtually complete and unaffected by accompanying 0.5M
alkali metal and 0.25M alkaline earth metal chlorides at a pH above 7. Below
pH 3, however, the adsorption of all rare earths was depressed similarly by
the accompanying salts. These results were interpreted to mean that the rare
earths, including europium, are ionic at pH below 3 and that they are taken
up principally by ion exchange on the soil. Above pH 7, they are precipitated.
The rare earths, including europium, appear to exist as tripositive ions.
Baetsle and Dejonghe (1962) gave a europium Kd range in Mol soil (mostly
quartz sand) at pH 7.7 as 228 to 705, and reported that at pH >3, europium
hydrolyzed. Therefore, above pH 3, the authors did not consider that adsorp-
tion was an ion exchange phenomenon. However, a plot of pH versus log Eu
activity (Figure 3-11) does not support the above conclusion.
Baetsle et al. (1964) gave Kd values for trace europium in three differ-
ent waters. The compositions of these waters were not given. Kd results are
shown in Table 3-41. The europium Kd values are much larger than the cesium
and strontium Kd values for the same soils.
TABLE 3-41. VARIATION OF TRACE EUROPIUM Kd VALUES '
ON MOL SOILS (BAETSLE ET AL., 1964)
Kd, ml/g
Soil Type
Eolian Sand
Horizon A
Horizon B
Horizon C
Mol White Sand
Mol Lignitic Sand
Deionized
Water, pH 3
Tap Water,
pH 7.7
214-244
61-75
69-89
_—
705
430
384
801
Groundwater,
pH 3
130-380
165-315
27-1830
-------
Serne and Rai (1976) studied the adsorption-precipitation behavior of
europium in soils and pure solutions. The pH of the various europium solu-
tions (0.15M CaCl2 plus trace to 50 mg Eu/1) was increased in 0.05 pH incre- .
ments until precipitation was observed by the Tyndall Beam method. At the
point of Eu(OH)3 precipitation, the europium activity and pH was determined.
The results are plotted in Figure 3-12 along with the theoretical Eu(OH)3
solubility line. The results indicate that:
1) the kinetics of precipitation of Eu(OH)3 is rapid,
2)
3)
the laboratory results agree well with the theoretical Eu(OH)-, solubility,
3+ - J
Eu does not appear to hydrolyze appreciably because total europium
3+
activity in equilibrium with Eu(OH)3 is similar to the calculated Eu
activity, and
4) for meaningful Kd values, the theoretical solubility of Eu(OH)3 should
not be exceeded.
^
LU
s
-2
-3
-4
-5
-6
-7
-8
\
\
e\
I \
• \
•\
\
\
\
\
\
\
\
\
\
\
\
V-
\Eu(OH)3
\
\
• \
3456 789
pH
Figure 3-12.
The influence of pH on the activity of europium
in solution (Serne and Rai, 1976)
3-89
-------
Europium adsorption experiments where Eu(OH)3 did not precipitate were
conducted, as shown in Table 3-42. The results show that the europium Kd
increases with an increase in pH. and decreases with increasing europium con-
centration, below the point of Eu(OH)3 precipitation. These reactions sug-
gest that ion exchange is an active mechanism beTow the point of Eu(OH)3
precipitation.
TABLE 3-42. EFFECT OF pH AND EUROPIUM CONCENTRATION
ON EUROPIUM Kd BY BURBANK SAND*
(SERNE AND RAI, 1976)
Eu, ppm**
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
0.05
Final
JEiL
4.82
4.82
4.90
5.00
5.05
5.05
5.08
5.10
5.10
5.10
5.10
5.10
5.15
5.15
5.15
5.15
5.20
5.20
5.20
5.25
5.30
5.30
5.50
4.75
4.75
4.80
5.00
5.00
5.00
5.08
5.08
5.10
5.25
Kd,
mg/1
46
48
51
75
70
69
75
84
78
78
82
75
93
90
81
93
94
98
98
107
106
97
153
32
38
38
63
53
53
56
62
67
71
Eu, ppm**
0.05
0.05
0.05
0.05
0.05
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
5.0
Final
nW
5.25
5.30
5.30
5.30
5.35
4.75
4.80
4.80
5.00
5.05
5.05
5.10
5.10
5.10
5.15
5.15
5.20
5.20
5,20
5.25
5.25
5.30
5.35
4.80
4.80
4.80
4.95
5.00
5.05
5.05
5.10
5.10
5.15
Kd,
mg/1
73
83
82
84
98
28
26
27
35
35
41
42
41
42
• 43
45
53
48
45
48
57
54
60
12.4
11.3
12.4
16.9
18.4
15.7
14.8
15.9
18.7
16.4
Eu, ppm**
5.0
5.0
5.0
5.0
50.0
50.0
50.0
50.0
50.0
50.0
50.0
50.0
50.0
50.0 .
50.0
50.0
50.0
50.0
50.0
Final
pH
5.20
5.20
5.25
5.32
4.70
4.70
4.73
4.88
4.90
4.90
4.90
4.90
4.90
5.00
5.00
5.00
5.15
5.15
5.20
Kd,
rnq/1
17.7
15.9
17.0
20.0
6.9
6.9
6.5
5.8
5.8
5.9
6.3
6.2
6.2
7.2
7.0
7.1
8.2
8.2
7.9
* One gram samples free of CaC03 and soluble salts were used
** Various concentrations of Eu were added to 0.15M CaCl2 solu-
tions spiked with Eul52. Samples were equilibrated first
for 48 hr and then for 24 hr after adjusting the pH.
3-90
-------
Migration Results
Field Studies--
Brookins (1976) reported that europium was retained in the shale that
enclosed the fossil nuclear reactor zones at the Oklo mine in Gabon along
with the other rare earths produced there by the nuclear reactions 1.8 bil-
lion years ago.
Laboratory Studies—
Baetsle et al. (1964) determined the K values associated with europium
movement in relation to groundwater movement is given in Table 3-43. These
data also may be compared to strontium and cesium K values for the same Mol
soils. The europium K values are very pH-sensitive, as the data of Fig-
ure 3-12 indicate.
TABLE 3-43. K VALUES OF EUROPIUM IN THE SATURATED
SUBSOILS AT MOL, BELGIUM (BAETSLE
ET AL., 1964) K IS DEFINED ON PAGE 2-30.
Soil Type pH 4 pH .3
"MoFsand 100 58
Perturbed Profile 1963 511
Mol Sand & Eolian Sand — 280
Mol Sand & Lignite — 997
Mol Sand 2472 225
Mol Sand 640 139
Mol. Sand _ 1282 78
Mol Sand 10600 74
Mol Sand 2145 53
Mol Sand 697 64
Mol sand is nearly pure quartz.
Summary
No information on possible europium compounds in soils and sediments is
available. Seme and Rai (1976) have shown that the kinetics of precipitation
3-91
-------
of Eu(OH)- from pure solutions is rapid. The concentration of europium in
equilibrium with Eu(OH)g decreases 1000-fold with an increase of one pH unit.
The activity of europium in equilibrium with Eu(OH)., is low under highly
alkaline conditions and Eu(QH)., may control europium concentration in alkaline
soils and sediments.
Bensen (1960) and Baetsle and Dejonghe (1962) reported that europium
appears to be taken up principally by ion exchange below pH 3. Baetsle and
Dejonghe (1962) suggested that ion exchange was not the adsorption mechanism
above pH 3 because of europium hydrolysis. However, data presented in Fig-
ure 3-11 and the results of Serne and Rai (1976) are at variance with the
above conclusion. Serne and Rai (1976) have shown that for meaningful
determination of Kd values in soils, one must be below the solubility line of
Eu(OH)g. Their results, thus obtained, showed that between pH 4.7 and 5.5
the Kd decreases with an increase in pH and Kd decreases as europium concen-
tration increases, which suggests an ion exchange as an adsorption mechanism.
References
Ahrens, L. H. 1952. The Use of lonization Potentials. Part I. Ionic Radii
of the Elements. Geochim. et Cosmochim. Acta. 2:155.
Baes, C. F., Jr. and R. E. Mesmer. 1976. The Hydrolysis of Cations. John
Wiley and Sons, New York.
Baetsle, L. and P. Dejonghe. 1962. Investigations on the Movement of
Radioactive Substances in the Ground. Part III: Practical Aspects of the
Program and Physiocochemical Considerations. IN: Ground Disposal of Radio-
active Wastes. TID-7628, pp. 198-210.
Baetsle, L. H., P. Dejonghe, W. Maes, E. S. Simpson, J. Souffriau, and
P. Staner. 1964. Underground Radionculide Movement. EURAEC-703.
Balashov, U. A., et al. 1964. The Effects of Climate and Facies Environment
on the Fractionation of the Rare Earths During Sedimentation. Geochemistry
International. 10:951-969.
Bensen, D. W. 1960. Review of Soil Chemistry Research at Hanford. HW-67201.
Brookins, D. G. 1976. Shale as a Repository for Radioactive Waste: The
Evidence from Oklo. Environmental Geology. 1:255-259.
Haskin, L. A. and R. A. Schmitt. 1967. Rare-Earth Distributions. IN.:
Researches in Geochemistry. Vol. 2. P. H. Abelson (ed.). John Wiley &
Sons, Inc. pp. 234-259.
3-92
-------
Haskin, L. A., et al. 1966. Meteoric, Solar, and Terrestrial Rare Earth
Distributions. IN: Physics and Chemistry of the Earth. L. H. Ahrens et
al. (eds.) Vol. 7. Oxford, Pergamon Press.
Latimer, W. M. 1952. The Oxidation States of the Elements and Their Poten-
tials in Aqueous Solutions. Prentice-Hall, Inc.
Seme, R. J. and Dhanpat Rai. 1976. Adsorption-Precipitation Behavior of
Eu in Soils and Standard Clays. Agronomy Abstracts, p. 132.
Schumm, R. H., D. D. Wagman, S. M. Bailey, W. H. Evans, and V. B. Parker.
1973. Selected Values of Chemical Thermodynamic Properties. Tables for the
Lanthanide (Rare Earth) Elements (Elements 62 through 76) in the Standard
Order of Arrangement. NBS Technical Note 270-7. Schwille, F., W. Lippok and
D. Weisflog. 1967. Model Experiments on Fluid Flow in the Transition Zone
from Unsaturated to Saturated Soil. IAEA-SM-93/11, pp. 151-159.
Walker, J. B., and G. R. Choppin. 1967. Thermodynamic Parameters of Fluoride
Complexes of the Lanthanides. Advances in Chemistry Series No. 71. The
American Chemical Society, Washington D.C. pp. 127-140.
Weast, R. C., Editor. 1976. Handbook of Chemistry and Physics. The Chemi-
cal Rubber Co., Cleveland, Ohio, p. B313.
3-93
-------
IODINE
Natural Soil and Rock Distributions
The iodine content of volcanic rocks is given in Table 3-44. Yoshida
et al. (1971) tried leaching the iodine from volcanic rocks, and usually
removed less than 20% of the total iodine in the rock. The average I value
for igneous rocks probably falls within the range of 75 to 150 ppb. Nearly
the same average value also would be applicable to metamorphic rocks.
TABLE 3-44. IODINE CONCENTRATION IN VOLCANIC ROCKS
Rock Type
Basalt (BCR-1)
Columbia River, WA
Basalts, Japan, 15
Andesites, Japan, 10
Andesite (AVG-1)
Guano Valley, OR
Oacites, Japan, 3
Rhyolite, Japan, 2
Obsidians, Japan, 3
Obsidian, Oregon
Obsidian, Arizona
Obsidian, Utah
Obsidian, Mexico
I. PPb
160
Reference
Becker and Manuel, 1972
29-140
22-260
270
43-220
26-320
20-65
730
730
540
1900
Yoshida et
Yoshida et
Becker and
Yoshida et
Yoshida et
Yoshida et
Seeker and
Becker-and-
Becker and
Becker and
al, 1971
al., 1971
Manuel, 1972
al., 1971
al., 1971
al., 1971
Manuel, 1972
Manuel-, -1972
Manuel, 1972
Manuel, 1972
The abundance of iodine in sedimentary rocks is shown in Table 3-45.
Note that the units have changed from ppb in igneous rocks to ppm in sedi-
mentary rocks. Pre-1960 iodine analytical values may have been as much as
35 times too high (Brehler and Fuge, 1974).
While the iodine content of marine sediments depends on the carbon to a
large extent, marine sediments are much richer in iodine than sedimentary
rocks. Walters and Winchester (1971) examined iodine binding in dry marine
sediments and found that most of the iodine was surface-adsorbed or covalent
bonded to carbon. About 23% could be extracted with organic solvents.
There is a marked increase in the soil iodine content compared to the
rocks they are derived from. For example, the average iodine content of
soils was noted as 5 ppm by Vinogradov (1959), who also suggested that much
of the soil iodine is atmospherically derived.
3-94
-------
TABLE 3-45. IODINE IN SEDIMENTARY ROCKS
Sedimentary Rock
I, ppm
Sandstones, 11,
Bashkiria, USSR .
Sandstone, white,
Klondyke, MO
Sandstone, red,
Potsdam, NY
Argillaceous Sandstone,
Portageville, NY
Argillaceous Shale,
Rochester, NY
Calcareous Shale,
Lima, NY
Limestones, 6,
Paleozoic, OK
Sediments,
0.5-1.5
0.068
0.1.4
37.6
13.0
38.0
4.2
405
Reference
SW Barents Sea
Itkina and Lygalova, 1964
Becker et al., 1972
Becker et al., 1972
Becker et al., 1972
Becker et al., 1972
Becker et al., 1972
Collins et al., 1971
Price et al., 1970
Many of the iodine studies have noted the correlation of iodine content
and humus in the soil (Sinitskaya, 1969; Irinevich et al., 1970; Pennington
and Lishman, 1971). De et al. (1971) contacted soils and soil clays with
iodide solutions at 20°, 30° and 40°C containing 0.127 g I"/l, 0.635 g l"/l
and 1.269 g l"/l. Only the clay minerals showed iodide uptake, with illite
adsorbing more iodide than kaolinite or montmorillonite.
Brief Chemistry
There are 24 known isotopes of iodine with 18 of these isotopes having
half-lives of less than 1 day. The only stable isotope is 127I. Only 129I,
with a half-life of 1.7 x 10 years is of long-term interest in high-level
waste disposal. The fission product I with a half-life of 8.07 days, and
I with a half-life of 60 days, are often a short-term disposal hazard
because of their anionic character and resulting lack of soil adsorption.
Although iodine is known in the (-1), (I), (III), (V) and -(VII) oxidation
states, its usual occurrence is as the (-1) (iodide) state in fresh waters with a
o
radius of 2.20 A (Goldschmidt et al . , 1926) and as iodate-iodine in marine or simv
lar aqueous environments (Sugawara and Terada, 1958). I" tends to be a dispersed
element in many environments as a result of its large size in comparison with
fluorine, chlorine, bromine and hydroxyl ions. Oxidation of iodide ions to
3-95
-------
produce the iodate ion (103) is easily accomplished in basic solution by
the reaction: 3I2 + 60H" * 51" + I0~ + 3H20 (Cotton and Wilkinson, 1962).
Iodine can form complexes with metal ions, but these are generally the least
stable of all the halide complexes, with a few exceptions. Iodine also is a
volatile element, subliming at atmospheric pressures without melting.
Examples of iodine minerals include marshite (Cul), iodargyrite (Agl) and
coccinite (Hglg), bellingerite [Cu(I03)2L salesite [CuI03(OH)] and lautarite
[Ca(I03)2] (Fleischer, 1966). Most such iodine minerals are confined to
unusual environments such as the Chilean nitrate deposits.
Solid Phase and Soil Solution Equilibria
Most of the compounds of iodine are very soluble. Some of the insoluble
or sparingly soluble compounds include the iodides of Pb and Pd, the hypo-
iodites of Ag and Hg, and Ba periodates (Pourbaix, 1966). In normal soils
the concentration of most of these elements (Pb, Pd, As, Hg and Ba) that
iodine forms compounds with is very low. Therefore, they would not be expected
to be present in soils.
Iodine can exist in (-1), (I), (III), (V) and (VII) oxidation states. Out
of all of these oxidation states, (-!•) (I", iodide) is the most important and
its domain of predominance extends all along the pH scale, almost completely
covering a large part of the stability domain of water (Pourbaix, 1966). Thus,
it is not surprising that iodine in aqueous solutions free from oxidizing
agents generally exists as iodide.
Experimental Adsorption Results
Raja and Babcock (1961) examined the behavior of I with two California
soils, kaolinite and bentonite. Results of pretreatment by autoclaving,
131
peroxide oxidation, alcohol digestion and extraction of I with various salt
solutions indicated that the bulk of the iodide retained by soils was due to
reaction with organic matter.
Goldberg et al. (1962) determined iodide adsorption on Rainier tuff
ground to 100 to 200 mesh from simulated groundwater. The value obtained
using 131I was 1.10 ral/g.
3-96
-------
1 2]
Kepak (1965, 1966) studied the sorption of I on hydrated aluminum
and ferric oxide. One wt% silver oxide was added to the ferric oxide adsor-
bent. The adsorbent of ferric oxide was 0.06 to 0.1 mm in grain size and
the influent solution contained 0.1M NaNO, at pH 6.9. Slow column break-
131 131
through of the I began immediately even from a I-distilled water solu-
tion. Sixty-six free column volumes (void fraction = 0.66) were loaded to
1% column breakthrough and 758 free column volumes to 90% I breakthrough.
The most extensive study of iodine retention by soils was made by
Wildung et al. (1975). Iodine was applied to the soils as iodide (I") and
methyl iodide (CHjI), a colorless transparent liquid when fresh, with one
part soluble in 50 parts of water. The solution normally remained nonionic
during the distribution studies. Twenty-two soil types were collected for
use in the study in Oregon, Washington and Minnesota with a range of proper-
ties as listed in-Table 3-46. The I" and CH3 I were used in trace quan-
tities in 0.01M CaCl2 solution.
TABLE 3-46. RANGE OF SURFACE SOIL PROPERTIES USED IN THE METHYL IODIDE
AND IODIDE ADSORPTIONS (WILDUNG ET AL., 1975)
Contents,_wt;L
Cation Exchange Soil Paste Organic
Capacity, meg/100 q pH Carbonate Carbon Sand Silt Clay
5.5 - 90.0 3.6 - 8.9 0 - 6.5 0.23-28.8 14.1-73.1 17.6-58.0 38-46.6
The equilibrium distribution coefficients were treated statistically
(Table 3-47). Methyl iodide retention by soils was largely a function of
organic carbon and clay content, while iodide adsorption was a partial function
of silt content. The mechanisms and causal relationship between the silt and
iodide were under further study but were not covered in the report. The
mechanisms underlying the cation exchange capacity-methyl iodide relationship
are difficult to envision. Regression equations were listed to obtain
iodide and methyl iodide Kd values over the soil parameter ranges studied.
Migration Results
Field Studies-
Brown (1967), in the study of groundwater beneath disposal sites at
125
Hanford, mentioned the presence of I in the groundwater. The groundwater
125
samples were taken from monitoring wells and the I concentration were all
less than 2 x 10"5 uCi/mK
3-97
-------
TABLE 3-47. Kd VALUES AND CORRELATIONS RELATING SOIL PROPERTIES
IODIDE AND METHYL IODIDE ADSORPTION (WILDUNG ET AL.,
1975)
Correlation Coefficients of Independent and
Dependent (Kd) Variables
Independent
Soil Variables Iodide Methyl Iodide
Silt wtS 0.47* 0.21
Clay wtS 0.22 0.43*
Organic Carbon wtS 0.06 0.85**
Cation Exchange Capacity 0.21 0.79**
pH -0.05 -0.41*
Statistically significant at the 5% confidence level
Statistically significant at the \% confidence level
Laboratory Studies-
Jacobs (1965) studied the desorption of radioiodine from clays. The
clays and Clinch River (Oak Ridge) floodplain sediment were contaminated with
iodine vapor from the oxidation of I-tagged KI. Desorption studies indi-
cated that the adsorption iodine was in the form of a mixture of HI, I,, and
adsorbed iodine. The rate of desorption decreased with .decreased surface
concentration, and after 24 hr, the iodine remaining was proportional to the
reciprocal of the absolute air temperature. Desorption occurs most rapidly
in high humidity air. Most of the iodine was readily removed with tap water.
After ten 25 ml increments of Oak Ridge tap water through a 10 g, iodine
contaminated Clinch River floodplain sediment, less than 5% of the iodine
remained on the sediment. However, removal from 10% to 5% required 60% of
the water leach volume, indicating that part of the iodine is relatively
tightly held. Some preliminary work with methyl iodide suggested that this
chemical form of iodine would move readily through the ground without being
adsorbed.
Summary
Because the anion exchange capacities of most soils are minimal over a
normal pH range of 6 to 8, the adsorption of iodide (I"), iodate (IQ^) and
organic-iodine molecules is normally also minimal (Wildung etal., 1975).
However, with soils of low pH (4 to 6), the iodide Kd value can rise to
3-98
-------
50 ml/g. Wildung et al. (1975) reported no statistical correlation between
organic content and iodide Kd, as they did for methyl iodide. The range of
methyl iodide Kd values was from 0.1 to 3.1, much lower than the iodide ion
Kd range, even though the correlation between methyl iodide Kd values and
soil organic content was statistically significant. Others (Sinitskaya, 1969;
Irinevich et al., 1970; Pennington and Lishman, 1971) have shown a positive
correlation between iodine content and organic material in the soil where the
iodine is apparently covalently bonded to carbon (Walters and Winchester, 1971)
It should be kept in mind that iodine, especially when present as an organic-
iodine molecule, can change to a vapor phase and migrate much more rapidly
than in an aqueous phase.
References
Becker, V. J., J. H. Bennett, and 0. K. Manuel. 1972. Iodine and Uranium in
Sedimentary Rocks. Chem. Geol. 9:133.
Becker, V. J. and 0. K. Manuel. 1972. Chlorine, Bromine, Iodine, and
Uranium in Tektites, Obsidians and Impact Glasses. J. Geophys. Res. 77:5353.
Brehler, B. and R. Fuge. 1974. Iodine. IN: Handbook of Geochemistry,
Vol. II/3, Springer-Verlag, New York. p. 53-1.
Brown, D. J. 1967. Migration Characteristics of Radionuclides Through Sedi-
ments Underlying the Hanford Reservation. ISO-SA-32.
Collins, A. G., J.. H. Bennett, and 0. K. Manuel. 1971. Iodine and Algae in
Sedimentary Rocks Associated with Iodine-Rich Brines. Bull. Geol. Soc. Am.
82:2607.
Cotton, F. A. and G. Wilkinson. 1962. Advanced Inorganic Chemistry. Inter-
science Publishers.
De, S. K., S. S. Rao, C. M. Tripathi, and C. Rai. 1971. Retention of Iodide
by Soil Clays. Indian Journal of Agricultural Chemistry. 4:43-49.
Fleischer, M. 1966. Index of New Mineral Names, Discredited Minerals and
Changes of Mineralogical Nomenclature in Vols. 1-50 of the American Mineralo-
gist. Am. Mineral. 51:1248.
Goldberg, M. C., V. J. Janzer, C. G. Angelo, and W. A. Beetem. 1962. The
Effect of Sodium Ion Concentration on Distribution Coefficients for Tuffs
from NTS. U.S.G.S. Technical Letter NTS-16.
Goldschmidt, V. M., T. Barth, G. Lunde, and U. W. Zachariasen. 1926. Geo-
chemische Verteilungsgesetze der Elemente. VII. Skrifter Norske Videnskaps-
Akad. Oslo, I. Mat-Naturv. Kl., No. 2.
3-99
-------
Irinevich, A. D., I. Z. Rabinovich, and V. A. FiTKov. 1970. Iodine in
Moldavian Soils. Pochvovedenie. 58-.
Itkina, E. S. and V. W. Lygalova. 1964. Geochemistry of Iodine and Bromine
in the Carboniferous Horizon and in the Domanik and Bavly Formations of Some
Petroliferous Areas of Bashikiria. IH: The Geochemistry of Oil and Oil
Deposits. L. A. Gulyayeva (ed.). Israel Program for Scientific Translations.
Jacobs, D. G. 1965. Desorption of Radioiodine from Clays. ORNL-P-1531.
Kepak, F. 1965. Sorption of Small Amounts of Radioiodine as Iodide Anions
on Hydrated Ferric Oxide Containing Silver. Collection Csechoslov. Chem.
Commun. 31:1493-1500.
Kepak, F. 1966. Sorption of the Radioisotopes 35S, 131I, and 06Ru on
Hydrated Oxides in Laboratory Columns. Collection Czechoslov. Chem. Commun.
31:3500-3511.
Pennington, W. and J. P. Lishman. 1971. Iodine in Lake Sediments in Northern
England and Scotland. Biol. Rev. 46:279-313.
Pourbaix, M. 1966. Atlas of Electrochemical Equilibria in Aqueous Solutions.
Pergamon Press, Oxford, England.
Price, N. B., S. E. Calvert, and P. G. W. Jones. 1970. The Distribution of
Iodine and Bromine in the Sediments of the Southwestern Barents Sea. J.
Marine Res. 28:22.
Raja, M. E. and K; L. Babcock. 1961. On the Soil Chemistry of Radioiodide
Soil Science. 91:1-5.
Sinitskaya, G. I. 1969. Iodine Content of the Zeya-Bureya Plain Soils.
Uch. Zap. Dal'nevost. Gos. Univ. 27:72.
Sugawara, K. and K. Terada. 1958. Oxidized Iodine in Sea Water. Nature.
182:250.
Vivogradov, A. P. 1959. The Geochemistry of Rare and Dispersed Chemical
Elements in Soils. Consultants Bureau,Inc.
Walters, L. J. and J. W. Winchester. 1971. Neutron Activation Analysis of
Sediments for Halogens Using Szilard-Chalmers Reactions. Anal. Chem. 43:1020.
Wildung, R. E., R. C. Routson, R. J. Serne, and T. R. Garland. 1975.
Pertechnetate, Iodide and Methyl Iodide Retention by Surface Soils. BNWL-
1950. Pt. 2. pp. 37-40.
Yoshida, M., K. Takahashi, N. Yonehara, T. Ozawa, and I. Iwasaki. 1971.
The Fluorine, Chlorine, Bromine and Iodine Contents of Volcanic Rocks in
Japan. Bull. Chem. Soc. Japan. 44:1844.
3-100
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NEPTUNIUM
Natural Soil and Rock Distributions
0-57 0-57 ooo 10
Very small quantities of "xNp r0/Npr°0U = 10"'^) were reported by
Keller (1971) but for practical purposes neptunium does not occur naturally
in soils or rocks.
Brief Chemistry
228 241
There are 16 known isotopes of neptunium from Np to Np (Keller,
1971). Only 237Np> a neutron reactor product of 238U (n,2n) and 235U (n.y)2,
with a half -life of 2.14 x 106 years and 239Np (2.35 days) are of interest.
236
The >5000 year half-life Np is a cyclotron product only while the 22-hr half-
23fi
life Np yields Pu as a daughter. Neptunium exists in aqueous solutions in
five oxidation states, Np(III), Np(IV), Np(V), Np(VI) and Np(VII). In the
absence of complexing agents, the first four oxidation states exist as hydrated
ions Np*3-aq, Np+4-aq, NpOt • aq and Npot2 • aq (Keller, 1971). Np(VII) is a
-3
strong oxidizing agent that is stable in strong alkaline solutions as
The most stable oxidation state in solution is the pentavalent one where nep-
tunium exists as a single charged neptunyl ion, NpO? »aq, with symmetrical
± ^
linear bonding -(0-Np-O) . It hydrolyzes only at a pH of greater than 7, dis-
proportionates only at high acid concentrations and forms no polynuclear complexes
(Keller, 1971). The NpO« ion is a poor complexing agent with inorganic ligands.
Solid Phase and Solution Equilibria
The thermodynami c data are lacking for neptunium compounds other than
oxide and hydroxides (Burney and Harbour, 1974). Figure 3-13 relates the
activity of NpOp to pH under an oxidizing environment (p02 0.68) in equilibrium
with neptunium oxide and hydroxides. Since Figure 3-13 was made using Np(V)
(Np02) as the y axis, the solubility of Np(IV) minerals would decrease and
Np(VI) minerals would increase with an increase in reducing conditions. In
an oxidizing environment CpOgfq) s °-68 atm^ tne soll'ds i° an increasing
order of stability are Np(OH)4, Np02, NpOgOH, Np02(OH)2. Under the oxidation-
reduction conditions represented by p02(g) > 3.8 atm, Np02 would be the most
stable compound. Since Np02 maintains very high concentrations of neptunium
in solution, especially in acidic and very oxidizing environments, it is
unlikely that Np02 would be found as a discrete solid in such terrestrial
environments. It may exist in very reducing conditions.
3-101
-------
S
-2
I I
NpOjOH; Np02 AT pO. • 2.08 atm; Np02(OH)2 AT p02 • 3.84 atm
10
PH
Figure 3-13. The relative stability of various neptunium solids in an
oxidizing soil environment [pO
2(g)
0.68 atm]
The activity of various solution complexes of neptunium in equilibrium
with Np02(s) in an oxidizing environment (p02 = 0.68) and assumed weather-
ing environment is given in Figure 3-14. Thermodynamic data for NpO« HPOT
were obtained from Sillen and Martell (1964). The remaining data were from
Burney and Harbour (1974). The solution complexes of Np(IV), Np(V), and
Np(VI) in order of increasing importance are Np(IV), Np(VI), and Np(V). The
activity of Np(IV) and Np(VI) complexes is so small that these complexes can
be safely ignored since they would not contribute significantly to the total
activity of Np in solution (Burney and Harbour, 1974). The lines of Np(V)
and Np(VI) complexes will shift downward with the increase in reducing envi-
ronment but still be parallel to the present lines. Under the conditions
assumed in Figure 3-14, NpO is the most dominant solution species in a pH
range of 0 to approximately 9. Beyond pH 9, NpOgHPO" and NpOpHCO^ would
control the total concentration of neptunium in solution.
is the usual form of charged neptunium species up to pH 9. Above
pH 9, an uncharged bicarbonate complex is formed. According to Keller (1971)
neptunium (V) does not hydrolyze below a pH of 7. Hence, a singly charged
3-102
-------
Figure 3-14. The activity of various neptunium species in equilibrium
with NpO?(s) in an oxidizing soil environment [p02(g) =
0.68 atm], pCT = pS042' = 2.5, pH03" = 3.0, pF~ = 4.5
and
5.0.
neptunyl ion is the usual form. Hence, NpO« would be expected to enter into
•f +2
ion exchange reactions. NpQ2 does not compete favorably with Ca and other
common divalent ions, and the distribution coefficients are usually rela-
tively low on most soils (Routson et al., 1976; Sheppard et al., 1976).
Experimental Adsorption Results
Bensen (1961) examined the adsorption on minerals of radionuclides in
237
reactor effluent cooling water. One of the radionuclides studied was Np,
which was adsorbed on 25 different common sulfide, silicate and carbonate
minerals, 0.05 to 0.25 mm, in equilibrium tests with trace amounts in Columbia
River water at 80°C. According to Bensen, neptunium was adsorbed poorly or
not at all by the minerals tested.
Robertson (1974) determined the speciation of neptunium in the cooling water
effluent from the Hanford N reactor. Through the use of filter membranes and
3-103
-------
cation and anion exchange resins, the following distribution was observed:
26% participate, 70% cationic, <3% anionic and <1%. non-ionic.
Routson et al. (1975, 1976) determined neptunium Kd values for a
Washington sand and a South Carolina sandy clay. The properties of these
soils are given in Table 3-48. The neptunium Kd values are given in Table 3-49.
Pre-equilibrations of the soils with nonradioactive solutions prior to contact
237
with the traced solution containing Np were carried out. Calcium nitrate
and sodium nitrate salts were used as calcium and sodium ion sources. Cation
exchange of NpO cannot be the principal adsorption mechanism because Na con-
centration essentially does not affect the neptunium Kd values from no sodium
competition to 3.0M Na competition. Calcium has some affect on the neptunium
Kd, but much less than there would be if neptunium removal and adsorption were
due to ion exchange.
TABLE 3-48. PROPERTIES OF SOIL SAMPLES
(ROUTSON ET AL., 1976)
CEC,
Soil CaCQq. mq/q Silt, wt% Clay. wt% meg/ 100 q pH
Washington Soil 0.8 10.1 0.5 4.9 • 7.0
(Surbank sandy loam)
South Carolina <0.2 3.6 37.2 25 51
Subsoil
TABLE 3-49. NEPTUNIUM Kd (ml/g) AS A FUNCTION
OF SOIL AND SOLUTION (ROUTSON
ET AL., 1976)
Ca Na
Soil
Washington
South Carolina
0.002 M
2.37
0.25
0.2 M
0.36
0.16
0.015 M
3.9
0.7
3.00 M
3.2
0.4
Sheppard et al. (1976) determined neptunium distributions on several
237
Washington and North Carolina soils. The distilled water-trace Np solu-
tions and soils were equilibrated over long periods of time to obtain a rate
of change per month in the distribution values. The 50-day values are given
in Table 3-50. A positive value for the monthly change in the distribution
3-104
-------
means an increase in the size of the distribution value. The distribution
was determined in the same way that a Kd value is determined, but does not
have the connotation that equilibrium has been attained as the Kd does.
TABLE 3-50. NEPTUNIUM 50-DAY DISTRIBUTION VALUES COMPUTED
FROM SHEPPARD ET AL. (1976)
Soil Identity
Muscatine
silt loam
Burbank
loamy sand
Ritzville
silt loam
Fuquay sand
0-5 'cm depth
Fuquay sand
5-15 cm depth
Fuquay sand
15-50 cm depth
Np Distribution, ml/g
127
15.4
20.2
33.7
37.2
32.4"
Monthly Change
in Distribution
+10%
+48%
+ 28%
+ 25%
+25%
+19%
The neptunium Kd value of about 3.9 ml/g was given by Routson et al.
(1975, 1976) for the Burbank sandy loam without sodium in the solution.
Sheppard et al. (1976), apparently for the same conditions, gives a neptunium
50 day distribution value of 15.4 ml/g for Burbank and states that this value
is increasing at the rate of +48%/month. At 100 days, the distribution was
approximately 50 ml/g on the Burbank loamy sand. The differences in neptunium
distribution between soil and solution probably resulted from the experimental
differences because the Burbank sandy loam of Sheppard was obtained from
Routson. However, Sheppard shows the Burbank of his sample to contain 21.2%
silt and 2.8% clay with a pH of 8.1 and a cation exchange capacity of
5.94 meq/100 g compared to Routson's Burbank soil sample (Table 3-48).
Sheppard1s Burbank must have contained calcite (CaC03) to attain a pH of 8.1.
3-105
-------
There is no evidence in Sheppard et al. (1975) that the soil samples were
pre-equilibrated in any way with the nonradioactive solution before the .
neptunium tracer was added, or that a blank solution without the soil was
run with the soil equilibrations for reference and use in counting as the
original solution condition without the soil present. In addition, Sheppard
237
reported that his Np(V) in the aqueous phase equilibrating with Ritzville
soil was partially filterable on Whatman number 50 filter paper. Routson,
on the other hand, reported that upon filtering previously centrifuged
samples of solution containing neptunium and in contact with Burbank or
South Carolina subsoil through 0.01 and 0.45 urn filters, no evidence of col-
237
loidal Np was found. This is further evidence that Sheppard1s high nep-
tunium distribution values require further elucidation in the light of
Routson's results.
Migration Results
Field Studies-
No reports of field studies of neptunium migration through soil or rock
were found in the literature.
Laboratory Studies-
No reports of laboratory studies of neptunium migration through soil or
rock were found in the literature.
Summary
The existing thermodynamic data (Figure 3-14) show that neptunium should
exist in an oxidizing soil environment as Np(V) in the form of NpO^. However,
existing adsorption results, essentially only a study by Routson et al. (1975,
1976) and Sheppard et al. (1976), inferred that there is little evidence of
NpO- ion exchange as a neptunium adsorption mechanism. Sheppard et al. (1976)
reported neptunium colloids present in the soil-solution environment, while
Routson et al. (1975, 1976) specifically showed that they were absent from
his solutions that had contacted soils. It is fairly certain that neptunium
Kd values are generally low. Little more can be said based on the adsorption
and migration studies now available in the literature. Additional research
is required on neptunium-soil and rock reactions.
3-106
-------
References
Bensen, D. W. 1961. Mineral Adsorption of Radionuclides in Reactor Effluent.
HW-69225.
Burney, G. A. and R. M. Harbour. 1974. Radiochemistry of Neptunium. NAS-
NRC Nuclear Sci . Ser. NAS-NS-3060.
Keller, C. 1971. The Chemistry of the Transuranium Elements. Vol. 3.
Kernchemie in Einzeldarstellungen. Verlag Chemie GmbH.
Robertson, D. E. 1974. Physicochemical Characterization of N-Reactor
Effluent Radionuclides in Soil and Water Systems. BNWL-1950, Pt. 2,
pp. 82-85.
99
Routson, R. C., G. Jansen, and A. V. Robinson. 1975. Sorption of Tc,
237^ and 241 Am on Two Subsoils from Differing Weathering Intensity Areas.
BNWL-1889.
qq 0-37
Routson, R. C., G. Jansen, and A. V. Robinson. 1976. * Tc, Np, and 'Am
Sorption on Two Subsoils from Differing Weathering Intensity Areas. BNWl-
2000. Pt. 2, pp. 50-52.
Sillen, L. G. and A. E. Martell. 1964. Stability Constants of Metal-Ion
Complexes. Special Publication No. 17. The Chemical Society, London.
%
Sheppard, J. C., J. A. ICittrick, T. L. Hart. 1976. Determination of Dis-
tribution Ratios and Diffusion Coefficients of Neptunium, Americium and
Curium in Soil -Aquatic Environments. RLO-221-T-12-2.
3-107
-------
PLUTONIUM
Natural Soil and Rock Distributions
Except for very small quantities of plutonium in the natural environments
given below, plutonium does not occur in any significant amount in soils or
239
rocks. Cherdyntsev et al. (1968), for example, lists Pu values in natural
-17 ?^Q -T5 239
rocks and minerals varying from 7 x 10 " g "*Pu/g to 1 x 10 g Pu/g.
Brief Chemistry
Fifteen isotopes of plutonium are known, including 10 isomers with dif-
ferent decay modes (Keller, 1971). Most of the isotopes are formed by multiple
neutron capture, and therefore, are not ordinarily encountered in natural
environments. An exception is the natural reactor at Oklo, Gabon (IAEA, 1975)
and very small amounts (Pu/U = 10" ) found in the uranium mineral pitchblende
(Cleveland, 1970).
The plutonium isotopes of concern to waste disposal, either because they
are contained in the fission product wastes or are a part of an impotant decay
chain for waste constituents, are listed in Table 3-51. In the short term,
almost any plutonium radionuclide would be of concern in waste-radionuclide
242 •
reactions although, on a weight basis, lon-ger-lived Pu is not nearly as
239
hazardous as Pu. Long irradiation times and high neutron fluxes are
242
required to produce significant amounts of Pu, so that a prevalence of
the higher mass Pu isotopes usually will not occur in waste solutions from
commercial fuel treatment facilities.
TABLE 3-51
Isotope
236,
PLUTONIUM RADIONUCLIDE DATA
(KELLER, 1971)
238
239
240
241
242
243
'Pu
Pu
Pu
Pu
Pu
Pu
Pu
HaIf-Life
2.85 years
86.4 years
24,400 years
6,600 years
14.1 years
387,000 years
4.98 hours
Decay Mode
a, SF
a, SF
a, SF
a, SF
8", a
a
3-108
-------
Plutonium exists in five oxidation states in aqueous solutions that
include Pu(III), Pu(IV), Pu(V), Pu(VI) and Pu(VII). These states can occur
+3 +4 + +2 -3
as the hydrated ions Pu • aq, Pu • aq, PuCL • aq, PuCL • aq and PuOg •
aq (Keller, 1971). The most stable state of plutonium in aqueous solution
is Pu(IV). However, disproportionation, in which Pu(III), Pu(IV), Pu(V) and
Pu(VI) states can coexist in unequal quantities, tends to further complicate
plutonium chemistry in its reactions with soils and sediments. According to
Cleveland (1970), plutonium in near-neutral solutions occurs in the order:
Pu > PuOl" > Pu > PuOt. Hydrolysis reactions lead to the formation of
positively charged to neutral colloidal polymers in a stepwise manner. For
example, PuH + OH" * PuOH*3 + OH" * Pu(OH)+2> etc., leads finally to
Pu(OH)4 which loses water to product thermodynamically stable Pu02 (Keller,
1971). The kinetics of the hydrolysis reactions leading to PuO« are not
known. It is known that depolymerization of colloidal Pu(IV) is a very slow
process, having a "depolymerization half-life" of 320 hr at 25°C in 5M HN03
(Keller, 1971). Complexing agents such as fluoride or sulfate ions can accel-
erate the process. Plutonium can form complexes with most of the ions com-
monly encountered in soil solutions. Thus disproportionation, complex forma-
tion and hydrolysis reactions, all relatively pronounced with aqueous plutonium
solutions, combine to add to plutonium chemistry complexity.
Solid Phase and Solution Equilibria
4+
Figure 3-15 relates the activity of Pu to pH in an oxidizing environ-
ment CpOo/n) = 0-68 atm] in equilibrium with various Pu solid phases. The
parameters indicated in parentheses after the mineral formulas refer to the
additional conditions of equilibrium. For example, (V&G) indicated after
the Pu(HPO.)2 mineral formula denotes that the mineral is considered to be
in equilibrium with variscite and gibbsite. Any solid in Figure 3-15 that
lies below another solid, at a given pH, is the more stable. Thus, for any
4+
two solid phases at a specific pH, the solid that maintains lower Pu activity
is more stable. Under the assumptions outlined in Figure 3-16 and at pH 6,
the plutonium compounds in order of increasing stability are: Pu(OH)3;
BPu203; PuF4; PuF3; Pu(HP04)2; PuOgOH; Pu02C03; Pu(OH)4; Pu02(OH)2; and
Amont the solids reported in Figure 3-15, Pu02 is the most stable mineral
under oxidizing conditions at all pH values.
3-109
-------
Pu(HP04)2
-------
The stabilities of plutonium solid compounds in a reducing environment
= 80 atm] are represented in Figure 3-16. In general, the reducing
environment substantially changes the stabilities of the solids. Comparing
Figures 3-15 and 3-16, it can be seen that the Pu(IV) solids remain unchanged,
while Pu(III) solids increase and Pu(V and VI) solids decrease in their sta-
bilities with the decrease in the oxygen pressure. Plutonium solids in the
reducing environment at pH 6 can be arranged in order of increasing stability
as follows: Pu02(OH)2; PuOgOH; PuF4; Pu(HP04)2; Pu(OH)3; 6Pu203; Pu02(OH)2;
PuF^; and Pu02. PuF-j would be the most stable solid compound below pH 4.
Above this pH value, Pu02 would be the most stable solid (Figure 3-1.6). Based
upon the Eh-pH diagrams for plutonium oxides, and hydroxides, Polzer (1971)
predicted Pu02(s) to be the most stable solid phase under conditions generally
found in the environment. However, the relative stabilities of Pu compounds
other than hydroxides and oxides, and predictions regarding weathering
sequences, were not considered by Polzer (1971).
Knowledge of the most stable compound present in the soil is necessary
in order to predict the behavior of fate of the element in the soil. An exten-
sive review on actinides indicated that numerous workers have studied the
plutonium concentration and distribution with depth in sofls. However, with •
the exception of Price and Ames (1975), none have made any attempt to identify
the solid phases of plutonium present in soils. Price and Ames (1975) reported
the presence of Pu02 under both alkaline and acidic oxidizing environments.
Their observations were made approximately 20 years after the acidic radio-
active waste was disposed to the alkaline soils. Tamura (1974) inferred the
presence of Pu02 in aridisols or entisols from the Nevada Test Site. Thus,
the Price and Ames, and Tamura results confirm theoretical predictions made
regarding the most stable plutonium compound expected to be present in soils.
However, if the soil solution activity of plutonium is lower than the solu-
bility product of Pu02 and any other possible compound, the final solution
activities will depend mainly on cation exchange processes and the copre-
cipitation of plutonium with other soil minerals.
Figure 3-18 demonstrates the activity of various plutonium species
expected in eauilibrium with PuO«(s) and in an oxidizing environment (p09 16
c 3+ 4+ +
atm). Plutonium in solution exists in four oxidation states (Pu , Pu Pu02,
3-111
-------
Pu02+) and forms complexes with OH', d", F", SO*', H2PO~ and Co*' ions. In
general, the activity of all positively charged ions and ion-complexes and a
few neutral ion-pairs decreases with the increase in pH, while the activity
of all negatively charged ion-complexes increases with the increase in pH.
In an oxidizing environment, Pu + and Pu and their ion-pairs have an insig-
nificant effect on the total activity of plutonium in solution. At any given
pH, the total activity of plutonium in solution can be obtained by adding
the appropriate activities of various plutonium species. Since the lines for
some plutonium species are based on concentration constants, the position of
the lines may change when converted to thermodynamic equilibrium constants.
Figures 3-17 and 3-18 were constructed by assuming a weathering environment
(pC02 = 3.53; pd" = p$0*~ = 2.5; pF~ = 3.5) and Pu02(s) as the compound that
may be present in the soil at equilibrium. When diagrams are constructed
for an actual weathering environment of a given soil, all the organic and
inorganic ligands present in the soil solution should be considered. Fig-
ures 3-17 and 3-18 were developed for selected inorganic ligands only.
-is
Figure 3-17.
The activity of various plutonium species in soil solution in
equilibrium with Pu02(s) at pH 8, pCOj = 3.52 atm, pCT =
- =2.5 and pHjPO^f =5.0 (Rai and Serne, 1977)
3-112
-------
Figure 3-18.
The activity of various plutonium species in soil solution in
with Pu02(s) in a mildly oxidizing soil environment [pOg = 16
atm], pC02 = 3.52 atm, pCl = 2.5, pF~ =3.5 and pHgPO^ =5.0
(Rai and Serne, 1977)
Figure 3-17 relates the activities of (a^, moles/liter) of various solu-
tion species at pH 8 to the oxidation potential as represented by the various
oxygen pressures. In general, with a decrease in oxygen pressure (increase
in reducing conditions), the kind of activity of the solution species
changes: 1) Pu species increase in activity, 2) Pu * species are unaffected
i OL
and 3) Pu02 and PuO^ species decrease in activity. However, in doing so,
the total activity of plutonium in solution at any given pH would not change
significantly because the activity of species for which thermodynamic
equilibrium constants are known is very large as compared to the others.
When the concentration constants are converted to thermodynamic constants,
the new lines would shift downward or upward but remain parallel to the
present lines. The activities of various species would still vary in the same
direction with the changes in pH as reported in Figure 3-18.
3-113
-------
A few of the reactions are not plotted in the figures because either
the activities of the species fall beyond the graph boundaries or the equilib-
rium constant data are not reliable. Cleveland (1970) suggested that the
2+
values for the formation constants of PuCO^ and Pu(OH)4° were suspect; the
values for Pu02(OH)3, Pu02(0)2°, and Pu02OH+ should be viewed with skepticism;
and the value for PuO,OH° should be considered only approximate. Under the
£ ' o
assumptions of Figure 3-18, the lines for Pu(OH).0 and Pu02(OH)2 would be
parallel to the X-axis with 10~12'51 and lO"19-1^ moles/1 activity, respec-
tively. Many researchers (Andelman and Rozzell, 1970; Fukai and Murray, 1974;
Grebenshchikova and Davydov, 1961, 1965; Sheidina and Kovarskaya, 1970) have
reported that plutonium in solution also exists as colloidal species over a
wide pH range. Colloidal species and neutral ion-complexes [Pu(OH)°»,
Pu02(OH)°2]-are not reported in Figures 3-17 and 3-18. If present in large
quantities, these species would significantly influence the nature of the
species and the total concentration of plutonium in solutions.
For the conditions specified in Figure 3-18, the dominant plutonium
species in low and high pH ranges are Pu02 and Pu02C03OH~, respectively.
The colloid complex of soils is mainly negatively charged. Therefore, posi-
tively charged species are usually absorbed by soils in cation exchange pro-
cesses. Since positively charged Pu02 species predominate only in solutions
of pH values <6, cation exchange adsorption of plutonium by soils would be
limited to acidic soils (pH <6). The negatively charged Pu02C03OH~ complex
that predominates above pH 8 [even when the activity of Pu02CO.,OH~ is compared
to Pu(OH)°4] should not be readily adsorbed by the soils. Thus, the activity
of plutonium in soil solutions of pH >8 would be relatively high because of:
1) little soil adsorption of negatively charged ions and 2) increases in the
activities of negatively charged ions with the increase in pH. Similarly,
under extremely low pH conditions the activity of plutonium in soil solutions
in equilibrium with PuOgU) may also be high. Rhodes (1957) found low adsorp-
tion of plutonium by soils in the pH range of approximately <2 and 8-13. Thus,
his results confirm part of the predictions based on Figure 3-18; continuous
decrease in adsorption is predicted with the increase in pH above 8, whereas
Rhodes' data indicated first a decrease in adsorption to pH 12 and then an
unexplained continuous increase in adsorption above pH 12.
3-114
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It should be mentioned that in addition to cation, exchange, the total
retention of plutonium wastes percolating through soils would depend on many
other factors, such as flow rate, precipitation of solid phases, physical
entrapment of particulate and colloidal plutonium, soil properties, and waste-
water composition. Therefore, the above adsorption predictions based only on
cation exchange should be viewed as an approximation. As the data on actual
adsorption isotherms and mechanisms and various soil-actinide interactions
accumulate, more accurate predictions will be possible.
Experimental Adsorption Results
Thorburn (1950) reported a sorption experiment in which plutonium in
water was percolated through a soil column. The soil was then wet sieved and
the activity of each particle size fraction counted. In general the decrease
in particle size increased the amount of Pu adsorbed (Table 3-52). In other ad-
sorption experiments, Thorburn placed small aliquots of actual plutonium wasta
on top of soil columns and eluted with distilled water or saturated magnesium-
calcium carbonate. There were no significant differences in the elution curves.
It is interesting to note that the elution curves had higher plutonium activity
in the first volume out of the column and always showed a small leakage which
Bright signify physical transport of fine particulates, neutral polymers or
colloids. Thorburn took the effluent from the distilled water leached column
and percolated it through a second soil column at various flow rates. The
data show that Pu concentration in effluents through the second column varied
with flow rate. The faster the flow the more plutonium in the effluent. Fil-
tration of the influent through a fine filter paper did not change the results.
Explanations of these phenomena are obscure but would seem to indicate plutonium
transport as polymer. Final adsorption studies of plutonium in water versus
pH by soil columns showed 87% adsorption at pH 3.8, 77% at pH 7.2 and 24% at
pH 10. Note that this falling of plutonium adsorption with increasing pH is
supported by the thermodynamic data of Figure 3-18.
Evans (1956) reported percentage of plutonium removed from a 0.001 N HNO^
solution by various grain sizes of quartz. The plutonium Kd values may be com-
puted from these data, and are listed in Table 3-53. Data also were given on
plutonium adsorption on different minerals of coarse clay size (2 to 0.2 um).
3-115'
-------
TABLE 3-52. PLUTONIUM ADSORPTION VERSUS
PARTICLE SIZE (THORBURN, 1950)
Size, nun Pu. d/min/g
>1.19 39.3
1.19 - 0.25
0.25 - 0.15 1,171
0.15 - 0.10 929
<0.105 8,578
Uncentrifugible
Silt and Clay 62,750
TABLE 3-53. PLUTONIUM Kd VALUES FOR QUARTZ OF VARIOUS
PARTICLE SIZES FROM A 0.001 N HN03-
PLUTONIUM SOLUTION (EVANS, 1956)
Quartz Particle Sizes Kd, ml/g
Coarse sand 5.6
Fine sand 10.0
Coarse silt (50-20 vat) 11.3
Medium silt (20-5 ym) 33.5
Fine silt (5-2 ym) 48.8
Coarse clay (2-0.2 ym) 80.9
Plutonium Kd values were computed from these data and are listed in Table 3-54.
Evans suggested that the lower Kd for plutonium by montmorillonite is probably
due to experimental difficulties and believed that Rhodes' (1955) data may
represent preferred plutonium Kd results. Rhodes' data, as reported by Evans,
are given in Table 3-55, assuming a 1 ml to 1 g solution to soil ratio. This
ratio was not given, but the results are proportional in any case.
Rhodes (1952, 1957a, 1957b) performed similar plutonium Kd determinations
on a Hanford sand soil (92% sand) with 2% calcium carbonate content. Using
a soil to solution ratio of 1:20 and neutral pH, Rhodes found rapid adsorption.
Analysis versus time yielded an equation for the amount adsorbed versus contact
3-116
-------
TABLE 3-54. PLUTONIUM Kd VALUES FOR DIFFERENT MINERALS
OF COARSE CLAY SIZE (2 to 0.2 urn) FROM
0.001 N HN03-PLUTONIUM SOLUTION (EVANS, 1956)
Mineral Kd, ml/g
Feldspar 170.4
Quartz 82.0
Glauconite «
Montmorillonite 157.0
Kaolinite 1091.0
TABLE 3-55. RHODES' PLUTONIUM Kd DATA AS
REPORTED BY EVANS (1956)
Material Kd. ml/g
Montmorillonite 4990
Fine micaceous sand 49
Kaolinite 42.5
Soil 36
Coarse sand .0.87
time, %PuADS = 81 + 5.2 log time/m,-n\- For example, after 1 min, 20 min and
60 min, 81%, 88%, and 90%, respectively, of the solution plutonium concentra-
tion was adsorbed. Increases in the soil to solution ratio were also observed
to increase the percent adsorbed. Rhodes also obtained adsorption versus pH
data and found rapid decreases in adsorption below pH 2 but almost total adsorp-
tion (Kd > 2000) between pH 3 and 8.5. Decreases at pH 10 to 12 were believed
to indicate changes in plutonium speciation. In other experiments, coarse sand
was observed to adsorb only 47% of the plutonium present compared to 97 to
99.8% for soils and clay minerals. Rhodes felt all these data indicated ion
exchange adsorption mechanisms. Yet plutonium adsorption from distilled water,
a synthetic high salt waste and an actual high salt waste all showed greater
than 98% removal. Plutonium solution loading experiments onto soil columns
at various pH's (1,4, 10) showed good removal at the first two pH's until the
3-117
-------
high acid case had destroyed the CaC03 buffering capacity. As the effluent
pH dropped to pH 1, increasing plutonium breakthrough was observed. Plutonium
breakthrough at pH 10 was highly variable and rapid, possibly caused by fine
Plutonium particulate dispersion or polymer formation. Plutonium adsorption
from distilled, water, 4M NaN03 and 2M(NH4)2HP04 solutions was above 96%, but
a 4M ammonium acetate solution showed only 59% adsorption. Acetate complexing
of plutonium was probably responsible for the diminished adsorption rather
than any effects from competing ions.
Further studies showed increasing the acetate concentration from 0.01M
to 8M dropped plutonium adsorption from 74% to 21%. Acetate solutions were
also observed to have the capability of extracting plutonium adsorbed on soils.
Bensen (1960) presented data showing that oxalate salts also inhibit plutonium
adsorption by soluble Pu-oxalate complex formation. Bensen also re-interpreted
Rhodes' data and concluded precipitation reactions, not ion exchange, probably
control plutonium adsorption on soils.
Other investigations on plutonium were performed at Savannah River by
Prout (1953, 1959). Plutonium adsorption by a soil (80% sand-20% clay) with
the clay mineral kaolinite dominant was determined versus pH for distilled
water traced with plutonium (MO M). The soil/solution ratio was 1:10 and
the mixtures shaken for 2 hr before phase separation by centrifugation. Adsorp-
tion of plutonium depended on pH and the valence state of the plutonium added
to solution. Adsorption was more than 90% complete from solutions of Pu(III)
and Pu(IV) between pH 2.5 and 12 and from solutions of Pu(VI) at pH's greater
than 6. Strong adsorption at mildly acidic to mildly basic pH's is probably
due to a combination of cation exchange and precipitation of hydrolysis products.
Decreased adsorption above pH 9 possibly represents the formation of soluble
negatively charged Pu polymers which do not readily exchange with the soil.
Knoll (1965, 1969) and Hajek and Knoll (1966) reported studies on high
salt wastes with organic contaminant? typical"of fuel reprocessing. An acid
(0.3M) high salt (5.4M NO^) waste spiked with actual americium and plutonium
wastes were percolated through soil columns. After 0.3, 1.5 to 4.0, and 20
column volumes of effluent, the plutonium breakthrough was 20 to 30%, 50%,
and 70 to 98%, respectively, for several columns. Batch adsorption tests on
the high salt acid wastes from storage tanks showed low adsorption Kd
3-118
-------
values of from 2.4 to 2.9 ml/g with resultant soil pH's of 2 to 3. If the
waste was neutralized and the supernatant was used in batch adsorption tests,
the Kd increased. Addition of suspected organic contaminants to a concentra-
tion of 20% by volume decreased the Kd for the neutralized waste to 1.4. Water
leachates of the sludge produced on neutralization contained some plutonium.
Upon soil contact the Kd for this leachate rose to 1540 ml/g. If the high salt
acid waste was neutralized in the presence of molar quantities of citric acid,
no precipitation was observed. If this citrate-neutralized waste was perco-
lated through soil columns, most of the plutonium was removed. A constant low
leakage percentage of plutonium was observed and may have been caused by a
solubility or particle transport mechanism.
In other experiments (Knoll, 1969) tap water spiked with plutonium was
percolated through soil columns to allow plutonium adsorption. The soil
columns were then leached with various organics: TBP-CCl^, DBBP-CC1,, Fab Oil,
D2EHPA in hydrocarbons C,Q-C-,^, and hydroxyacetic acid. These organics are
all commonly used in various process streams during fuel reprocessing. TBP in
carbon tetrachloride after 130 column volumes had caused only 5% of the soil
sorbed plutonium to leach; DBBP removed 40% of the soil sorbed plutonium in
-80 column-volumes; Fab Oil removed less than 4% in .30 column volumes; D2EHPA
removed about 30% of the plutonium; and hydroxyacetic acid removed no less
than 40% of the plutonium depending on the acid concentration. When the plu-
tonium was added to the organics and then percolated through the soils, very
little soil adsorption was observed.
Many investigations have been performed on plutonium adsorption by sedi-
ments in both fresh and marine waters. Schneider and Block (1968) found
• j.0 iO J.
plutonium to adsorb more than I and Sr but less than Zn or Cs on Rhine
River sediments. Equilibrium was reached within 24 hr.
The location of plutonium at ground disposal sites on the Hanford Reserva-
tion has been closely monitored for over 20 years. Liquid wastes disposed to .
the unsaturated sandy soil in several instances included trace amounts of
plutonium. Typically, plutonium has not been found to percolate through the
soils to any extent. Typical data (Brown, 1967) show 99.9+% of long-lived
radionuclides such as plutonium to be adsorbed within the first 10 m of soil.
3-119
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Crawley (1969) reports on plutonium penetration at two facilities." Plutonium
in a sand to which slightly acidic high salt wastes were released had pene-
trated at least 18 ft. Plutonium in a sand to which neutralized low salt
water had been released penetrated less than 2 ft.
Desai and Ganguly (1970) determined the ability of humic and fulvic acid
in ammoniacal solutions to solubilize plutonium. In this basic solution with-
out humic or fulvic extract, only 13% of added plutonium remained in solution;
with humic extract 54% of the added plutonium remained in solution; and
further testing showed that the soluble plutonium was noncationic. The fulvic
acid extract solubilized 58% of the plutonium, and again the complex was
noncationic.
Rozzell and Andelman (1971) studied the adsorption-desorption of aqueous .
-8 -7
plutonium (10 to 10" M) on quartz and other silica surfaces. Adsorption
continued typically for 12 to 15 days before apparent equilibrium was reached.
At pH 7, adsorption increased with increasing ionic strength, but decreased
when HCOg was added. The amount of adsorption varied at pH 5 and 7, but dif-
ferently at high and low ionic strengths, as well as with age of the solution.
Plutonium desorption indicated that there were two basically different adsorbed
species. There was a great net desorption at pH 5 versus pH 7 to 9.
Tamura (1972) discussed the adsorption of plutonium on several materials
including the soil constituents listed in Table 3-56. The suspension pH repre-
sents the final system pH.
TABLE 3-56. REMOVAL OF PLUTONIUM FROM pH 7 WATER
BY SEVERAL SOIL MINERALS (TAMURA, 1972)
Constituents Kd, ml/g Suspension pH
Attapulgite 4,370 9.60
Montmorinonite 630 9.20
Alumina, activated 755 8.35
Kaolinite 352 6.25
Illite 129 5.90
Quartz 0 6.35
3-120
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Fried et al. (1973, 1974, 1974) found that plutonium adsorbed on lime-
stones, shales, and sandstones would very slowly migrate through the pores.
The velocity of migration relative to the leaching water was 3 x 10 for
-5
limestone and 6 x 10 for basalt. If the leaching solution was saturated
with C02 an increased elution from the limestone and sandstone was observed.
Studies of plutonium movement along a fissure crack in basalt shows plutonium
+4
movement in two phases: the slower moving phase adsorbs similar to Pu , but
a faster moving phase with only about 1/10 the tendency to adsorb was also
observed. Fried et al. hypothesized a polymeric state for this phase.
Fried et al. (1974, 1974) immersed disks of limestone and basalt in
Pu(N03)4 solution at 4 x 10 M Pu. After equilibration the activity in the
solutions was re-measured. Losses in plutonium were equated to adsorption on
the rocks. Basalts adsorbed more plutonium than limestone on a surface area
basis. The effects of salt solutions of Na , Ca , Sr , La , and Zr were
also studied. In all cases the salts displaced some of the plutonium from
the rocks. The higher the concentration and the larger the valence charge,
the easier was the displacement. This was similar to the trends expected for
ion exchange mechanisms. A leaching study showed eight column volumes of
0.5M HC1 removed 30% of the plutonium adsorbed on_shale. Eight additional
column "volumes of 4M HC1 removed an additional 20%.
Two review articles which contain discussions on adsorption reactions
of plutonium and soils are available, Francis (1973) and K. Price (1971, 1973).
Ames (1974), Price and Ames (1975) and Ames (1976) characterized the
actinide-bearing sediment underneath a liquid waste disposal facility which
received high salt, acidic wastes. The distribution in plastic impregnated
core samples showed a surface soil plutonium concentration of near 0.5 mCi
239 " 239
Pu/g. This concentration decreases to about 0.6 uCi Pu/g within the
first 2 m of underlying sediment and to less than 60 pCi/g at the maximum
depth sampled (9 m). Examination of the contaminated soils showed at least
two types of plutonium present: 1) discrete Pu02 particles (>70 wt% Pu02) and
2) plutonium (<0.5 wt% Pu02) associated with silicate hydrolysis. The Pu02
particulate form "filtered out" within the upper 1 m of sediment. The
"hydrolysis" type of plutonium penetrated deeper within the sediment and
deposited in association with silicate hydrolysis products as the soil minerals
3-121
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reacted with-the acid waste. The highest concentrations of nonparticulate
Plutonium about 1 m below the bottom of the disposal facility receiving acidic
wastes were associated with smectites present in the soil before waste solu-
tion disposal. The smectites higher up in the column tended to be partially
to totally altered because of the acidic nature of the waste solutions. Plu-
tonium was uniformly distributed in the sectioned smectite particles, suggest-
ing adsorption of charged plutonium species by the clays.
Emery and co-workers (1974, 1974, 1975) have studied the ecological
behavior of one Hanford liquid waste pond. Ninety-five percent of the incoming
low-level waste percolates through the desert sands. The sediments are the
principal repository, with the top 10 cm containing an average 390 pCi Pu/g.
The overlying water contains 0.01 pCi Pu/1. Assuming that these values con-
stitute equilibrium conditions, an estimate of the plutonium Kd can be made.
The resultant Kd value is 3.7 x 10 ml/g. Plutonium in the interstitial
waters appears to be mainly cationic or nonionic forms. Plutonium in the pond
water appears to be mainly fine particulates possibly signifying that the waste
stream plutonium content is predominantly Pu02 particles.
Miner et al. (1973, 1974), Polzer and Miner (1974, 1977) and Glover et al.
(1977) described-experiments in progress to measure the adsorption of both
"soluble" plutonium and particulate plutonium (Pu09) in numerous soils at
6 8
three plutonium levels between 10 and 10 M. The plutonium adsorption was
rapid and quite high: 59% of the time, adsorption was greater than 99% (Kd >
430) and only 5% of the time was it below 90% adsorption (Kd < 39). Correla-
tions were found between soil characteristics associated with soil ion exchange
or acidity. The adsorption also depends on the initial plutonium concentrations
for which several possible causes are discussed including precipitation with
increasing plutonium concentration and the formation of carbonate complexes.
Experiments at Savannah River (Savannah River Laboratory, 1975) with
degraded TBP-DBBP-kerosene wastes containing plutonium showed rapid migration
through soils (Kd ^ 0.4 ml/g). When small portions of the solvent waste were
allowed to dry on soils and leached by groundwater, however, very little plu-
tonium migration was observed (Kd ^ 6000 ml/g) in the effluent.
3-122
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Duursma et al. (1971, 1973, and 1974) have determined the adsorption of
Plutonium in the marine environment. The adsorption of plutonium is inter-
mediate; less than Cs, Rb, Zn, Fe, Zr-Nb, Ru and Pm and greater than Ca and
Sr. Plutonium has been observed to penetrate sediments at least to 8 to 14 cm,
and in the early 1970's, inventories in the sediment appear to show losses.
Biological processes or upward migration in interstitial waters as a reduced
species are two possible explanations, although Mo and Lowman (1975) showed
data which find reducing conditions tend to lower plutonium mobility. Changes
in the americium to plutonium ratio from waters to sediments in the ocean led
Duursma to conclude americium is preferentially adsorbing or not being remobilized
as fast as plutonium. In seawater Pu02(C03)3~ may predominate and cause the
237
lower plutonium adsorption. Using a Pu tracer in valence states +3, +4, and
+6 in both oxic and anoxic seawaters at pH 7.8 to 8.0, Duursma found the'Kd
for marine sediments to be about 10 ml/g. Valence states and solution Eh
did not cause marked differences.
Fukai and Murray (1974) performed plutonium adsorption-desorption experi-
ments for both freshwater and saline conditions. Plutonium (III) adsorption
experiments from river water to river sediments (100 ml traced solution/g
sediments) were performed. The initial stage of adsorption was rapid; 50% .
within 10 hr. For pH 3 to 11 over 90% of the Pu(III) was adsorbed by the
river sediment. Desorption experiments were performed wherein the sediments
used in the adsorption experiments were suspended in seawater. Over the pH
range 4 to 10 about 10% of the adsorbed plutonium was released without apparent
pH differences. Similar adsorption-desorption experiments were performed with
river sediments contacted with various dilutions of freshwater sewage effluent
at pH 8.1. The Pu(III) adsorption increased from 85% in river water to over
95% at 100% sewage effluent. These sediments were next contacted with sea-
water and desorption monitored. At equilibrium, river sediments that had not
been contacted with sewage effluent released only 5% of the plutonium, whereas
sediments contacted with 5 'to 100% by volume sewage effluent released 12%.
Van Dal en et al. (1975) determined the plutonium distribution coefficient,
Kd, for several Dutch subsoils from 90% saturated NaCl solutions between pH 7
4
and 8. Clay samples of mainly illite and kaolinite had a Kd of MO ml/g and
3-123
-------
for a river sand a Kd of 200 ml/g. Gypsum-bearing and clay-bearing sandstones
were intermediate. The pH dependence on adsorption of plutor.ium was minor
between pH 5 and 8.
Mo and Lowman (1975) placed piutoniurn-contaminated calcareous sediment
in aerated seawater and anoxic seawater and stirred the system until solution
activities were at equilibrium. The resultant plutonium Kd values were 1.64
x 104 and 3.35 x 105 ml/g.
Tamura (1975, 1975) evaluated the plutonium content of various size frac-
tions of contaminated Nevada Test Site and Oak Ridge soils and Mound Laboratory
sediments. These results should be considered exploratory in nature because
the results are based on only a few observations; the source terms differed;
and the particle size fraction separations were technique dependent. The NTS
soils showed different particle sizes with plutonium associations depending on
location from ground zero. In general, the coarse silt (20 to 53 um) and medium
silt (5 to 20 ym) fractions contained about 80 to 85% of the total plutonium.
The Oak Ridge plutonium was more evenly distributed in the silt and clay sizes,
with clay slightly higher. The Mound Laboratory sample showed 70% of the
plutonium in the clay fraction.
Bondietti et-al., (197-5.)-- found, purified soil humates" to absorb'plutonium
strongly, 99.9+% while reference clays (montmorillonite and kaolinite) adsorbed
about 98 and 96% of plutonium added to a water solution at 163 iag Pu/g of solid.
If the organic matter and iron oxide coating were removed from clay material,
plutonium adsorption was 95% versus 99.9% while intact. Leaching solutions,
citrate, DTPA, and fulvate removed small amounts of adsorbed plutonium as
shown in Table 3-57. Thus, it appears that natural soil organic matter can
very strongly adsorb plutonium. The ability of soil organic matter to alter
the oxidation state of plutonium in soil water solutions was also verified.
Pu(VI), and by inference Pu(V), were unstable in the presence of fulvic acid,
polygalacturonic acid and alginic acid. Positive proof of reduction to Pu(IV)
was shown at environmental pH's 6.5 to 8. Plutonium (IV) was the most stable
valence upon interaction with these organics. Further reduction of Pu(IV) to
Pu(III) occurred in the presence of humic or fulvic acids, but was not observed
3-124
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above pH 3.1 under aerobic conditions. Bondietti also demonstrated that plu-
tonium is at least partially associated with humic materials in ORNL soil
contaminated 30 years ago.
TABLE 3-57. PERCENT PLUTONIUM LEACHED BY EXTRACTANTS
(BONDIETTI ET AL., 1975)
Percent Leached
Extractant
Humate
0.3
30
5
Montmorillonite
8
15
0.9
Kaolinite
11
20
3
10"3M Citrate
10"3M DTPA
10"2M Fulvate
Bondietti (1974) determined the plutonium Kd for clays separated from a
soil (Miami silt loam) and montmorillonite using 5 x 10" M calcium solutions
at pH 6.5 for two plutonium valence states. The results are shown in Table
3-58. The high adsorption of Pu(VI) on the soil clay, in contrast to or as
compared with montmorillonite, indicated that plutonium reduction by clay com-
ponents must have occurred, since adsorption values similar to the Pu(IV) were
obtained. Pu(VI) appeared to adsorb less than Pu(IV).
TABLE 3-58. PLUTONIUM Kd AS A FUNCTION OF OXIDATION STATE
(BONDIETTI, 1974)
Material Initial Oxidation Kd, ml/g
Soil Clay Fraction Pu(IV) 1.0 - 1.7 x 105
Pu(VI) 7.5 x 104
Montmormonite Pu(IV) 2.1 x 104
Pu(VI) 2.5 x 102
'Relyea and Brown (1975) studied the adsorption-desorption relations of
plutonium solutions added to water at pH 2. The ratio of plutonium adsorbed
by the soil to that in the equilibrium solution was greater than 99.1. Dif-
fusion of plutonium in four soils was found to be much slower than for exchange-
able soil cations.
3-125
-------
237
Volesky and Friedman (1976) used Pu to study the adsorption of plu-
tonium on Niagara limestone surfaces (k), and the effect of sodium chloride
237
concentrations on k values. A k is defined as ( Pu activity/ml of solution)/
237 2 2
( Pu activity/cm of limestone), with the dimensions ml/cm . The k value
is essentially constant until NaCl is 1M in concentration. The k then
increases rapidly to 3M NaCl and again becomes constant. As the k value
237
increases, Pu is being desorbed from the limestone and re-entering the
solution. The desorption may result from the formation of plutonium chloride
complexes.
Rai and Seme (1977) used thermodynamic data to develop solid phase-soil
solution equilibria diagrams which can be used to estimate plutonium behavior
in terrestrial environments. Pu02 was the most stable phase in the pH and
oxygen partial pressure ranges found in soil environments. Plutonium experi-
mental results found in the literature were used to verify empirical predic-
tions, based on the diagrams, of plutonium adsorption on soils.
Migration Results
Field Studies—
Magno et al. (1970), in investigations of radionuclide migration from the .
Nuclear Fuel Services plant in New York State, estimated that more than 90%
0-30
of the "°Pu, Pu and Pu contained in the plant effluent remained in the
sediments of the lagoon system.
A series of wells were drilled into the contamination produced by the
1973 leakage of 435,000 liters of high level waste from a Hanford storage tank
(Anon., 1973). Core sample analysis showed the relative mobilities of the
radionuclides contained in the high pH, high salt solution. Plutonium had
migrated the shortest distance from the tank rupture of all the radionuclides
in the contaminated sediments.
Emery et al. (1974a, 19746, 1975) have discussed the plutonium inventory
2
of a 56,657 m reprocessing waste pond at Hanford. The pond has received
239 240
wastewater since 1944, which included about 1 Ci of ' Pu. It has
retained greater than 99% of this plutonium in the upper 20 cm of its sediments
even though no water leaves the pond except by evaporation or soil percolation
and the mean residence time of the water is only 40 hr (Emery et al., 1974).
The plutonium is probably present in a parti cul ate form.
3-126
-------
The plutonium resulting from operation of the natural reactor at Oklo
1.8 billion years ago was reported to have remained in the vicinity of the
original reactor site (Brookins, 1976).
Bondietti and Reynolds (1975) investigated the species of plutonium found
in Oak Ridge seepage water from a sol'id waste burial ground. Pu(VI) was iden-
tified as the main Pu species in the seepage solutions from the burial ground.
The implication is that Pu(VI) is more mobile than Pu(IV).
Water infiltration and plutonium migration at Maxey Flats burial grounds
were studied by Meyer (1975). A hydrostatic head of 4 to 6 m was found in
some trenches, causing lateral and downward migration of leachates. Surficial
plutonium migration also had occurred from trench leakage. Above background
plutonium concentrations have been detected in onsite and offsite surface
soils and stream sediments, in soil cores to 90 cm and in wells 15 to 27 m
deep. Other radionuclides accompanying the plutonium included tritium, Co,
89'9°Sr and 134-137Cs.
Radiological measurements at the Maxey Flats burial site, Kentucky, also
have indicated some physical radionuclides transport on particulate matter
(Meyer, 1975; Montgomery et al., 1977). The authors mentioned that 54Mn, 60Co,
90 137 238 239
Sr, Cs, Pu and Pu were detected in stream sediments, although the
water levels were quite low. In E-Series test well measurements, all plutonium
within the detection limit was associated with sediment. The authors suggest
that this raises a question as to the mechanism of plutonium migration from
the disposal trenches.
Laboratory Studies—
Lindenbaum and Westfall (1965) prepared several dilute solutions of
colloidal plutonium with citrate to plutonium ratios of 1800 (3.4 x 10" M
citrate) to 1. The colloid was about 80% ultrafilterable at pH 4.0 and a
citrate to plutonium ratio of 1. The cellophane used in the ultrafiltration
had a pore size that corresponded to a plutonium hydroxide molecular weight
of about 200,000. At pH 10, the citrate to plutonium ratio 1 colloid fell to
65% ultrafilterable and 5% at pH 11. The citrate to plutonium ratio 1800
colloid behaved in a similar fashion falling from 92% ultrafilterable at pH 4
3-127
-------
to 72% at pH 11. With no citrate present, the plutonium was about 13% ultra-
filterable at pH 4 and fell to near 0% at pH 11. Qualitatively, Zr(IV), Th(IV)
and the rare earths may behave in a similar manner, peptizing at higher pH,
as suggested by the works of Rhodes (1957). The dispersed polymers are less
soil-reactive and should migrate farther through the soil. Hajek (1965) deter-
mined that the migration rate in soil, relative to groundwater, was 1/100 and
1/1000 times the rate of groundwater flow for strontium and cesium, respectively.
The diffusion velocities of plutonium and americium were 1/10,000 times the
transporting solution velocity (Hajek, 1966). The apparent self-diffusion
-12 2
coefficient for plutonium was 4.8 x 10 cm /sec, so that the average distance
diffused in 329,000 years was 10 cm (Hajek, 1966). The 329,000 years was
239
13.5 half-lives for Pu. Relyea (1977) determined the diffusion coefficient
for plutonium in the same soil using a quick freeze method developed by Brown
et al. (1964). He obtained an'apparent plutonium diffusion coefficient of
-10 2
5 x 10 cm /sec at 791 hr on the same Burbank loamy sand.
Hajek (1966) reported that plutonium mobility from contaminated soils was
low in the Hanford 216-Z-9 covered trench. About 0.1% of the plutonium could
be leached with groundwater, but about 7.5% of the resident americium was
removed. No further plutonium was leached after a throughput of 13 column
volumes of groundwater. At least 3.5% of the plutonium on the soil was removed
by 20 column volumes of IN NaNOg solution. The leach rate was constant after
20 column volumes. One column volume was equivalent to 61 cm of leaching
solution.
The role of microorganisms in the movement of soil plutonium has been
under study by Au (1974). He reported that the plutonium discrimination factor
was about 4 between agar growth medium at pH 2.5 and Aspergillus spores. The
pH 2.5 probably aids in maintaining plutonium mobility. However, the usual
discrimination factor between soil and plants, for example, is 10 to 10
(Francis, 1973). Acid-producing molds that could increase plutonium solubili-
zation were also identified in the soil.
Fried et al. (1974), in their plutonium distribution studies in limestones
and basalts, indicated that the adsorption coefficients were dependent on the
types and amounts of other ions in solution. At least two chemical forms of
3-128
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Plutonium were migrating in neutral solutions. One of these, possibly a
polymerized plutonium oxide, migrated ten times faster through rock fissures
than the other form.
Jakubick (1975) modeled the migration of plutonium down the soil profile
and found that the migration rate of PuCL (0.8 cm/yr) was about 100 times
faster than plutonium applied to the soil as Pu(N03)4-
Nishita et al. (1976) determined the extractability of plutonium from a
contaminated soil as a function of pH. The pH influences hydrolysis and pre-
cipitation of the plutonium. In the study, 2 g of contaminated soil was agi-
tated in 25 ml of solution, in duplicate, and the pH adjusted with NaOH or
HN03. The Kd values are given in Table 3-59 for the Aiken clay loam, a
kaolinitic soil, at several pH values. It is not clear from the report whether
these values are initial or final system pH measurements. The pH values, at
the high and low ends of the range especially, usually drift considerably
during equilibrations with soils so a final pH measurement is most meaningful.
The Kd values are given in Table 3-59 because it should be of little importance
as to which direction equilibrium is approached.
TABLE 3-59. Kd VALUES OF 238Pu IN UNTREATED" AIKEN
CLAY LOAM (NISHITA ET AL., 1976)
_pji_
1.21
2.12
2.56
4.69
5.55
7.08
"°Pu Kd, mg/Jl
430 ± 44
862 ± 33
796 ± 59
2591 ± 591
2347 ± 220
3086 ± 19
pH
8.54
9.43
10.31
11.25
12.22
13.25
"°Pu Kd, mq/2,
963 ± 130
512 ± 14
3.02 ± 1.6
213 ± 5.8
138 ± 4.8
207 ± 14
The migration of plutonium and americium in the lithosphere was investi-
gated by Fried et al. (1976). Most of the plutonium and americium were strongly
held on tuff, basalt and limestone. Migration rates of 10 m/m of water flow
were measured in one case, and 30 um/m of water flow for plutonium migration
in another instance into limestone. A small amount, less than 35%, of the
plutonium appeared to migrate at a rate about 10 times as fast as the rest of
3-129
-------
the plutonium. This rapidly moving plutonium may be adsorbed on fine mineral
particles. Also, the chemical environment can drastically increase the migra-
tion rate of plutonium and americium. Even dilute solutions (10 M) of highly
charged ions cause substantial desorption and migration of plutonium and
americium adsorbed on rock surfaces.
Relyea and Brown (1975) measured the diffusion coefficients of plutonium
in soil and in water by a capillary tube diffusion cell suggested by Brown et
al., (1964). The soil diffusion coefficients were measured by contacting cells
containing contaminated and uncontaminated soils for a known time interval,
freezing instantly with liquid air, and slicing and counting of sections of
the untagged soil perpendicular to the diffusion axis. The apparent diffusion
-8 2
coefficients were low compared to normal soil cations, 1.4 x 10 cm /sec in a
-11 2
sandy soil to less than 5 x 10 cm /sec in a silt loam. Associated Kd values
were 300 to 500 ml/g for the sandy soil and 10,000 ml/g for the silt loam.
EDTA and DTPA reduced most Kd values by a factor of up to 100. The aqueous
-7 2
diffusion coefficients varied from 3.1 x 10 cm /sec in solution extracted
-5 2
from the silt loam to 2.7 x 10 cm /sec in a solution extracted from the sandy
soil.
Fried et al. (1977) reported the upper 'limits of the relative migration
ratios (radionuclide distance moved/water distance moved) of plutonium and
americium through rocks to be 100 um/m and 500 um/m of water flow, respectively.
Using a migration ratio of plutonium of 125 um/m of water flow, and a water
flow of 0.6 km/yr, the plutonium would migrate no more than 12.5 cm/yr. The
241
Am moves faster but has a shorter half-life. The authors concluded that
deep burial in geological formations would provide effective isolation of
actinides from the environment. Several rather large assumptions are contained
in this conclusion. One of these assumptions is that the small cores used to
obtain the migration data perfectly model a geologic waste repository.
Further work on the rapidly migrating form of plutonium was reported by
Friedman et al. (1977). In one case, a HOEHP column was used to remove all
polymeric plutonium species before the experiment, and the rapidly moving plu-
tonium was eliminated. In another treatment, the plutonium solution was evapo-
rated repeatedly with HN03 with the same experimental results. Thus, the.
identification of the rapidly moving plutonium form as a polymer was verified.
3-130
-------
A model was devised to predict the flow of wastes containing plutonium through
rock fissures. The model will be field tested by core drilling a former dis-
posal site at Los Alamos and comparing predicted and actual movement of the
actinides through the tuff.
Fried et al. (1977) also studied the retention of Pu(VI) on Los Alamos
tuff with a stock solution of plutonium containing 268 ug Pu(VI)/ml. The
+2
adsorption spectrum of PuO at 830 nm was used to verify the purity of
the stock solution. A 10 ml tuff cylinder, coated with wax on the exterior,
was -used at a flow rate of 2 ml/hr. Prior to use, the tuff column was vacuum
evacuated under water to ensure saturation. Measured amounts of Pu(VI) were
placed on top of the column and allowed to percolate downward. Water was
slowly passed down the column to simulate the passage of groundwater. The
effluent water was collected in several fractions and analyzed. The value
obtained for Pu(VI) migration was 26,500 um/m of water flow through the tuff.
The comparable Pu(IV) migration was 100 um/m of water flow. After eight free
column volumes (2.4 ml/free column volume) passed through the tuff, a switch
to 4N HNO^ rapidly stripped most of the remaining plutonium. A small amount
of residual plutonium, evenly distributed in the tuff column, may or may not
have been Pu(VI).
Wildung et al. (1973, 1977) have considered the influence of soil micro-
bial processes on the long-term solubility, form and plant availability of the
transuranic elements. Diffusible plutonium in soils, usually less than 0.1%
of the total plutonium, appeared to be particulates of hydrated oxide but
there was evidence of microorganisms affecting the solubility, form and plant
availability of transuranics. Plutonium toxicity was a function of plutonium
solubility in soil but microorganisms were generally resistant, with toxicity
due to radiation effects. Highly resistant bacteria, fungi and actinomycetes
were isolated from the soil and shown to transport plutonium into the cell,
altering its form to a soluble, negatively-charged complex. The plutonium
forms were not well defined, but were dependent on organism type, carbon source
and plutonium exposure time during growth.
3-131
-------
Summary
In most oxidizing environments, ionic plutonium is expected to be present
in solution predominantly as Pu(V) or Pu(VI) (Figures 3-17 and 3-18). With
high plutonium concentrations, PuCL could be expected as a stable solid phase
(Figures 3-15 and 3-16). The plutonium adsorption data indicate that more
than one mechanism is operating during reactions with rocks and soils. There
is a direct correlation of plutonium adsorption results with soil cation
exchange capacities (Thorburn, 1950; Evans, 1956; Rhodes, 1952, 1957a, 1957b;
Van Oalen et al., 1975; Glover et al., 1977) indicating that part of the
adsorption is due to ion exchange. However, little effect on adsorption was
obvious from competing cation concentrations that were very high (Rhodes,
1957a; Mo and Lowman, 1975). Perhaps the principal correlation of plutonium
adsorption is with pH (Rhodes, 1952, 1957a, 1957b; Knoll, 1965; Hajek and
Knoll, 1966; Knoll, 1969). A relatively rapid soil column breakthrough of
plutonium was obtained with a low pH (<1) solution (Knoll, 1965). Strong
adsorption occurs over the pH range of from 4 to 8 (Prout, 1958, 1959), but
above pH 8 the formation of negative complexes or polymers with low charge
densities can cause lowered adsorption (Thorburn, 1950; Rhodes, 1957a; Polzer
and Miner, 1977). In addition, plutonium is easily complexed with- humic acids
(Desai and Ganguly, 1970; Bondietti et al., 1975) and with oxalate (Bensen,
1960), acetate (Rhodes, 1957a) and a large number of phosphorus-bearing organics
used in the nuclear fuels reprocessing industry (Knoll, 1969). Complexed plu-
tonium also can result from soil fungal or microbial activities (Wildung
et al., 1973, 1977; Au, 1974) all of which result in more mobile plutonium.
Even the precipitated or adsorbed plutonium can be readily moved through the
aqueous environments in the form of particulates or suspended sediments
(Magno et al., 1970; Emery et al., 1974a, 1974b, 1975; Meyer, 1975;
Montgomery et al., 1977). If the Eh environment allows it, Pu(IV) can be
oxidized to Pu(V) or (VI), with much lower plutonium adsorption values result-
ing (Bondietti, 1974). The plutonium polymers usually are adsorbed readily,
but can migrate more rapidly than soluble plutonium in some instances where
the surface charge density is low (Friedman et al., 1977). Much experimental
work remains to be done to separate and understand the several reactions
mechanisms occurring in plutonium adsorption on rocks or soils. A controlled
Eh environment for the adsorption work is important as an experimental condi-
tion when an element disproportionates as readily as plutonium.
3-132
-------
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Prout, W. E. 1959. Adsorption of Fission Products by Savannah River Plant
Soil. DP-394.
Rai, Dhanpat and R. J. Seme. 1977. Plutonium Activities in Soil Solutions
and the Stability and Formation of Selected Plutonium Metals. Journal
Environmental Quality. 6:89-95.
Relyea, J. F. 1977. .The Behavior of Plutonium in Soil. Ph.D. Thesis,
University of Arkansas, Fayetteville, AR.
Relyea, J. F. and D. A. Brown. 1975. The Diffusion of Pu-238 in Aqueous and
Soil Systems. Agronomy Abstracts, p. 124.
Rhodes, D. W. 1952. Preliminary Studies of Plutonium Adsorption in Hanford
Soil. HW-24548.
Rhodes, D. W. 1957a. The Adsorption of Plutonium by Soil. Soil Science.
84:465-471.
Rhodes, D. W. 1957b. The Effect of pH on the Uptake of Radioactive Isotopes
from Solution by a Soil. Soil Science Society of America, Proceedings.
21:389-392.
Rozzell, T. C. and J. B. Andelman. 1971. Plutonium in the Water Environment.
II. Sorption of Aqueous Plutonium on Silica Surfaces. Advances in Chemistry
Series. TQ6:280-98.
3-136
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Savannah River Laboratory Quarterly Report Waste Management. April-June 1975.
DPST-75-125-2.
Schneider, H. and W. Block. 1968. On the Question of the Capacity of the
Rhine for Radioactive Nuclides. Sorption of Radionuclides by Sediments of
the Rhine. Gas Gasserfach. 109:1410-15 (in German).
Sheidina, L. D. and E. N. Kovarskaya. 1970. Colloidal State of Pu(IV) in
Aqueous Solutions. Sov. Radio. Chem. 12:229-233.
-Tamura, T. 1974. Distribution and Characterization of Plutonium in Soils
from Nevada'Test Site. p. 24-42. IN.: P. B. Dunaway and M. G. White (ed.)
The Dynamics of Plutonium in Desert Environments. NVO-142.
Tamura, T. 1972. Sorption Phenomena Significant in Radioactive - Waste
Disposal. IN: Underground Waste Management and Environmental Implications.
T. D. CookTid.). Am. Assoc. Petrol. Geol. pp. 318-330.
Tamura, T. 1975. Characterization of Plutonium in Surface Soils from Area
13 of the Nevada Test Site. IN: The Radioecology of Plutonium and Other
Transuranics in Desert Environments. M. G. White and P. B. Dunaway (eds.).
NVO-153, pp. 27-41.
Tamura, T. 1975. Physical and Chemical Characteristics of Plutonium in
Existing Contaminated Soils and Sediments. IAEA-SM-199/52.
Thorburn, R. C. 1950. Absorption on Hanford Soil and Related Soil Properties.
HW-15655. ' - . .
Van Dalen, A., F. DeWitte, and J. Wiskstra. 1975. Distribution Coefficients
for Some Radionuclides Between Saline Water and Clays, Sandstones and Other
Samples from the Dutch Subsoil. Reactor Centrum Nederland. pp. 75-109.
Volesky, A. F. and A. M. Friedman. 1976. Niagara Limestone Absorption of
Plutonium from Salt Solutions. ANL.
Wildung, R. E. and T. R. Garland. 1973. Influence of Soil Microbial Activity
on the Uptake and Distribution of Plutonium in the Shoots and Roots of Barley.
BNWL-1850, Pt. 2, pp. 22-25.
Wildung, R. E. and T. R. Garland. 1977. The Relationship of Microbial Pro-
cesses to the Fate and Behavior of Transuranic Elements in Soils and Plants.
IN: The Transuranium Elements in the Environment. W. C. Hanson (ed.).
3-137
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PROMETHIUM
Natural Soil and Rock Distributions
Promethium has not been found to occur naturally in soils and rocks.
Brief Chemistry
There are 16 known radioactive isotopes and isomers of promethium with
mass numbers from 141 to 154 and half-lives ranging from 34 sec to 17.7 years
(Lavrukhina and Pozdnyakov, 1970). Promethium-147 (half-life 2.5 years) is
235
the only promethium isotope obtained in large amounts from U fissioning by
thermal neutrons, and is present in fuel reprocessing fission products.
Promethium is a lanthanide and like most lanthanides exists in a trivalent
state. Only the (III) oxidation state of promethium is stable in aqueous soil
solutions. Starik and Lambert (1958) studied the state of micro-amounts of
promethium in aqueous solutions at different pH values by ultrafiltration and
centrifugation. The results showed that promethium was present as Pm at
pH < 3. With an increase in pH, hydrolysis products gradually accumulate as
multiply-charged aggregates.m After a maximum is reached at pH of about 7, the
charged aggregates begin to flocculate. At higher pH values, the aggregates
become neutral or negatively charged. Electrophoteric measurements showed that
from pH 5.3 to 6.7, the bulk of the polynuclear particles was positively
charged. At pH 9, no mobility is seen, indicating neutrally charged particles.
At pH 5 to 8, no promethium can be centrifuged, indicating a colloid in solu-
tion. At pH 9, uncharged forms such as [Pm(OH)x(N03)3_x] predominate (Starik
et al., 1959).
Solid Phase and Soil Solution Equilibria
Promethium forms crystalline compounds with F", Cl", Br", I, OH", P04 ",
AsO 3~, B033~, Cr04, Cr03> Mo042~, W042", V043~, Sc033" (Lavrukhina and
Pozdnyakov, 1970; Wheelwright, 197.3). However, the thermodynamic data are
available for Pm(OH)3 (solid) (Sillen and Martell, 1964) and Pm203 (Pourbaix,
1966). It is evident that Pm(OH)3 is more stable than Pm203 (Figure 3-19).
The Pm activity in equilibrium with Pm(OH)3 decreases a thousand-fold with
a one unit increase in pH.
3-138
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Promethium is a lanthanide and like most lanthanides exists in a trivalent
No data were found on solution species of promethium except for
state (Pm3+).
(PnT + NOZ * PmNO, log K° at I = 1 is 2.48) (Lavrukhina and
J 3+
Pozdnyakov, 1970). The solution species is probably predominantly Pm , much
like europium. If it assumed that promethium does not form any other complexes,
then the concentration of promethium in solution will be governed mainly by
Pm . Under these conditions, ion exchange can be an important promethium
removal mechanism in acidic environments and at trace promethium concentrations.
Figure 3-19. The relative stability of various promethium solids.
Experimental Adsorption Results
Bensen (1960) reported the complete uptake of promethium and other triva-
lent rare earths at pH > 6. Promethium behavior in soil systems was said to
be identical to cerium behavior. The uptake of promethium is nearly complete
and unaffected by 0.5M alkali metal and 0.25M alkaline earth metal cations at
above pH 7. Below pH 3, the adsorption of all rare earths including promethium
is depressed similarly by accompanying salts. The results indicate that
promethium and other rare earths are ionic species below pH 3, and begin to
precipitate as charged polymers above pH 3. Bensen (1960) also reported that
the species and amount of accompanying anion affect promethium adsorption by
soils. 'Citrate, acetate and carbonate ions were found to inhibit rare earth
3-139
-------
adsorption on soils. This was attributed to formation of complex ion species.
Chloride and nitrate ions had no appreciable effect on adsorption, and sulfate
had very little effect.
Migration Results
Field Studies—
Magno et al. (1970) investigated migrating radionuclides from the Nuclear
Fuel Services plant in western New York State. They estimated from analytical
147
data that greater than 90% of the Pm discharged from the plant was deposited
by sedimentation in the lagoon system. These results imply that most of the
promethium was a precipitate, probably Pm(OH)3, or adsorbed on other solids
in the water that settled out in the lagoons.
Laboratory Studies—
Ames (1960) reported the removal of promethium during a replacement
reaction of calcite with phosphate in solution. The promethium was finally
2+ 3+
immobilized in apatite, Ca^POJgOH, where 2 Ca are replaced by a Pm and
Na or seme other univalent cation.
Schulz (1965) classified promethium with other rare earths as immobile
in soils due to either very strong adsorption by clay particles or present
as insoluble hydroxide.
Summary
Low concentrations of promethium are expected in equilibrium with
in slightly acidic to slightly alkaline conditions (Figure 3-9) so that
PmfOHjj precipitation may control promethium concentrations in soil and sedi-
ment solutions. Promethium, like europium, would be expected to be ion
exchangeable below the concentration required for Pm(OH)- precipitation. Only
limited data are available on promethium adsorption on soils and sediments,
but these data tend to support the above predictions.
The removal of promethium is nearly complete on soils and rocks above
pH 6 (Bensen, 1960). The effect of competing cations on promethium adsorption
above pH 7 is minimal. Below pH 3, the adsorption of all rare earths including
promethium are similarly depressed by competing salts. Charged polymers can
form under the proper conditions during Pm(OH)3 precipitation (Starik et al.,
3-140
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1959; Bensen, 1960). Citrate, acetate and carbonate ions inhibited rare
earth adsorption on soils., probably due to the formation of complex ion species.
Over 90% of the promethium discharged to a lagoon system remained in the lagoon
(Magno et al., 1970) as strongly adsorbed on clay particles or present as an
insoluble hydroxide (Schulz, 1965). Promethium can be removed from solution
during replacement reactions (Ames, 1960).
References
Ames, L. L., Jr. 1960. Some Cation Substitutions During the Formation of
Phosphorite from Calcite, Economic Geology. 55:354-362.
Benson, D. W. 1960. Review of Soil Chemistry Research at Hanford. HW-67201.
Lavrukhina, A. K. and A. A. Pozdnyakov. 1970. Analytical Chemistry of
Technetium, Promethium, Astatine and Francium. Translated by R. Kondor.
Ann Arbor - Humphrey Science Publishers.
Magno, P., T. Reavey, and J. Apidianakis. 1970. Liquid Waste Effluents from
a Nuclear Fuel Reprocessing Plant. BRH-NERHL-70-2.
Pourbaix, M. 1966. Atlas of Electrochemical Equilibria in Aqueous Solutions.
Pergamon Press, Oxford, England.
Schultz, R. K. 1965. Soil Chemistry of Radionuclides. Health Physics.
11:1317-1324. . '
Sillen, L. G. and A. E. Martell. 1964. Stability Constants of Metal-Ion Com-
plexes. Special Publication No. 17. The Chemical Society, London.
Starik, I. E., N. I. Ampelogova, F. L. Ginzburg, M. S. Lambert, I. A. Skulskii,
and V. N. Shchebetkovskii. 1959. On the Molecular State of Ultra-Small
Quantities of Radioelements in Solutions. Radiokhimiya. 1:370-378 (in
Russian).
Starik, I. E. and M. S. Lambert. 1958. State of Microquantities of Promethium
in Aqueous Solutions. ZhNKh. 3:136.
Wheelright, E. J. 1973. Promethium Technology. Am. Nucl. Soc. Publication.
3-141
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RADIUM
Natural Soil and Rock Distributions
The radium content of various igneous and sedimentary rocks is given in
Table 3-60. Vinogradov (1959) reported that the radium content of soils
ranged from 0.5 x 10" to 1.1 x 10" ppm. The small amount of radium present
in soils has become separated from parent thorium and uranium during rock
weathering processes, and most closely follows barium. The radium content of
igneous rocks increases by 100 times from ultra-basic to granitic rocks. The
radium content of most sedimentary rocks is about the same as that for
granites.
TABLE 3-60. AVERAGE RADIUM CONTENT OF VARIOUS ROCK TYPES
Rock Type
Ultrabasic Igneous
Basic Igneous
Intermediate, Igneous
Granitic Igneous
Sandstones
Shales
Limestones
Ra, ppm
Reference
0.009 x 10
0.6 x 10
0.917 x 10"
1.395 x 10
0.71 x 10
1.08 x 10
0.42 x 10
-6
-6
-6
,-6
-6
,-6
Davis, 1947
Evans et al., 1942
Senftle and Keevil, 1947
Senftle and Keevil, 1947
Bell et al., 1940
Bell et al., 1940
Evans and Goodman, 1941
Brief Chemistry
213 230 218
There are 16 isotopes of radium from Ra to Ra with no Ra or
229
Ra. All of the isotopes of radium are unstable, and all of the naturally-
228 '24
occurring radium isotopes occur in the thorium decay series ( Ra, " Ra),
23fl 226 235
the U-radium decay series ( Ra), and the U-actinium decay series
223
( Ra) (Vdovenko and Dubasov, 1973). Radium radionuclide data are given in
228
Table 3-61. All are alpha emitters except Ra which is a beta emitter.
With the present uranium fuel cycle and the much longer half-life, only the
226
Ra isotope is of long-term concern in waste disposal. If future energy
228
production includes thorium fuels, then Ra also must be included in waste
disposal management plans.
3-142
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TABLE 3-61. RADIUM RADIONUCLIDE DATA
(VDOVENKO AND DUBASOV, 1973)
Isotope
223Ra
224Ra
225Ra
226Ra
228Ra
Half-Life Decay Mode
11.43 days
3.64 days
14.8 days
1622 years
5.77 ± 0.02 years
ct
a
6"
a
B"
Radium is a homolog of the alkaline-earth elements, with a (II) oxidation
state. The radius of the +2 radium ion is 1.52 A compared to 1.43 A for Ba+2.
The compounds formed by radium and their solubilities are similar to barium.
For example, the solubility product of radium sulfate is 4.25 x 10 at 25°C
(Vdovenko and Dubasov, 1973) and the solubility product of barium sulfate is
1.08 x TO ° at 25°C (Weast, 1976). Metallic radium dissolves in water with
the evolution of hydrogen and the formation of readily soluble Ra(OH)9. Of
+2
the alkaline-earth metal cations, Ra shows the least tendency for complex
formation, although 1:1 complexes with citric, tartaric, succinic and several
other acids were detected at pH 7.2 to 7.4 by Schubert et al. (1950). It may
+2
be assumed that Ra is not hydrolyzed in aqueous solutions, in an analogy
+2
with Ba , although there is no specific literature on the subject.
Solid Phase and Solution Equilibria
The thermodynamic data for radium compounds are available only for radium
nitrate, chloride, iodate and sulfate (Parker et al., 1971). However, all
of the compounds except sulfate are very soluble. Therefore, solid phase
diagrams are not presented for radium compounds. The solubility product for ,
RaS04 is 10~10'37 compared to 10"9'96 for BaS04-
No thermodynamic data were located for radium hydrolysis or complex ion
species. It is expected the radium will behave in the soil solution much like
2+
strontium does. The species Ra is expected to be the most important over
the normal soil pH range from 4 to 8.
Experimental Adsorption Results
Stead (1964) gave a radium Kd value of 6700 ml/g for NTS tuff. Arnold
and Grouse (1965) ran batch adsorption tests on some exchange materials that
included the natural zeolites, clinoptilolite and chabazite, as represented
by a pelletized molecular sieve (AW-500) and barite (barytes), a natural BaSOd.
3-143
-------
The results of the adsorption tests were recomputed as distribution coefficients
or Kd values. The solution was a lime-neutralized waste that contained 4100
pCi 226Ra/l, 500 mg/1 Ca+2, 80 mg/1 Mg+2, 1000 mg/1 Na+, 2500 mg/1 SO^2 and
900 mg/1 Cl" at pH 7.7. The radium Kd values are given in Table 3-62.
TABLE 3-62. RADIUM Kd VALUES FROM LIME-NEUTRALIZED WASTE;
1.25 g EXCHANGER/1 OF WASTE (ARNOLD AND GROUSE,
_J965)
Exchanger Loading,
Exchanger Mesh Size _ pCi/g _ Kd,ml/q
Clinoptilolite 20-50 2650 646
Chabazite 20-50 2900 707
Barite 20-50 2000 490
R. J. Seme of PNL, 1974, used soils from Utah and simulated river water
to determine radium distribution coefficients. The soils were pre-equi libra ted
by four washings with the simulated river water composition shown in Table
?2fi 226
3-63 minus "the RaCl. The fifth solution contained the Ra as well as
the other constituents, and was used for the radium Kd determinations. The
Kd values were determined in triplicate to allow measurement of precision.
The Utah soils contained 2 to 5% calcite, with 'quartz and feldspar constituting
the bulk of these sandy, arid soils. Minor constituents included hydromica
and a small amount of a smectite clay. The radium Kd results are listed in
Table 3-64. The Kd correlated with the cation exchange capacities of the soils.
TABLE 3-63. SIMULATED RIVER WATER COMPOSITION
(SERNE, 1974)
Constituent
Ca
Mg
Na
K
. HC°3
so4
Cl
u
Ra
Added as-
CaS04-2H20, CaCl2
Mgso4
NaCl
KC1
NaHC03
CaS04»2H20, MgS04
CaCl2, NaCl, KC1
uo3
RaC19 in HC1 solution
mg/1
82
26
75
3.4
171
246
57
1
7 ug/1
3-144
-------
TABLE 3-64. RADIUM DISTRIBUTION COEFFICIENTS WITH THE SOLUTION
OF TABLE 3-63 (SERNE, 1974)
Son
i
ii
in
IV
Final pH
7.9
7.9
8.0
7.6
7.7
7.6
7.8
7.9
7.8
7.8
7.6
7.8
Kd,
354
289
467
214
ml/g
± 15
t 7
t 15
± 15
Migration Results
Field Studies--
Granger et al. (1961) and Granger (1963) showed that radium had migrated
out of Ambrosia Lake, New Mexico, uranium ores and had been partly reconcentrated
in barite (BaSOJ and cryptomelane (KMn +Mn % 9C01C'H90) found in and near
226 2+ /.to 10 (.*,
some of the ore bodies. The Ra occurred in the Ba position in barite
.. 2 L ' ? 2 fi
and the Mn position in cryptomelane. The high concentrations of. Ra-were
not associated with parent uranium, which is good evidence for the recent
migration of radium. The strongly to weakly radioactive cryptomelane partially
replaced mudstone that occurred near the ore bodies and was relatively low
in uranium content and enriched in lead. This suggests that the lead is
226
radiogenic and has also migrated with the Ra. The mechanism involved in
??fi 226 2+
reconcentration of the Ra is the substitution of Ra for the chemically
2+ 2+
very similar Ba in barite and Mn in cryptomelane. Analyses of the outer
5 to 10 cm of mudstone layers near ore disclosed anomalously high radioactivity
coupled with an abnormally high lead content. Within mudstone layers, however,
the radioactivity was essentially in balance with the uranium content,' and
the lead content was low.
Hansen and Huntington (1969) determined radium and thorium distributions
in a series of morainal soils in Bench Valley, California. Thorium accumulated
immediately beneath horizons, containing a high amount of organic material.
The thorium apparently migrated as organic complexes. Radium was distributed
with the uranium in the high organic layers.
3-145
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The ground and surface water sampling and analyses for radium given by
Wruble et al. (1964) and Kaufmann et al. (1975, 1976) for waters of the
Colorado Basin and Grants area, New Mexico, respectively, also illustrate the
tendency of the radium daughter to become separated from uranium due initially
to uranium mining and milling operations, and to continue migrating due to
inherent geochemical differences between uranium and radium.
Laboratory Studies —
226
Several reports have been concerned with the leaching of Ra from
uranium recovery process tailings and wastes (Whitman and Porter, 1958; Anonymous,
1960; Feldman, 1961), and other authors have shown that radium can be leached
from stream sediments, minerals and uranium mill wastes (Starik and Polevaya,
1958; Starik and Lazerev, 1960). The factors that influence Teachability of
226
Ra from uranium mill waste solids and river sediments were investigated by
Shearer (1962) and Shearer and Lee (1964). Less than 1.5 wt% of the river
sediments and mill waste solids were greater than plus 20 mesh size and less
than 30 wt% were minus 140 mesh in particle size. The amounts of radium
leached with distilled water versus time showed that essentially no radium
was leached after 15 min. Diffusion of radium from the interior of the parti-
cles was relatively insignificant. By varying the liquid to solid ratio in
distilled water-solid Teaching equilibria, it was shown that the ratio affects
the amount of radium leached. The largest ratio effect was shown with leach-
ing of the acid leach process tailings, less with alkaline leach process tail-
ings and the least from river sediments. It was demonstrated that sulfate was
present in the waste solids and that the sulfate was easily solubilized. Trace
amounts of barium present led to precipitation of BaSO, and the coprecipitation
of RaSO,. If radium was added prior to the solids-distilled water equilibria,
it too was removed from solution by coprecipitation with BaSO.. One-hundredth
molar solutions were used in leaching equilibria (100 ml/g river sediment) to
determine effects on radium leaching. MgClo. KC1 , NaCl , HC1 and water solu-
tions all leached less than 1 uug of radium while CaCl2 leached 1.2 uug Ra,
SrCl9 6.3 uug Ra and Bad, 30 uuRa. Apparently the radium on river sediments
+2
is in the form of Ra and is exchangeable.
Havlik et al . (1968a) investigated the leaching of Ra from uranium mill
solids and uranium ores. The first report concerned the effects of pH on
3-146
-------
leaching rates. The authors found, Vike Shearer and Lee (1964), that equilib-
rium leaching was rapidly accomplished in 15 to 30 min. Homogenized uranium
ores (540 pCi Ra/g and 85 pCi Ra/g) and milling tailings (14 pCi Ra/g) were
shaken for varying lengths of time as a 3g solid/30 ml solution ratio. The
pH was modified from 1 to 14 with hydrochloric acid, boric acid and sodium
2?fi
hydroxide. At pH 1, 22% of the Ra was liberated. At pH 9, the amount
226
leached had decreased to 2.8%. At pH 13, the amount of leached Ra increased
to 5%.
The second report by Havlik et al. (1968b) studied the leaching of radium
from the same solids as affected by leaching solution composition in addition
to acidity. Unlike the leaching results of Shearer and Lee (1964), Havlik
et al. found that the largest concentrations of radium were leached by IN KC1
and IN NaCl solutions. Bad7. SrCU and CaCU also were used, but with much
less radium liberated. The IN KC1 leached 100% of the radium in mill tailings
and NaCl, 95%. Uranium ore leaching results were lower, showing 22% and 31%
leached by KG! and 14% and 17% leached by NaCl. In all cases, the one normal
salt solutions removed more radium than ten normal salt solutions.
Summary
'~ • i • —
Radium is present as Ra over the normal soil pH range (4 to 8) and
shows little tendency to form complex species (Schubert et al., 1950). Radium
would be expected to substitute for other divalent cations during replacement
-or precipitation reactions (Granger et al., 1961; Granger, 1963). A direct
correlation of cation exchange capacity with adsorption (Arnold and Grouse,
1965; Serne, 1974) and leaching studies with different types of competing
cations (Starik and Polevaya, 1958; Starik and Lazerev, 1960; Shearer, 1962;
Shearer and Lee, 1964; Havlik et al.,.1968b) suggests that an important reac-
tion mechanism for radium adsorption is cation exchange. Radium could be
expected to migrate in much the same manner as strontium.
References
Anonymous. 1960. January, 1960. Summary Report on I. The Control of Radium
and Thorium in the Uranium Milling Industry, II. Radium- 226 Analysis Princi-
ples, Interference and Practice, III. Current Winchester Laboratory Projects.
WIN-112.
3-147
-------
Arnold, W. D. and D. J. Grouse* 1965. Radium Removal from Uranium Mill
Effluents with Inorganic Ion Exchangers. I&EC Process Design and Development.
4:335-337.
Bell, K. G., C. Goodman and W. L. Whitehead. 1940. Radioactivity of Sedimen-
tary Rocks and Associated Petroleum. Bull. Am. Assoc. Petroleum Geol. 24:1529.
Butler, J. N. 1964. Ionic Equilibrium. Addison-Wesley, Reading, MA.
pp. 180-183.
Davis, G. L. 1947. Radium Content of Ultramafic Igneous Rocks. I. Laboratory
Investigation. Am. J. Sci. 245:677.
Evans, R. D. and C. Goodman. 1941. Radioactivity of Rocks. Bull. Geol. Soc.
Am. 52:459.
Evans, R. D., C. Goodman, N. B. Keevil. 1942. Radioactivity: The Earth's
Heat and Geological Age Measurements. IN: Handbook of Physical Constants.
Geol. Soc. Am. Special Papers 36, Sec. 18, p. 267.
Feldman, M. H. 1961. Summary Report, 1959-1961, Winchester Laboratory,
Winchester, Massachusetts and Grand Junction, Colorado. WIN-125.
Granger, H. C. 1963. Radium Migration and Its Effect "on the Apparent Age
of Uranium Deposits at Ambrosia Lake, New Mexico. U.S. Geol. Survey Prof.
Paper 475-B, pp. B-60-B-63.
Granger, H. C., E. S. Santos, B. G. Dean and P. B. Moore. 1961. Sandstone--
Type Uranium Deposits at Ambrosia Lake, New Mexico - an Interim Report.- Econ.
Geol., vol. 56, no. 7, pp. 1179-1209.
Hansen, R. 0. and G. L. Huntington. 1969. Thorium Movements in Morainal
Soils of the High Sierra, California. Soil Science. 108:257-265.
Havlik, B., J. Grafova, and B. Nycova. 1968a. Radium-226 Liberation from
Uranium Ore Processing Mill Waste Solids and Uranium Rocks into Surface
Streams - I. Health Physics. 14:417-422.
Havlik, B., B. Nycova, and J. Grafova. 1968b. Radium-226 Liberation from
Uranium Ore Processing Mill Waste Solids and Uranium Rocks into Surface
Streams - II. Health Physics. 14:423-430.
Kaufmann, R. F., G. G. Eadie and C. R. Russell. 1975. Summary of Ground Water
Quality Impacts of Uranium Mining and Milling in the Grants Mineral Belt,
New Mexico. U.S.E.P.A., Off. Rad. Prog.-Las Vegas Facility. Tech. Note
ORP/LV-75-4.71 p.
Kaufmann, R. F., G. G. Eadie and C. R. Russell. 1976. Effects of Uranium
Mining and Milling on Ground Water in the Grants Mineral Belt, New Mexico.
Ground Water. 14(5):296-308.
3-148
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Parker, V. B., D. D. Wagman, and W. H. Evans. 1971. Selected Values of
Chemical Thermodynamic Properties. U.S. Department of Commerce. NBS Technical
Note 270-6.
Schubert, A. J., E. R. Russell, and L. S. Myers. 1950. Dissociation Constants
of Radium Inorganic Acid Complexes Measured by Ion Exchange. J. Biol. Chem.
185:387-398.
Senftle, F. E. and N. B. Keevil. 1947. Thorium-Uranium Ratios in the Theory
of Genesis of Lead Ores. Trans. Am. Geophys. Union. 28:732.
Shearer, S. D., Jr. 1962. The teachability of Radium-226 from Uranium Mill
Solids and River Sediments. Thesis. Univ. Wisconsin.
Shearer, S. D., Jr. and G. F. Lee. 1964. teachability of Radium-226 from
Uranium Mill Solids and River Sediments. Health Physics. 10:217-227.
Starik, I. E. and K. F. Lazarev. 1960. Effect of Crushing of Minerals on
the Extraction of Radioactive Elements. Radiokhimiya. 11:749-752.
Starik, I. E. and N. I. Polevaya. 1953. The Leachability of ThX and Rd Th
from Minerals. AEC-tr-4208, pp. 108.
Stead, F. W. 1964. Distribution in Groundwater of Radionuclides from Under-
ground Nuclear Explosions. IN: Proc. Third Plowshare Symp. Engineering with
Nuclear Explosives. April 21^23, 1964. TID-7695, p. 127-138.
Vdovenko, V. M. and Yu. V.- Dubasov. 1973... Analytical Chemistry of Radium.
Israel Program for Scientific Translations.
Vinogradoy, A. P. 1959. The Geochemistry of Rare and Dispersed Chemical
Elements in Soils. Consultants Bureau.
Weast, R. C. 1976. Handbook of Chemistry and Physics. The Chemical
Rubber Co.
Whitman, A. and E. S. Porter. 1958. Chemical Stream Pollution from Uranium
Mills. WIN-99.
Wruble, D. T., S. D. Shearer, D. E. Rushing and C. E. Sponagle. 1964. Radio-
activity in Waters and Sediments of the Colorado River Basin, 1950-1963. Rad.
Health Data, 5(ll):557-567.
3-149
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RUTHENIUM
Natural Soil and Rock Distributions
Few data on natural distribution of ruthenium are available. Data avail-
able on ruthenium in rock-forming minerals are listed in Table 3-65. Platinum
group metals, including ruthenium, are known to be concentrated in chromite;
but these are older analytical values and may be much higher than modern
neutron activation ruthenium analyses. There are no ruthenium data on sand-
stones and limestones. However, the ruthenium contents are probably very low.
The meager data on ruthenium content of shales are given in Table 3-66. To
put the ppb values into perspective, 100 ppb amounts to 1 x 10 g of ruthenium
per gram of rock or mineral. The behavior and distribution of natural ruthe-
nium during weathering processes is largely unknown.
TABLE 3-65. ABUNDANCES OF RUTHENIUM IN IGNEOUS ROCKS
AND ROCK-FORMING MINERALS IN ppm
Mineral
01ivine
Chromite, Norway
Chromite, Texas
Chromite, Penn.
Columbite, Norway
Ilmenite
Rutile
Tantalite
Ounite, Russia
Dunite, Russia
Apodunite, Russia
Mica shonkinite
Ru. ppb
10
500
500
500
10
4
10
2
2-30
(5 average)
3-6
(4.5 average)
0.9-4
(2.2 average)
0.5
Reference
Noddack and Noddack, 1931
Qoldschmidt and Peters, 1932
Goldschmi'dt and Peters, 1932
Goldschmidt and Peters, 1932
Noddack and Moddack, 1931
Noddack and Noddack, 193;
Noddack and Noddack, 1931
Noddack and Noddack, 1931
Razln et al., 1965
Razin et al., 1965
Razln et al., 1965
Razin et al., 1965
3-150
-------
TABLE 3-66. RUTHENIUM CONTENTS OF SHALES
Shale Ru, ppb Reference
Kupferschiefer, 3 Noddack, 1936
Mansfeld, Germany
Black Shale (3) 60 Tischendorf, 1959
80
40
Bleached Shale (3) 40 Tischendorf, 1959
5
20
Brief Chemistry
The 16 isotopes of ruthenium include the radioactive fission products
93Ru, 94Ru, 95Ru, 97Ru, 103Ru, 105Ru, 106Ru, 107Ru and 108Ru. The natural,
stable ruthenium isotopes and their abundance are shown in Table 3-67.
TABLE 3-67. ATOMIC PERCENTAGES OF STABLE RUTHENIUM
ISOTOPES (HEATH, 1976)
96 98
Atom % 5.51 1.87
Ruthenium-103 with a 39.6-day half-life and 106Ru with a 367-day half-
life are the only radioisotopes of ruthenium tha't persist long enough to be
of concern in waste disposal.
Ruthenium exhibits several oxidation states varying from Ru(-II) to
Ru(VIII), with Ru(III) and Ru(IV) the most common oxidation states in aqueous
solutions. Ru(II) is less common. For ruthenium, there is little evidence
for simple aquo ions. Nearly all aqueous solutions, whatever the anion, may
be considered to contain complex ions (Cotton and Wilkinson, 1962). Ru(II),
for example, is known to form a unique series of nitrosyl complexes, nitric
oxide complexes, ammonia and sulfur complexes, carbonyl complexes, etc.
In an alkaline medium, Ru(IV) forms the insoluble hydrated oxide Ru02 •
XH20 which does not tend to be reduced to the trivalent state. Ruthenium (III)
hydroxide is readily oxidized in air to Ru(IV) (Ginzburg et al., 1975).
Ruthenium exists in nitric acid solutions either as the Ru(III) nitrosonitrates
3-151
-------
[RuNO(N03)n(H20)5_3]3"n or [RuNO(N03)5_nHn(OH)m • (HgO^P', or as the poly-
meric Ru(IV) aquohydroxo cations [Ru(OH)x(H20)g-x] (N03)4_x (Ginzburg et al.,
1975). Thus, from acidic nitrate solutions, which are the usual form of fuel
reprocessing wastes, ruthenium solutions may be obtained that contain cationic,
anionic and neutral complexes. Hence, the chemical forms of the ruthenium in
a given media are a product of the chemical history of that media.
Solid Phase and Solution Equilibria
Ruthenium is generally present in association with platinum group metals.
The abundance of ruthenium in various rocks and soil forming minerals is sum-
marized in Tables 3-60 and 3-61. Ruthenium also forms discrete solid compounds
such as Ru02, RuO^, Ru(OH)3, Ru(OH)., RuCl3, and RuS2> The relative stability
of the discrete solids in oxidizing conditions (p02/q\ = 0.68 atm) is given
in Figure 3-20. The thermodynamic data for Ru02 and RuO, were obtained from
Wagman et al. (1969). The data for Ru(OH)3 and Ru(OH)4 were selected from
Sillen and Martell (1964). The solids in increasing order of stability for
the conditions outlined in Figure 3-20 are: Ru(OH)3> RuO^, Ru(OH)4, and Ru02
(amorphous, hydrated). RuCU and RuS2 do not appear in Figure 3-20 because
they are too soluble and fall beyond the boundaries of the figure. From
Figure 3-20 it can be ascertained that Ru02 (amorphous, hydrated) would be
most stable over the pH range and up to reducing environments equivalent to
p02 of 50. In extremely reducing conditions (p02 = 80) where HLS (aq) would
be the most stable sulfur species, Ru$2 would maintain lower activities than
Ru02 and hence would be the most stable.
The activity of various solution complexes of ruthenium in equilibrium
with Ru02 (amorphous, hydrated) in oxidizing (p02 0.68) and other conditions
spelled out in the graph are given in Figure 3-21. The thermodynamic data
used to construct Figure 3-21 were selected from several sources including
Sillen and Martell (1964) for RuOH3*, Ru4* and Ru3"1", Baes and Mesmer (1976)
for Ru04(OH)~, and Wagman et al. (1969) for the remaining species. Baes and
Mesmer (1976) indicate that the thermodynamic data on ruthenium complexes'are
either not available or is not very accurate. However, it was decided to
construct Figure 3-21 using the existing data to estimate the nature of the
species that may be present in solution.
3-152
-------
-8
-10
-12
-14
-16
-18
-20
- Ru0;
AND
p02 • SO. 8
uO,
3
10
pH
Figure 3-20.
The relative stability of various ruthenium solids in an oxidiz-
ing soil environment [p02(g) = 0.68 atm], pCl" = pS042~ = 2.5
and pH$ =3.0.
-8
,-12
-16
-20
-24
Ru04°
6 7
pH
10
Figure 3-21. The activity of various ruthenium species in equilibrium with
(amorphous hydrate) in an oxidizing soil environment
3-153
-------
Figure 3-21 indicates that in an oxidizing environment, significant
4+
activity of uncomplexed Ru will be present only in very acidic conditions
(pH <2). Some of the ruthenium species which may be present in solution are
Ru04~, Ru042~, Ru04°, Ru(OH)22+, Ru04(OH)-, RuCl5(CH)2~, Ru4+, and Ru3+.
Among these species, Ru would be predominant in pH <2, Ru(OH)? in pH 2 to
o Cm
5.2, Ru04 in pH 5.2 to 10.5, and Ru04 in pH >10.5. With an increase in
reducing conditions, the activity of Ru would increase, the activity of
o-2-
RuO. , RuO. , RuO* and Ru(LOH would decrease and the rest of the species
would remain unchanged in activity. The nature of the predominant solution
species is highly dependent upon the oxidation-reduction conditions, as is
the most stable solid compound. The predominant ruthenium species are cationic
below pH 5 and anionic above pH 5. Hence, ion exchange on soils and rocks,
mainly cation exchange, would be of importance only below pH 5 under these
conditions.
Experimental Adsorption Results
Rhodes (1957) reported the variation in Ru distribution coefficients
with solution pH on a typical Hanford subsoil. The soil had a cation exchange
capacity of 5 meq/100 g, 2 wt% calcium carbonate, with about 2% clays of
mainly montmorillonite and lesser amounts of chlorite. The Kd values for
106Ru are presented in Table 3-68. It should be noted that the 106Ru was
obtained from Oak Ridge National Laboratory, probably as RuCU in HC1 solution.
TABLE 3-68. RUTHENIUM DISTRIBUTION COEFFICIENTS ON A HANFORD SOIL AS
A FUNCTION OF pH (RHODES, 1957)
pH Kd, ml/g pH Kd, ml/g
40
39
44
92
27
101
752
3-154
0
1.3
2.6
4.7
6.6'
7.4
8.5
0
0
26
101
752
460
274
9.2
10.1
10.4
11.0
11.3
13.0
14.0
-------
The soil adsorption of Ru in a chloride medium cannot be similar to the
adsorption of the nitro- and nitrato complexes of nitrosylruthenium normally
present in acidic nitrate waste solutions. Rhodes and Nelson (1957) make this
same point when they discuss the soil adsorption of ruthenium present in
uranium recovery plant waste. Sampling of groundwater beneath a discontinued
disposal facility showed that the concentration of Ru remained constant for
several months. This indicated little adsorption after the initial break-
through of Ru. Klechkovsky (1956) also investigated Ru adsorption on
soils but did not specify the type of ruthenium or soils used.
Wilding and Rhodes (1963) reported that the addition of 100 ppm citrate
to a synthetic waste-cationic ruthenium solution reduced the Kd value for
ruthenium on Idaho well sediments to 27 ml/g from 472 ml/g with no citrate.
For the same solutions, the reduction on the natural zeolite clinoptilolite
was to a Kd value of 4 ml/g from a Kd value of 48 ml/g with no citrate present.
Addition of 100 ppm EDTA to the synthetic waste (MTR cooling water) lowered
the ruthenium Kd value from about 472 ml/g with no EDTA to 100 ml/g with EDTA.
• o 4.3
Kepak (1966) reported the adsorption Kd for Ru(NO) and Ru on hydrated
3 3
ferric oxide as 8.4 x 10 and 4.8 x 10 , respectively.
Jenne and Wahlberg (1968) examined the sediments present in White Oak
Creek, Tennessee, identifying their mineralogy and determining a Kd for several
elements. The Kd was determined by analyzing the radionuclide content of
White Oak Creek water and sediment, and assuming that an equilibrium between
the two was attained. The water contained 230 pCi Ru/1, while no measur-
able Ru was found in the sediment. It was concluded that the Ru in
White Oak Creek water was not being adsorbed by the contacting sediment.
The form of ruthenium in the Clinch River, downstream of White Oak Creek,
was suggested by Pickering et al. (1966) to be nitrosyl ruthenium hydroxide.
They reported 16 Ci of Ru in Clinch River bed sediment as of 1962.
Lomenick and Gardiner (1965) reported 1,038 ± 88 Ci of Ru concentrated in
a small area of White Oak Lake coinciding with two inlet streams that drain
the waste seepage pit area. White Oak Lake is above White Oak Creek in the
drainage system around Oak Ridge, Tennessee.
3-155
-------
Izrael and Rovinskii (1970) studied the groundwater leaching of radio-
nuclides from Russian bomb blast rubble and debris. They used dialysis,
electrodialysis extraction and ion exchange to examine the chemical state of
several radionuclides. Ruthenium was found to be 19.4% colloidal, 5.1%
cationic and 75.5% anionic. The complex makeup of Ru showed nearly the
same distribution in groundwater solutions of different origin, and had an
insignificant adsorptive capacity.
Arnold and Grouse (1971) described problems associated with radioactive
contamination of copper recovered from ore fractured with nuclear explosives.
The fractured ore was leached with dilute sulfuric acid and the copper was
recovered by replacement of scrap iron as copper cement. The Ru consti-
tuted 42% of the total activity in the leach solution and only the Ru was
removed along with the copper during the iron replacement reaction.
Amy (1971) studied the adsorption and desorption rates of three ruthenium
compounds on acidic and calcareous soils (see Brown, 1976, for a summary of
Amy's work). The three compounds were cationic Ru NON02 (N03)2 (H20)2,
anionic Ru NO(N02)2 OH(H20)2 and RuCl3 which produced a preponderance of
anionic forms.
R. J. Serne of PNL, 1973, determined Ru equilibrium distribution
coefficients with the simulated tank waste compositions given in Table 3-9,
and seven well sediments described in Table 3-10 under antimony. The Ru
was added to the simulated tank solutions as RuCU in HC1 solution and Kd
1 nc
values obtained as shown in Table 3-69. The form of the Ru after addition
to the.simulated tank solutions is unknown. The traced solution was allowed
to remain at 70°C for 48 hours before contact with the sediments.
TABLE 3-69. EQUILIBRIUM DISTRIBUTION COEFFICIENTS BETWEEN SEVEN GLACIO-
FLUVIATILE SEDIMENTS AND THE SOLUTIONS IN TABLE 3-9 CONTAINING
106Ru. EACH VALUE IS AN AVERAGE OF THREE EQUILIBRATIONS.
Solution
merits
1
2
3
4
5
6
7
I_
2.14
2.01
2.18
2.60
1.04
1.03
1.19
II
0.13
0.09
0.16
1.79
0.30
0.24
0.41
III
0.27
0.23
0.18
0.28
0.53
0.40
0
IV
0.60
0.53
0.30
0.03
0.58
0.54
0.15
v_
0.34
0.24
0.29
0.53
0.68
0.44
0.68
•3-156
-------
Migration Results
Field Studies—
Spitsyn et al. (1958) used an alkaline solution (4 to 8 g NaOH/Jl,
200 g NaN03A) and an acidic solution [6 to 8 g HNOg/A, 200 g Al(N03)3/&]
'in field studies to determine radionuclide migration rates through soils.
Practically no ruthenium was adsorbed from acidic solutions, and no anionic
ruthenium was adsorbed by soils. Various radionuclides including ruthenium
were injected into the ground in a mixture and migration followed by wells
along the migration route. Ruthenium traveled about half as far as strontium.
Haney, Brown and Reisenauer (1962) and Brown and Haney (1964) estimated
from the direction and rate of the groundwater flow that 7 to 8 years were
required for Ru contamination in the groundwater to move from Hanford Purex
Plant site to the Columbia River, a distance of nearly 17 miles. The half-
life and travel time of Ru result in less than 1% of the Ru reaching
the Columbia River.
Lomenick (1963) discussed ruthenium flow into the drained White Oak Lake
bed at Oak Ridge from the intermediate waste pits by surface runoff and ground-
water movement. Movement of groundwater in the upper 2 ft of lake bed soil
was from T to 5 ft/day, while at 2 to 5 ft depth, movement was 0.05 to 0.25 ft/
day. The lake bed then (1963) contained about 1200 Ci of ruthenium, with
about 70% of the ruthenium in the top 2 ft, or rapid water movement zone,
presumably originating at the intermediate level waste pits.
Raymond (1964, 1965) investigated the vertical migration and lateral
spread of radionuclides beneath three Hanford waste disposal sites by using
scintillation logging of nearby monitoring wells over a period of time.
g
A disposal site received 7.6 x 10 liters of condensate waste containing
about 4500 gross beta curies, mostly Ru, to 1965. One week after disposal
started, contamination was present to 105 ft. Ruthenium had migrated to
130 ft by June 1958, and the entire soil column was contaminated with ruthe-
nium by June 1959. Ruthenium traveled at nearly the same velocity as the
soil water.
3-157
-------
Brown (1967) gave some profiles of Ru at disposal sites. The Ru
moved at essentially the same velocity as the groundwater, with little or no
soil adsorption. Radiochemical analyses of the groundwater yielded a ruthe-
nium pattern of contamination comparable to the size and shape of the tritium
and technetium contamination patterns. Limits defined by the 1.0 pCi Ru/
3
cm isoconcentration contour in 1966 are given. The disposal site ruthenium
concentration was 1000 pCi Ru/cm of waste.
Brendakov et al. (1969) added Ru to a mountain meadow soil and followed
the vertical migration due to precipitation. The average displacement distance
from the surface was 3 mm a year after amendment, 7 mm 2 years after amendment
and 14 mm 3 years after amendment. The average displacement was defined as:
?* * *
X* = ^P1X1, where pi = fractional % of the cadionuclide in a layer (3 mm thick-
ness) and xi = the depth of the layer in mm. The displacement distance of
Ru was about the same as the Cs and Sb also used in the study. These
results are not typical for Ru migration that ordinarily has been reported
to much exceed the migration velocity of Cs.
In 1973, the 241-T-106 underground liquid high level waste storage tank
at Hanford leaked about 435,000 liters of waste into the surrounding sediments
(Anon., 1973). A series of wells were drilled to ascertain the locations and
movements of the various radionuclides in the sediment column. Monitoring and
core sample analysis showed the relative mobilities of the radionuclides in
the high pH, high salt solution. Ruthenium was the most mobile of the radio-
nuclides, having migrated to 27 m (35 m above the water table). The one yCi
Ru/1 soil isopleth was used to delineate the contaminated zone of about
25,000 m of sediments. One of the reactions of the high pH waste with the
sediment column involved the chemical attack and partial dissolution of soil
minerals (Shade, 1974).
Magno et al. (1970), investigating the radionuclides from the Nuclear
Fuel Services plant in Western New York that were migrating through the
effluent lagooning system, estimated that 70% of the Ru passed through the
lagoon system and into nearby surface streams.
Brookins (1976) reported that ruthenium was retained in the shale that
surrounded the natural reactor site at Oklo in Gabon. Approximately 1.8 bil-
lion years had elapsed since the reactor was active.
3-158
-------
Laboratory Studies—
Nishita et al. (1956) leached ruthenium from soils and clay minerals with
water and a salt solution. The Ru was the most readily leached radionuclide
from the soils with water. The Ru also was the most easily water-leached
from bentonite and kaolinite. The ruthenium tended to be the least exchange-
able in both the soils and the bentonite and kaolinite. A persistent fraction
of the bentonite and kaolinite adsorbed ruthenium was nonexchangeable.
Essington et al. (1965) used fallout material on soils to determine the
leaching effects of water, chelating agents and HC1 on Ru. Chelating
agents (DTPA, CDTA, EDDHA), increased the soluble radionuclides by a small
amo
103
amount as compared to water. The Ru tended to remain in the column. Some
Ru was reported in leachates in CaDPTA-leached soils.
According to Schulz (1965), ruthenium tends to react like manganese,
cobalt, zinc, iron and chromium in reactions in a soil environment. They are
most soluble in acidic conditions and tend to precipitate as hydroxides as
the pH rises. Ruthenium tends to form complexes that are uncharged or nega-
tively, charged and would therefore tend to be quite mobile in soils.
Dlouhy (1967) gave ruthenium elution curves from Casaccia soil that were
given as the volume of water put through the column versus the fraction eluted
at flow velocities of 136 cm/hr and 270 cm/hr. The characteristics of Casaccia
soil were not given. The velocity of the ruthenium down the column was
obtained in relation to water velocity and a theoretical profile on the soil
column derived for field application and estimates. The Kd values for the
3 3
ruthenium ranged from 1.8 to 2.4 cm /g on the soil and 10 to 15 cm /g on the
tuff. The velocity of the ruthenium divided by the velocity of the water was
experimentally determined as 0.11 at the 136 cm/hr water velocity and 0.10 at
the 270 cm/hr water velocity.
Eichholz et al. (1967) studied the distribution of ruthenium between
suspended sediment particles and aqueous solutions. Natural water samples
were characterized and used in the adsorption studies as sources of suspended
solids, as shown in Table 3-70. A fission product mixture was added to the
natural water samples, the system equilibrated and the water and sediments
separated. The adsorption-of Ru is shown in Table 3-71. A portion of the
3-159
-------
6Ru was associated with the suspended sediments, though to a lesser extent
than the other radionuclides with the exception of I.
Source
Colorado River,
Utah
Camp McCoy,
Wisconsin
Bayou Anacoco,
Lousiana
Lodgepole Creek
Nebraska
Chattahoochee
River, Georgia
Billy's Lake
Georgia (swamp).
TABLE 3-70.
Suspended
Solids, ppm
229
12
24
PROPERTIES OF NATURAL WATER SAMPLES
(EICHHOLZ ET AL., 1967)
Dissolved Conductivity
Solids, ppm pH pmhos
350
60
63
68
7.5
6.9
6.2
4.2
540
80
60
965
131
200
31
6.8
7.3
300
45
45
Solids
Composition
95% quartz,
5% calcite,
feldspar, illite,
kaolinite
30% quartz +
feldspar,
6% kaolinite, 24%
illite
30% quartz,
2% kaoUnite,
6% smectite
20% illite,
80% smectite
33% quartz,
44% kaolinite
23% illite
2% quartz,
balance was
amorphous
and/or organic
TABLE 3-71. ADSORPTION OF 106Ru ON SUSPENDED SOLIDS IN NATURAL WATERS
(EICHHOLZ ET AL., 1967)
Source
Colorado River
Camp McCoy
Bayou Anacoco
Chattahoochee River
Billy's Lake
-EL
7.5
6.9
6.2
7.3
4.2
106
Ru Adsorption, %
25.0
5.7
-0
10.4
0
It is obvious that ruthenium adsorption was a function of pH of the water.
The higher percentages of adsorbed ruthenium are directly correlated with
water pH.
3-160
-------
Nishita and Essington (1967) continued work on Ru migration in dif-
ferent types of soils, comparing the relative movement of Zn, Sr, Y,
Ru, Cs, Ce and Pm. Irrespective of the soil type, water leaching
resulted in the greatest movement for Ru.
Bovard et al. (1968) reported that the free functional groups (carboxylic,
phenolic) were the cause of organic matter complexing and fixation of radio-
nuclides. For humic acids, the mean order of complexing was Ce>Fe>Mn>Co>Ru>^
Sr>Cs. In the case of fulvic acids, no universal order existed. Two types of
soils were used in the studies, a calcareous alluvial soil and a humoferrugi-
nous podzol. Leaching of fission products, particularly ruthenium, by rain-
water occurred essentially in the form of fulvates.
Amy (1971) studied the migration rates of different ruthenium forms in
calcareous and acidic soil in columns 30 cm in height by 4.5 cm in diameter.
The nitrbdinatrato complexes of ruthenium, RuNONO,(NO^)., (H.,0),, were much
137
more mobile than cesium in calcareous soils. About 98 to 100% of the Cs
remained in the top 4.5 cm of the soil columns, while 38 to 80% of the nitro-
dinatrato complexes were found in the top 4.5 .cm with distribution down the
column to 1 to 2% of the ruthenium in the last column section. From 2 to 33%
of this-ruthenium form passes through the column. For comparison, ruthenium
as RuCl3 adsorbs only 10 to 50% in the top 4.5 cm of the soil column, and the.
dinitro complex, RuNOtNOgJgOhKHgOjg, only 11 to 36%. The fraction from RuCl3
that migrates down the entire column length varied from 3 to 68%, depending
on the soil type. In summary, in calcareous soils maximum ruthenium mobility
was attained when the ruthenium was in the form of an anionic complex in per-
meable soils with a sandy texture and low in organic material. Three forms
were identified in the column work referred to as A, B, and C. Form A was
cationic with a very rapid adsorption rate. Form B was adsorbed very slowly by
the soil and was essentially adsorbed to the same degree throughout the column.
It may be an anionic form that is evolving toward a. cationic form. Form C
was present only at the bottom of the soil column if flow rates are very slow.
This also was predominately an anionic form also changing toward a more easily
adsorbed form. The RuNONO^NO-j^HgO^ contained Forms A and B in about equal
proportions. The RuCl3 contained 15% of Form A, 75% of Form B and 15% of
Form C. The RuCNOg^OHMH?^ contained only 5* °f P0'™ A and 85% and
3-161
-------
Forms B and C, respectively. Less than 2.5% of any of the above forms passed
through the Barthelasse soil which was rich in organic matter. The acidic
soils always adsorbed less ruthenium in any form than the calcareous soils.
However, like the calcareous soils, the adsorption of ruthenium decreased with
the increasing proportion of anionic forms. The calcareous soils, in order to
adsorb 50% of the ruthenium, required from 1 to 4 min for the nitrodinatrato
complexes, 100 to 300 min for RuCl3 and 200 min to 1 day for the dinitro com-
plexes. The acidic soils, on the other hand, required 7 to 15 min to adsorb
50% of the ruthenium from the nitrodinatrato complexes, 500 min to 1 day from
RuCl3 and 1 to 3 days from dinitro complexes. There were indications that the
organic matter in acidic soils favored the mobility of all forms of ruthenium.
Champlin and Eichholz (1976) have shown that finely particulate soil
materials, such as kaolinite, can adsorb a variety of trace radionuclides
including Ru and adhere to the soil matrix surface by Van der Waal's forces
(Champlin, 1969). Further, these fine particulates were mobilized by laundry
detergents and softening agents, so that they were transported quickly through
the. soil column without reaction with contacting soils.
Summary
3+
Adsorption of Ru peaks at pH 6.6 and 14, with a definite correlation
between pH and ruthenium Kd for all ruthenium species (Rhodes, 1957; Eichholz
et al., 1967; Amy, 1971). Adsorption of all ruthenium species is low at low
pH, and little if any adsorption of anionic species occurs at any pH (Spitsyn
et al., 1958; Haney et al., 1962; Brown and Haney, 1964; Raymond, 1964, 1965;
Brown, 1967; Magno et al., 1970; Schulz, 1965). Some instances of migration
on suspended solids have been noted (Champlin and Eichholz, 1976; Eichholz
et al., 1967). Hydrated ferric oxide is a good adsorbent for ruthenium (Kepak,
1966). Low ruthenium Kd values have been reported in high salt solutions
(Anonymous, 1973; Serne, 1973), probably as a result of nitrate complexintj:
Migration of most ruthenium complexes through the soil column is relatively.
rapid (Pickering et al., 1966; Izrael and Rovinskii, 1970; Amy, 1971). Most
of the thermodynamic data for the ruthenium complexes do not exist. Thus an
estimation of their relative importance in waste disposal operations or spills
is difficult, requiring experimental studies for evaluation in^each case.
3-162
-------
References
Amy, J. P. 1971. Contribution to the Study of the Migration of Ruthenium-106
in Soils. RFP-Trans-140.
Anonymous. 1973. 241-T-106 Tank Leak Investigation. ARH-2874.
Arnold, W. D. and D. J. Grouse. 1971. Radioactive Contamination of Copper
Recovered from Ore Fractured with Nuclear Explosions.
Avtokratova, T. D. 1963. Analytical Chemistry of Ruthenium. Israel Program
for Scientific Translations.
Baes, C. F., Jr. and R. F. Mesmer. 1976. The Hydrolysis of Cations.
John Wiley and Sons, New York.
Bovard, P., A. Grauby and A. Saas. 1968. Chelating Effect of Organic Matter
and Its Influence on the Migration of Fission Products. IN: Isotopes and
Radiation in Soil Organic Matter Studies. CONF-680275, pp. 471-495.
Brendakov, V. F., A. V. Dibtseva, V. I. Svishcheva, and V. N. Churkin. 1969.
Vertical Distribution and Evaluation of the Mobility of Fission Products in
Certain Soil Types of the Soviet Union. AEC-tr-7030, pp. 152-170.
Brookins, D. G. 1976. Shale as a Repository for Radioactive Waste: The
Evidence from Oklo. Environmental Geology. 1:255-259.
Brown, D. J. 1967. Migration Characteristics of Radionuclides Through Sedi-
ments Underlying the Hanford Reservation. ISO-SA-32.
Brown, D. J. and W. A. Haney. 1964. Chemical Effluents Technology Waste
Disposal Investigations July-December, 1973 - The Movement of Contaminated
Ground Water from the 200 Areas to the Columbia River. HW-80909.
Brown, K. W. 1976. Ruthenium: Its Behavior in Plant and Soil Systems.
EPA-600/3-76-019.
Champlin, J. B. F. 1969. Transport of Radioisotopes by Fine Particulate
Matter in Aquifers. PB-232179.
Champlin, J. B. F. and G. G. Eichholz. 1976. Fixation and Remobilization of
Trace Contaminants in Simulated Subsurface Aquifers. Health Physics.
30:215^219.
Cotton, F. A. and G. Wilkinson. 1962. Advanced Inorganic Chemistry.
Interscience Publishers.
Dlouhy, Z. 1967. Movement of Radionuclides in the Aerated Zone. IN:
Disposal of Radioactive Wastes into the Ground. IAEA-SM-93/18, pp. 241-249.
Eichholz, G. G., T. F. Craft, and A. N. Galli. 1967. Trace Element Fractiona-
tion by Suspended Matter in Water. Geochim. et Cosmochim. Acta. 31:737-745.
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Essington, E. H., H. Nishita and A. J. Steen. 1965. Release and Movement of
Radionuclides in Soils Contaminated with Fallout Material from an Underground
Thermonuclear Detonation. Health Physics. 11:689-698.
Ginzburg, S. I., N. A. Ezerskaya, I. V. Prokof'eva, N. V. Fedorenko,
V. I. Shlenskaya, and N. K. Bel'skii. 1975. Analytical Chemistry of Platinum
Metals. Translated by N. Kaner. Israel Program for Scientific Translations.
Goldschmidt, V. M. and C. Peters. 1932. Zur Geochemie der Edelmetalle.
Nachr. Adad. Wiss. Gottingen, Math-Physik. K1.IV:377.
Haney, W. A., D. J. Brown, and A. E. Reisenauer. 1962. Fission Product
Tritium in Separations Wastes and in the Ground-Water. HW-74536.
Heath, R. L. 1976. Table of Isotopes. IN: Handbook of Chemistry and
Physics. R. C. Weast (ed.), Chemical Rubber Co.
Izrael, Yu. A. and F. Ya. Rovinskii. 1970. Hydrological Uses of Isotopes
Produced in Underground Nuclear Explosions for Peaceful Purposes. UCRL-
Trans-10458.
Jenne, E. A. and J. S. Wahlberg. 1968. Role of Certain Stream-Sediment
Components in Radioion Sorption. U.S.G.S. Professional Paper 433-F.
Kepak, F. 1966. Sorption of the Radioisotopes 35S, 131I and 106Ru on
Hydrated Oxides in Laboratory Columns. Collection Czechoslov. Chem. Commun.
31:3500-3511.
Klechkovsky, V. M. 1956. On the Behavior of Radioactive Fission Products in
Soil. Their Absorption by Plants and Their Accumulation in Crops. AEC-TR-2867.
Lomenick, T. F. 1963. Movement of Ruthenium in the Bed of White Oak Lake.
Health Physics. 9:835-845.
Lomenick, T. F. and D. A. Gardiner. 1965. The Occurrence and Retention of
Radionuclides in the Sediments of White Oak Lake. Health. Physics. 11:567-577.
Magno, P., T. Reavey, and-J. Apidianakis. 1970. Liquid Waste Effluents from
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Nishita, H. and E. H. Essington. 1967. Effect of Chelating Agents on the
Movement of Fission Products in Soils. Soil Science. 102:168-176.
Nishita, H., B. W. Kowalewsky, A. J. Steen, and K. H. Larson. 1956. Fixation
and Extractability of Fission Products Contaminating Various Soils and Clays.
Soil Science. 81:317-326.
Noddack, I. 1936. Uber die Allgegenwart der Chemischen Elementer. Angew.
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Noddack, I. and W. Noddack. 1931. Die Geochemie des Rheniums. Z. Physik.
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R. W. Andrew. 1966. Radioactivity in Bottom Sediments of the Clinch -
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Raymond, J. R. 1964. Investigation of the Disposition and Migration of Gross
Gamma Emitters Beneath Liquid Waste Disposal Sites. HW-81746, pp. 4.32-4.36.
Raymond, J. R. 1965. Cesium and Strontium Distribution Beneath Liquid Waste
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Schulz, R. K. 1965. Soil Chemistry of Radionuclides. Health Physics.
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ARH-CD-176.
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3-165
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STRONTIUM
Natural Soil and Rock Distribution
Average strontium contents of igneous and sedimentary rocks are given in
Tables 3-72 and 3-73. Note that the strontium contents are often high when
compared to the trace concentrations of strontium that may come in contact
with them in waste solutions. Vinogradov (1959) reports the average strontium
90
content of soils as 300 ppm. The 300 ppm present as Sr in a waste solution
90
would be 42.3 Ci Sr/1, for example. Limestone and other calcium-rich sedi-
ments and sedimentary rocks tend to be the highest in strontium content due
to limited substitution of strontium for calcium in calcite. Strontium car-
bonate crystallizes in an aragonite type structure when present in sufficiently
high concentrations.
TABLE 3-72. AVERAGE STRONTIUM CONTENT OF IGNEOUS ROCKS
(TUREKIAN AND WEDEPOHL, 1961) IN ppm
Granitic
Ultramafic Basaltic High Ca Low Ca Syenite
1.0 465 440 100- 200
TABLE 3-73. AVERAGE STRONTIUM CONTENT IN SEDIMENTARY ROCKS-AND SEDIMENTS
(TUREKIAN AND WEDEPOHL, 1961) IN ppm
Pelagic
Sandstone Limestone Carbonate Clay
20 610 . 2000 180
Brief Chemistry
84
The four naturally-occurring isotopes of strontium include Sr (0.55%),
QC Q7 00 Q7
bbSr(9.75%), b/Sr (6.96%) and abSr (82.74%). The isotope b/Sr can be locally
87
variable in abundance because it is a daughter of Rb. The other isotopes
of strontium between Sr and Sr are radioisotopes. Only Sr fission
90
product (half-life 28.1 years), decaying to Y, is of sufficiently high yield
and long half-life to be of concern in waste disposal operations. The
3-166
-------
on
radionuclide Sr also is obtained in high yield but the half-life is too
85
short (52 days) to create a persistent disposal problem. The isotope Sr
(half-life 64 days) is not a fission product, and is usually not encountered
except in laboratory waste solutions.
Strontium (II) is the only oxidation state to be encountered in soil-
+?
contacted solutions. According to Ahrens (1952) the ionic radius of Sr w is
1.12 A, very close to that of Ca+2 at 0.99 A and Ba"1"2 at 1.34 A. Strontium
tends to follow calcium in weathering and other geochemical processes with
some exceptions. For example, when present in sufficient concentrations,
strontium does form a carbonate that is not isostructural with CaC03. The
other main strontium ore mineral, celestite (SrSO,), is isostructural with
its calcium counterpart, anhydrite (CaS04). However, limestone or calcite
(CaCQ3) can allow the substitution of several hundred parts per million
strontium before there is any tendency for strontianite (SrC03) to form. As
the limestone recrystallizes, there is a tendency for the strontium content
to decrease.
There is little, if any, tendency for strontium to form complexes with
inorganic ligands. Izrael and Rovinskii (1970), for example, studied by
electrodialysis the chemical state of strontium leached by groundwater from
rubble produced in a nuclear explosion. They found that 100% of the strontium
+2
was Sr , with no colloidal or anionic strontium present in the leachate.
Solid Phase and Solution Equilibria
Strontium is an alkaline-earth and forms several salts. Thermodynamic
data for various strontium compounds [SrO, Sr(OH)2, SrF2, SrClg, SrCl2-2H20,
•SrCl2-3H20, SrCl2-H20, SrHP04, Sr3(P04)2, SrS04, Sr(N03)2, SrC03> SrSi03,
SrSi04, SrZr03] were considered in order to determine their stability. Except
for SrSi03, SrHP04, Sr3(P04)2, SrC03, and SrS04, the strontium solid phases
are too soluble to show in Figure 3-22. The thermodynamic data for all of
these compounds except Sr3(P04)2 were obtained from Parker et al. (1971).
The data for Sr3(P04)2 were selected from Sillen and Martell (1964). In an
acidic environment, most of the strontium solids will be highly soluble, and
2+ -4
if the activity of Sr in solution exceeds 10 moles/liter, SrS04 can
3-167
-------
precipitate and would be a stable phase. However, in alkaline conditions,
SrC03 would be the stable phase and the compound that may maintain strontium
activity in soi.l solutions.
CVi
r-Sr3(P04)2(V&G)
Sr3(P04)2(OCP)
2+
Figure 3-22. The relative stability of various strontium solids at pCa =
pS042' = 2.5, pH4Si04 = 3.1, and pC02(g) = 1.52 atm in equilib-
rium with Variscite and Gibbsite (V and G), Dicalcium Phosphate
Dihydrate (DCPD) and Octacalcium Phosphate (OCP). Redox
potential has little effect on strontium solubilization.
The activity of various solution species in equilibrium with SrC03 and
o M
at anionic activities of pCl" = pS04 " = 2.5, pN03 » 3.0, pF = 4.5,
pH2P04" = 5 and pC02/ > = 1.52 atm are plotted in Figure 3-23. The thermo-
dynamic data for SrN03^ and SrP04~ were obtained from Sillen and Martell
(1964). All the other data were obtained from Parker et al. (1971). Although
strontium forms complexes with the various anions listed above, the solution
complexes do not contribute significantly to the total strontium activity in
solution. Sr2+ would be the dominant solution species over the environmental
3-168
-------
pH range of interest and would be expected to be adsorbed by an ion exchange
reaction. Izrael et al. (1970) have indicated that strontium was likely to be
100% cationic and that the principal reaction with soils and rocks is ion
exchange.
co
01
o
The activities of various strontium species in equilibrium with
SrC03(s) in the soil at pCl = pS042' = 2.5, pN03~ = 3.0, pF" =
4.5, pC02(g) « 1.52 atm and pHgPCty" = 5.0.
Figure 3-23.
Experimental Adsorption Results
QQ
Amphlett and McDonald (1956) used Sr to study cesium and strontium
removal by a soil made up of montmorillonite, illite and kaolinite. The
+ Sr"1"2 H> 2 Cs"1"
equilibrium constant for the reaction 2
determined as 9.3 x 10 .
Srsoil was
Rhodes (1957) showed the effect of system pH on the adsorption of stron-
tium from a distilled water solution by a fluviatile Hanford soil. A soil
paste pH was 8.6, and the calcite content was 2 wt%. The clay fraction of
3-169
-------
less than 2 ym was 2 wt% and consisted of primarily montmorillonite, with some
mica and chlorite. The system was pH adjusted with NaOH or HC1. The strontium
Kd rose from 5 ml/g at pH 6 to 30 ml/g at 8, and 120 ml/g at pH 10. Above
pH 10, the sodium added in the.NaOH used for pH adjustment began to compete
for exchange sites with the strontium, and strontium adsorption began to level
off. In 4M NaNO.,, strontium adsorption was much less affected by pH. At pH 8,
for example, strontium Kd was about 5 ml/g and rose to about 10 ml/g at pH 10.
Rhodes and Nelson (1957) determined the effects of strontium concentra-
tion on the strontium distribution between soil and synthetic uranium recovery
plant scavenged waste. The soil used here was the same as that used in the
above study (Rhodes, 1957). On a log-log plot of strontium concentration in
the waste versus strontium concentration on the soil, a straight line relation-
ship resulted indicating ion exchange as a reaction mechanism.
Prout (1958, 1959) showed the effects of pH and strontium concentration
on strontium adsorption by Savannah River soil. This soil was 80% sand and
20% clay, chiefly kaolinite. The results were very similar to those of Rhodes
(1957) in that a maximum strontium adsorption was reached at about pH 10,
although this maximum was much higher (Kd = 700 to 800 ml/g). The effects of
competing cations including several sodium nitrate, concentrations were also
given. The 30% NaN03 solution is about the same as Rhodes (1957) 4M NaN03
solution, and the results also are comparable in strontium distribution
coefficients. The amount of competing sodium is apparently so high that the
soil differences are of lesser importance.
Spitsyn and Gromov (1958) showed that the adsorption of strontium by
montmorillonite was an ion exchange process. Various competing cations were
added to the strontium-clay system and the strontium distribution coefficient
_2
redetermined. These authors also added small amounts of anions such as CO,
-2
and C,0A that were said to form radiocolloids with strontium. They believed
-5
the enhanced removal of strontium reported by McHenry (1955), when 3 x 10 to
-4 -3
3 x 10 M PO* was added to wastes contacting calcareous soils, could be
attributed to the formation of precipitates as a result of the natural calcium
on soil exchange sites reacting with the added phosphate ions. This may be
true at some phosphate concentrations, but later study also demonstrated a
replacement reaction (Ames, 1959) where calcite reacted with the phosphate to
form an apatite.
3-170
-------
McHenry (1955, 1958) extensively studied strontium adsorption by a Hanford
soil. The Hanford soil had an exchange capacity of 5.90 ± 0.16 meq/100 g, a
clay content of 1.15%, a calcite content of 1.50% and a paste pH of 8.4. The
soil was calcium-based before use. A comparison of the Hanford composite soil
with others is made in Table 3-74 as strontium Kd values from distilled water.
The effect of initial strontium concentration on strontium Kd values is also
given. The influence of am'ons on strontium Kd values is shown in Table 3-^75.
It is probable that replacement of the carbonate in the calcite by oxalate and
*3 9 *3
phospha'te ions was underway at the 3 x 10" M £7^4 an(* ^4 concentrations,
leading to enhanced removal of strontium during the reaction (Ames et al., 1958).
TABLE 3-74. STRONTIUM Kd VALUES FOR WESTERN SOILS
(McHENRY, 1958)
Soil
Ringold, E. Wash.
Warden, E. Wash.
Bowdoin, E. Mont.
Hall, C. Nebraska
Hanford Composite
Kd, ml/q
400
270
135
600
70
CEC,
meq/100 g
34.0
6.0
13.6
26.3
5.8
Mechanical Analysis, %
>50 um
18.6
78.1
48.5
5.9
84.6
50-2 um
39.0
18.9
34.6
60.6
• 11.9
<2 yun
42.4
3.0
16.9
33.5
3.5'
Paste
7.2
7.3
8.3
6.5
8.3
TABLE 3-75. THE INFLUENCE OF ANIONS ON THE STRONTIUM Kd IN THE HANFORD
COMPOSITE SOIL. ANIONS, AS SODIUM SALTS, WERE PRESENT IN
CONCENTRATIONS 10 TIMES AS GREAT AS THE STRONTIUM CATIONS.
(McHENRY, 1958)
Kd, ml/q
S/2,
3 x
3 x
3 x
3 x
10
10
10
10
M
-4
-5
-6
-7
NOJ
18.
49.
65.
65.
Cl"
0
8
6
6
25
55
70
72
.7
.9
.3
.9
C2042
574
98.
65.
67.
9
6
9
sa
24
61
72
78
-2
.3
.4
.9
.6
PQ43
847
115
70.3
75.7
Klechkovsky et al. (1958) determined strontium distribution coefficients
on a number of Russian soils, and the effects on strontium Kd of competing
Ca and K . The com)
their concentrations.
+2
Ca and K . The competing cations lowered the'strontium Kd in proportion to
3-171
-------
Sorathesn et al. (1960) obtained some strontium distribution coefficients
on Clinch River sediments that were in the ten thousands, much higher than for
standard clays equilibrated with the same solution. The difference was thought
to be due to the organic material in the river sediment. Benson (1960)
reviewed the strontium work with soils at Hanford up to 1960.
Baker and Beetem (1961) studied the distribution coefficients for carrier-
free strontium on an Alaskan wind-blown silt with various concentrations of
competing cations which were to be expected at the site. One or two of the
Kd values exceed the experimental limits by large amounts, but the remaining
values conform to a straight line plot of strontium Kd versus the inverse sum
of the competing calcium and magnesium concentrations as required by ion
exchange theory (Kaufman, 1963).
Cohen and Gailledreau (1961) investigated strontium adsorption on Saclay
soil consisting of 45% montmorillonite, 45% kaolinfte and 10% illite. The
cation exchange capacity was 210 meg/100 g. Graphs were presented of trace
strontium-macrocalcium, trace strontium-macrosodium and trace strontium-
macropotassium systems. Some work was performed on three-component ion
exchange systems. It was shown that the Kd values plotted versus the stron-
tium concentration for the above Saclay ~soil yielded-a straight line on log-
log scales as required by ion exchange theory.
Goldberg et al. (1962) measured the strontium distribution between
ground Rainier Tuff and a synthetic groundwater solution at 1-day and 7-days
equilibration times. There was marked time dependency on strontium adsorp-'
tion by the 100 to 200 mesh tuff. Comparable solutions yielded a 7-day
strontium Kd of 3000 ml/g and a 1-day strontium Kd of 500 ml/g.
Angelo et al. (1962) measured strontium Kd values on 32 to 80 mesh
basalt and obtained values ranging from 16 ml/g to 135 ml/g, depending on
the amount of salt in the contacting solution. The lesser values correlated
with the higher salt contents.
Janzer et al. (1962) examined strontium adsorption on Project Gnome
rock samples. The strontium Kd ranges obtained are listed in Table 3-76.
The rock types are described in Table 3-24. One sample part was equilibrated
with non-traced solution, with the tracer added later. The five equilibrating
3-172
-------
solutions varied from 1,000 to 45,000 mg/1 in dissolved solids and each experi-
ment was run in triplicate. The strontium Kd values are given as minimums for
the highest dissolved solids to maximums for the lowest dissolved solids. A
membrane technique was used to separate solids and solution.
TABLE 3-76. SUMMARY OF STRONTIUM Kd VALUES
(JANZER ET AL., 1962)
Kd, ml/g
Sample
45
95
320
331
350
365
380
479
510
5-26
552
620
653
680
Baetsle et al.,-(1962, 1963) gave the range in strontium Kd values with
soils at Mol, Belgium, that consisted chiefly of quartz sand as 1.7 to 38 ml/g
at pH 7.7 with tap water. Dynamic or column Kd values were compared to static
determinations. Problems were encountered in obtaining adequate Kd values
derived from column breakthrough data because physical flow phenomena were
superimposed on the effects due to radionuclide adsorption reactions. Batch
adsorption data were easier to obtain and free of problems associated with
column physical flow phenomena.
Beetem et al. (1962) determined strontium Kd values on granodiorite from
the Climax stock with synthetic groundwater. The 100 to 200 mes'h samples
3-173
Minimum
8
8
10
3
8
9
6
8
5
8
7
6
0.6
0.7
Maximum
20
23
29
22
16
16
12
17
14
17
15
14
9
11
-------
were equilibrated for 72 hr and yielded strontium Kd values of 4 to 9 ml/g.
The 0.5 to 1.0 mm granodiorite was equilibrated for 720 hr and yielded Kd
values of from 11 to 23 ml/g, probably due to the longer contact time.
Kokotov et al. (1961, 1962) gave Kd values for strontium on six Russian
soil types as a function of pH. The replacement series Sr > Ca > Mg > K s NH.
> Na was given for strontium competing cation efficiency.
Stead (1963) referred to strontium Kd values of 5 to 14 ml/g on a dolo-
mite from the Rustler formation when the groundwater had 4250 mg/1 of total
dissolved solids. The average Kd was 10 ml/g.
Berak (1963) used the same solution and conditions shown in Table 3-20
under cesium to determine strontium Kd values of a number of rocks and min-
erals. The diameter of the solid particles was 0.1 to 0.2 mm. The concen-
tration of carrier strontium was 1 x 10 M as SrCl^- The strontium distribu-
tion coefficients are given in Table 3-77. Strontium discrimination factors
were determined on several rocks and minerals by using a solution containing
10"5M Sr (NO,), + 90Sr and 2.5 x 10"3M Ca"1"2. After equilibration, 90Sr was
J +2
determined by radioassay and Ca by complexometric titration. The fraction
90 +2
of Sr on the solid divided by the fraction of Ca retained on the solid
constitutes the strontium discrimination factor at the strontium and calcium
concentrations and ratios used. The results are listed in Table 3-78.
TABLE 3-77. STRONTIUM DISTRIBUTION COEFFICIENTS FOR SEVERAL ROCK,
MINERAL AND SOIL TYPES
Rock
Alluvium,
Central Nevada
Silty Clay, Idaho
Silty Clay, Idaho
Clinch River, Tenn.
bottom sediments
Sr
Kd
, mg/1 Condition
48-2454 500-4000
545
1690
1438
659
3537
4024
240
220
:
-
*
1
3
7
1
3
7
Quartz ,
illite,
apatite
Quartz ,
illite,
apatite
hr
days
days
hr
days
days
pH 6
pH 9
ym
Reference
Nork et al.,
1971
Wilding and
Rhodes, 1963
Wilding and
Rhodes, 1963
Carrigan et al.,
1967
3-174
-------
TABLE 3-77. (continued)
Reck
Sr Kd, mg/1
Condition
Reference
Sandstone, fine,
light gray
Shaley siltstone,
carbonaceous, black
Sandstone, fine,
light gray
Sandstone, very fine,
silty, dark gray
1.26
1.13
1,88
8.32
9.41
9.56.
1.37
1.36
2.08
7.76
8.37
9.79
4000 ym
500 ym
<62 ym
4000 ym
500 ym
<62 ym
4000 ym
500 ym
<62 um
4000 ym
500 ym
<62 ym
Sokol ,
Sokol ,
Sokol ,
Sokol ,
1970
1970
1970
1970
Shaley limestone,
New Mexico
Sandstone,.
New Mexico
Alluvium, INEL, Idaho
Tuff, Rainier
Mesa, Nevada
Tuff, Rainier
Mesa, Nevada
Tuff, Rainier
Mesa, Nevada
Tuff, NTS, Nevada
Limestone,
Yucca Flat, Nevada
Granodiorite,
Climax stock
Granite, Central
Nevada
Basalt, Buckboard
Mesa, Nevada
Basalt, Amchitka
Basalt, Amchitka
8.32
1.37
>4000 ym
>4000 ym
7.2-10.5 (Laboratory)
40 (Field)
260
2070-3480
1700=4300
4000
0.19
4-9
11-23
1.7
16-135
220
>400 ym
100-200 mesh
>4000 ym
100-200 mesh
0.5-1 mm
>4000 ym
32-80 mesh
500-4000 ym
1220 500-4000 ym
1.1 (seawater) 500-4000 ym
Nork and
Fenske, 1970
Nork and
Fenske, 1970
Schmalz, 1972
Nork and
Fenske, 1970
Goldberg et al.,
1962
Kaufman,
1963
Stead, 1964
Nork and
Fenske, 1970
Beetern
et al., 1962
Nork and
Fenske, 1970
Angelo et al.,
1962
Essington and
Nork, 1969
Essington and
Nork, 1969
3-175
-------
Rock
TABLE 3-77. (continued)
Sr Kd. mg/1 Condition
Reference
SUPERSATURATED ROCKS (contain quartz)
Rhyolites,
Czechoslovakia
Volcanic Glasses, East Slovakia
.
Central Slovakia
Iceland
Rhyodacites, East Slovakia
Dacites, Slovakia
Quartz Porphyries,
Central Bohemia
Bohemia
North Bohemia
SATURATED ROCKS (contain no quartz
Trachytes, North Bohemia
Slovakia
Andesites, Central Slovakia
56 0.1-0.2 mm
15
7
6
33
30
25
41
12
4
12
9
9
21
12
31
37
0
7 0.1-0.2 mm
32
1
11
13
3
6
11
16 0.1-0.2 mm
43
21
27
18 0.1-0.2 mm
8
47
6
8
or feldspathoids)
39 0.1-0.2 mm
12
11 0.1-0.2 mm
13
Berak, 1963
Berak, 1963
>•
Berak, 1963
Berak, 1963
Berak, 1963
Berak, 1963
3-176
-------
TABLE 3-77. (continued)
Rock
Pyroxenic, East Slovakia
Garnetic, East Slovakia
Altered,
Pyroxenic,
Propylletized
Granul. pyrox., vitreous
Spillites, albitized
Proterozoic basalt, -
Central Bohemia
Diabases, Bohemia
North Moravia
Melaphyres, basalt
North Bohemia
Slovakia
Basalts, North Bohemia
Vitreous
Sr Kd, mg/1
17
13
37
16
30
32
21
48
20
50
Condition
Reference
8
0
1
10
25
32
30
39
45
1
212
35
5
10
187
6
22
80
33
32
UNDERSATURATED ROCKS (Nonfeldspathoidal)
Picrites, North Moravia
64
64
43
85
85
64
51
108
9
0.1-0.2 mm Berak, 1963
0.1-0.2 mm Berak, 1963
0.1-0.2 mm Berak, 1963
0.1-0.2 mm Berak, 1963
0.1-0.2 mm Berak, 1963
3-T77
-------
Rock
TABLE 3-77. (continued)
Sr Kd, mg/1 Condition
Reference
Teschenites with olivine,
North Moravia
.
Picrites, North Moravia
Peridotite, North Moravia
Pyroxenite, NW Bohemia
Monchi quite with olivine,
North Moravia
Ouachitite, with apatite
North Moravia
Olivine basalts, N. Bohemia
East Slovakia
Central Slovakia
North Bohemia
UNDERSATURATED ROCKS (with
Phono! ites, North Bohemia
Tephrites, North Bohemia,
chabazite
leucite
chabazite
nepheline
nepheline-leucite
..leucite
leucite
Basanites, North Bohemia,
nepheline-leucite
Nephelinites, Mel i lite,
NW Bohemia
leucite
olivine-nepheline basalt
74
30
22
38
25
46
92
13
35
0
0
35
7
52
321
322
324
334
feldspathoids)
15
• • 7
140
75
86
26
44
143
83
65
80
32
0
22
38
59
__.0
5
22
28
60
0.1-0.2 mm Berak, 1963
0.1-0.2 mm Berak, 1963
0.1-0.2 mm Berak, 1963
0.1-0.2 mm Berak, 1963
. . ... . .
0.1-0.2 mm Berak, 1963
0.1-0.2 mm Berak, 1963
0.1-0.2 mm Berak, 1963
3-178
-------
TABLE 3-77. (continued)
Rock
Leucitites, NW Bohemia
Sr Kd, mg/1
Condition
Reference
nepheline
Melitites, olivine-hauyne basalt
MINERALS
Quartz
Agate
Chalcedony
Opal
Olivine
Chondrodite
Thaumasite
Leucophanite
Zircon
Topaz
Kyanite
Sillimanite
Staurolite
Garnet, almandine
Garnet, grossular
Garnet, andiadite
Garnet, uvarovite
Hydrogarnet
Vesuvianite
Sphene
Rinklte
Axinite
Hemimorphite
Beryl
Dioptase
Tourmaline
Woliastonite
Rhodonite
Hypersthene
Aegirine
Augite
Diopside-
' Enstatite
Jeffersonite
Tremolite
Chrysotile
Sepiolite
Attapulgite
Palygorskite
Zoisite
56
82
33
16
29
5
9
0
0
3
10
0
18
35
8
0
9
8
0
0
0
2
24
0
0
10
0
14
0
15
0
8
5
4
14
2
3
3
16
6
2
30
5
8
0
0.1-0.2 mm Berak, 1963
0.1-0.2 mm
0.1-0.2 mm
Berak, 1963
Berak, 1963
0.1-0.2 mm Berak, 1963
3-179
-------
TABLE 3-77. (continued)
Rock Sr Kd, mg/1 Condition Reference
Talc 6~~
Pyrophyllite 0
Biotite 5
Muscovite - 5
Chlorite Penninite 2
Chlorite, Dellesite o
Sericite 2
Illite 100
Glauconite 25
Celadon He 35
Serpentine 0
Kaolinite 15
Ha Hoy site 58
Allophane 71
Hisingerite 97
Smectites 104
150
96
138
150
217
163
Nontronite 187
Apophyllite 10
rMelvlite -.--. 5 ' •
Feldspars, Orthoclase 10
Sanidine 23,11
Albite , 10
Oligoclase 11
Andesine 10
Laboradorite 8
Scapolite 11
Leucite 11
Anal cite 25
Pollucite 7
Nepheline 0
Sodalite 3
Cancrinite 14
Zeolites:
Mordenite 27 0.1-0.2 mm Berak, 1963'
Stil bite 166
25
Heulandite 47
Faujasite 34
Harmotone 12
Chabazite ' 210
Natrolite • 8
Scolecite 4
3-180
-------
TABLE 3-78. MICROSTRONTIUM-MACROCALCIUM DISCRIMINATION ON SOME 0.1 to 0.2 mm
ROCKS AND MINERALS (BERAK, 1963)
Strontium
Discrimination
Mineral Coefficient
Halloysite ' 1
Montmorillonite 1
Sepiolite 2.3
Chlorite 1
Oellesite • 9
Socialite >9
Anal cite 2
Cancrinite 12
Leucite 3
Orthoclase 9
Orthoclase 11
Microcline 8
Albite 9
Labradorite 9
Mordenite with quartz 1.6
Heulandite . 48
Heulandite 55
Heulandite 19
Mordenite 1.3
Faujasite 4
Harmotome >9
Brewsterite >24
Phillipsite 8
Chabazite 14 .
. Natrolite 9
Scolecite >9
Thomsonite >34
Thomsonite 10.2
SILICATE ROCKS
Basalt "17
Nepheline tephirite 2
Nephelinite . 10
Nepheline-leucite tephirite 4
01ivine hauyne nephelinite 5
Leucite tephrite 5
Leucite tephrite 4
Leucitite with phillipsite 8
Tephrite with chabazite 5
Basalt mineralogy - augite, 40%
plagioclases, 50%
magnetite, 8%
apatite, 1%
Nephelinite - Ti-augite, 68%
magnetite, 12%
nepheline, 10%
plagioclase, 7%
leucite, 3%
3-181
-------
The data of Table 3-77 show that the secondary silicates, such as the
clay minerals and certain zeolites, give much higher Kd values for strontium
than the primary igneous minerals found in rocks, such as quartz and feldspars
Some of the secondary silicates can also discriminate between geochemically
similar calcium and strontium, as seen in Table 3-78.
Tamura and Struxness (1963) outlined several strontium removal mechanisms
that could apply to soil solutions. Two ion exchange reaction types were
postulated. One type involves simple ion exchange on a relatively pH-stable
substrate such as zeolite. The other type was exemplified by alumina that
is amphoteric as a substrate for ion exchange. The differences were further
elucidated by Tamura (1964). Strontium Kd values as high as 30,000 were mea-
sured in alkaline systems on alumina. The strong influence of soil sesqui-
oxides in strontium removal was emphasized.
Baetsle et al. (1964) have given data (Table 3-79) on Kd values for
strontium with Belgian soils. The strontium exchange work on Mol, Belgium,
soils was with deionized water, mains water (tap water?) and groundwater.
None of the water compositions were given. The pH was given probably as mea-
sured in the eluting solution prior to the equilibrations. The Kd values
89
were obtained by adsorbing a band of Sr on the top of a soil column and
eluting with the solution indicated. Samples of the effluent were counted
and the Kd determined from the C/CQ = 0.5 as indicated above. The results
were part of a study to compare static and dynamic Kd values.
TABLE 3-79. VARIATION OF TRACE STRONTIUM Kd VALUES WITH pH
(BAETSLE ET AL., 1964)
Kd. ml/g
Deionized Mains Water Groundwater
Soil Type Water, pH4 pH 7.7 pH 3 pH 4 pH 5 pH 7
Eolian Sand
Horizon A 16-22 38 °7 2.0-2.8 3.6-5.8 5-7
Horizon B 1.5- 1.6 10 13
Horizon C 0.5- 0.9 3.5 12
Mol White Sand 0.95 1.7 4.5 2.1-2.5 5.4-5.6 4.7-4.9
Mol Lignitic Sand -- ' -- — 2.0-3.1 5.0-5.5 6.8-7.4
3-182
-------
The mains water Kd values are higher due to the higher pH. However, the
higher mains water Kd values at pH 3 were difficult to rationalize. They
should be lower than both the deionized water Kd values and the mains'water Kd
values at pH 7.7.
A Burns, Mississippi, montmorillonite, with the composition shown in
Table 3-80, was used in macroion systems of potassium, calcium and magnesium
to obtain Kd values and calculated Kd values in mixed two and three ion
systems (Baetsle et al., 1964). The selectivity coefficients given were
KSr ' °-096
""mont. + Sr> and " °'98
-------
TABLE 3-80. COMPOSITION OF BURNS, MISSISSIPPI, MONTMORILLONITE.
CAPACITY WAS 1.04 meq/g (BAETSLE ET AL., 1964).
Constituent
Sl°2
A1203
Fe2°3
FeO
Ti02
CaO
MgO
H20"
H20+
Total
Wt 1,
51.29
15.96
2.46
0.14
0.18
2.21
3.95
0.05
0.13
18.00
5.37
99.74
TABLE 3-81. COMPARISON OF EXPERIMENTAL AND COMPUTED STRONTIUM Kd VALUES
FOR BURNS MONTMORILLONITE (MERCER, 1967)
Concentration, meq/ml
<-K . CMg ^ca
0.0855 .
0.0404
0.0143
0.00702 .
0.00240
«
—
—
• --
--
0.0885
0.0411
0.0155
0.00612
0.00582
—
—
—
—
..
0.01382
0.00962
0.00604
0.00412
0.00282
0.0180
0.0146
0.0106
0.0068
0.0043
--
«
—
—
—
0.0852
010388
0.0100
0.0046
0.0015
_.
—
«
--
~
0.0769
0.0348
0.0096
0.0032
0.0009
0.0124
0.00850
0.00520
0.00360
0.00248
0.0128
0.0114
0.0074
0.0056
0.0036
Experimental
Kd, ml/g
24.8
59.5
135.5
207.0
369.7
• 12.17
20.6
56.9
114.3
193.5
18.0
50.0
116.0
193.0
336.0
12.3
• 25.8
58.6
100.6
244.1
Baetsle
Kd, ml/g
21
41
102
177
368
12.1
21.3
53.2
106.8
221.8
22.0
48.0
m.o
203.0
300.0
14.5
24.5
55.0
107.0
209.0
Mercer
Kd, ml/g
10
33
102'
191
352
11
21
53
109
212
10
33
102
206
285
11
22
62
104
206
3-184
-------
Hawkins and Short (1965) completed an experimental design to allow mul-
tiple regression analysis of the effects on strontium and cesium Kd values of
+2 ' •+•
the following ions and their concentration ranges: Ca , 5-500 ppm; Na ,
6-650 ppm; K+, 1-100 ppm; NHt, 1-100 ppm; Mg"1"2, 1-125 ppm; Cs+, 0.05-5 ppm;
-3
PO, , 0.1-10 ppm; and pH 4-9. The equation for strontium adsorption may be
applicable to other soils besides INEL, Idaho soils, if they do not contain
abundant sesquioxides.
Wahlberg and Dewar (1965) studied ion exchange adsorption of trace con-
centration strontium (10 N Sr) as affected by two competing cations. The
Kd values calculated with the mass-action equation agreed most closely with
high exchange montmorillonite than with lower capacity kaolinite.
Dlouhy (1967) investigated strontium and other cations adsorbed by
Cascaccia soil and tuff as a function of pH, time and solution composition.
Kd values were used in chromatographic equations to show the radionuclide
distribution beneath disposal sites. Strontium Kd values ranged from 6 to
15 ml/g in the soil to 45 to 75 ml/g in the tuff.
Hajek and Ames (1968) compared batch and small column Kd values for
strontium and cesium on a Burbank loamy fine sand containing up to 2 wt%
calcite. The soil to solution ratio was varied between 100 to 5 for the
batch Kd determinations. Strontium Kd values tended to decline as the solu-
tion to soil ratio declined. The small columns did not yield a real Kd value
for strontium unless it was preleached with at least 450 column volumes of
3N NaNOg. It is probable that calcium from the calcite was solubilizing and
competing with the strontium during the batch tests. Pre-equilibration of
the soil with the solution to be used minus the strontium radioisotope is a
must before Kd values are measured. Gardner and Skulberg (1964) had similar
Kd changes with solid-solution ratio changes.
90
Jenne and Wahlberg (1968) found that the adsorption of Sr was controlled
by in situ precipitation of calcium carbonate in a small stilling basin sample
in White Oak Creek, Tennessee. The clay minerals controlled only the adsorp-
tion of cesium while manganese and iron oxides controlled Co removal.
Cerrai et al. (1969) reported that the strontium Kd on a marine sediment
from a core at 50 m off La Spezia, Italy, was 6.3 ml/g.
3-185
-------
R. J. Serne (PNL), 1973, determined strontium distribution coefficients
from high salt, high pH, synthetic tank solutions with Hanford core well sedi-
ments. The chemical composition of the synthetic tank solutions are given in
Table 3-9 and the sediment characteristics in Table 3-10 under antimony. The
strontium Kd values in Table 3-82 are the average of three values. The rela-
tively low strontium Kd values obtained were principally due to the high con-
centrations of competing cations present.
TABLE 3-82. STRONTIUM Kd VALUES (ml/g)
Solution
Sediment
1
2
3
4
5
6
7
I
3.91
2.90
2.93
12.40
1.87
2.50
1.86
II
1.74
1.17
1.22
5.44
0.81
0.95
0.87
III
0.023
0.047
0.072
0.074
0.13
0.10
0.08
IV
0.0
0.0
0.0
0.0
0.17
0.14
0.23
V
18.40
8.40
12.87
173.84
12.70
39.20
20.16
Routson (1973) determined strontium adsorption characteristics of three
Hanford project soils. The soils were well-character!zed as shown in
Table 3-83. The clay fractions (0.2 to 2 ym and <0.2 urn) were analyzed by
X-ray diffraction and found to be very similar for the three soils. The
Ritzville silt loam, Burbank loamy sand and Ephrata sandy loam showed mont-
morillonite as the dominant mineral in the 0.2 to 2 ym fraction, with lesser
amounts of vermiculite and mica. Montmorillonite was the only clay mineral
detected in any of the <0.2 urn fractions. There was no evidence of clay
mineral weathering after deposition of the parent materials. The only indica-
tion of soil weathering is the accumulation of calcium carbonate in the upper
C horizons of all three soils. The strontium Kd values are given in Table 3-84.
The original soil samples had chiefly calcium in the exchange positions. Note
that the sodium based soils showed higher strontium Kd values, indicating that
calcium displaced strontium when present in the system. It can also be seen
in Table 3-84 that calcite removal from the soils was the most effective way
to raise the strontium Kd values.
3-186
-------
TABLE 3-83.
CHARACTERISTICS OF THE SOILS USED IN STRONTIUM Kd
DETERMINATIONS (ROUTSON, 1973)
Particle Size
% of
Sail Fraction
Ritzville
All
0-2 in
A12 2-18 in.
81
Cca
C
Burbank
All
18-27
27-41
41-72
0-3 in
in.
in.
in.
.
A]2 3-16 in.
AC
AC2
1C
Ephrata
All
A12
Bl
IB2
IIC
16-22
22-30
30 in.
0-3 in
in.
in.'
.
3-12 in.
12-15 in.
15-19 in.
19-24 in.
TABLE 3-84.
Sand
23.
47.
54.
39.
55.
76.
91.
88.
82.
79.
44.
59.
57.
62.
92.
3
0
7
8
0
8
2
7
5
8
2
0
4
9
1
TRACE
FROM 0
Silt Clay
64.1 12.6
46.0 7.0
42.5 2.5
57.7 3.3
42.0 3.0
16.8 6.4
7.6 1.2
10.6 0.7
16.4 1.1
16.2 4.0
47.1 8.7
36.9 4.1
39.3 4.3
31.4 5.7
5.4 2.5
CaC03
% of Total %
-
• o
1
8
4
-
0
1
5
3
-
0
0
0
3
-
.8
.3
.9
.2
-
.7
.6
.5
.5
-
.57
.70 .
.87
.28
Fe2°3 Organic Carbon
of Total % of Total
1
1
0
l'
1
0
0
0
0
0
1
1
1
0
0
.50
.44
.89
.09
.11
.80
.84
.62
.52
.39
.22
.25
.20
.98
.46
—
0.34
0.27
0.13
0.17
--
0.16
0.14
0.13
0.20
—
0.18
0.13
0.29
0.26
STRONTIUM ADSORPTION CHARACTERISTICS
.2N NaCl (ROUTSON, 1973)
Equilibrium Distribution Coefficients (ml/a)
No
Pretreatment
Ritzville
A12
Cca
C
Burbank
A12
AC
AC,
1C
Ephrata
A12
81
IB2
IIC
21
27
27
24
12
15
21
17
12
12
17
23
Na-Based '
109
131
84
58
57
41
53
48
73
68
32
55
Organic Carbon
Removed
37
46
40
43
20
28
35
38
21
22
33
55
Calcite
Removed
189 -
257
71
82
106
91
67
81
115
126
166
145
Organic C
and Calcite
Removed
96
101
61
66
45
61
65
52
44
57
96
70
3-187
-------
Jackson (1976) used tracers to study adsorption of seven radionuclides,
including strontium, on copper ore and decomposition products during copper
leaching, as a function of ore particle size, solution composition, pH and
liquid/solid ratio. The reactions were carried out in autoclaves for up to
8 months in a controlled oxygen pressure and 363 K. The solids continued to
concentrate strontium from the liquid for a long period of time as the ore
and gangue surfaces continued to change and grow other mineral phases. A sub-
stantial amount of activity was adsorbed by the ore and gangue decomposition
products (gypsum, jarosite, anhydrite). As an example, in one long-term
experiment, the strontium Kd was 0.25 ml/g at time zero and 499.6 ml/g at
zero plus 173.6 days.
Migration Results
Field Studies —
Mawson (1956) and Evans (1958) reported the results of studies on fission
product radionuclide movements after disposal in sand at Chalk River, Ontario.
In dilute, neutral pH wastes ruthenium moved much more rapidly through the
sand than did strontium. When acid wastes were disposed into the same sand,
strontium and ruthenium both moved rapidly and at about the same speed through
the sand.
Spitsyn et al.. (1958) used an alkaline solution (4 to 8 g NaOH/1,
200 g NaN03/7) and an acidic solution (6 to 8 g HN03/1, 200 g A1(N03)3/1)
in field studies of radionuclide migration. A mixture of radionuclides were
injected into the sandy soil and the movement traced by wells along the migra-
tion route. Strontium traveled the farthest, with ruthenium traveling about
half as far as strontium. The authors concluded that the groundwater into
which the strontium is disposed should have a low calcium and general dissolved
solids content.
Parsons (1961) described the movement of fission products from a surface
disposal site in Chalk River, Ontario, through the near-surface groundwater,
resulting from a 1954 experimental disposal of 5700 liters of waste containing
2100 kg of ammonium nitrate. The high concentration of ammonium ions caused
all of the strontium to move off relatively rapidly with the groundwater.
Horizontal movement of the waste was rapid enough to reduce vertical dispersion,
3-188
-------
and the strontium remained close to the top of the water table and traveling
at 0.027 times the velocity of the groundwater.
Parsons (1963) described a 1954 disposal of 3800 liters of acid waste
90
containing complexing agents and over 1000 Ci of Sr. The waste was poured
into a pit lined with limestone. Sampling of the sandy soil and near-surface
groundwater allowed the movement of fission products to be followed. The
ruthenium migrated rapidly, traveling at nearly the same velocity as the
90
groundwater. The Sr became separated from the ruthenium, developing into
90
a continuous tongue 198 m long, 3 to 4 m thick and containing 800 Ci of Sr.
90
The ruthenium entered a shallow lake in 1957, while the Sr remained more
deeply buried and traveling much more slowly.
90
The migration of Sr as seepage from Oak Ridge burial pits was investi-
gated by Lomenick et al. (1967). They reported that most of the strontium on
the side walls and bottom of the pit in Conasauga shale was tied up in the
precipitated calcite (calcium carbonate). Water moved 0.15 m/day while the
90
Sr moved 0.24 m/yr.
Mortensen and Marcdsiu (1963) incubated a silty clay loam with 40 uCi
90
Sr for 60 days and obtained hot water extractions or hydrolyzed with 6N HC1.
90
The'supernatant liquid was electrodialyzed with Sr migrating to the anode
90
compartment suggesting negatively charged Sr-organic matter complexes. Gel
90
filtration showed that Sr was complexed or held on exchange sites by high
molecular weight polymers. The radioactivity correlated with the presence of
polyuronides. Phenols, amino acids and keto acids were separated from the
6N HC1 hydrolysate by paper electrophoresis, with no clear correlation between
radioactivity and these components.
Brown (1967) reviewed the hydrology and geology of two contaminated waste
disposal trenches at Hanford. .Several wells were drilled in and around the
90
disposal sites and the Sr and moisture determined in the sediment column
QQ
beneath the disposal sites. The bulk of the Sr activity was contained in a
15 meter vertical section immediately below the bottom of the disposal sites.
-1 90
The average strontium concentration was 10 uCi Sr/g sediment in the
15 meter high strontium concentration zone and fell off rapidly below this
-4 90 90
zone to 10 uCi Sr/g sediment. Spatially anomalous Sr distributions
3-189
-------
were correlated with changes in lithology. Contaminated sediment columns con-
go
taining 0.15 yCi Sr/g were used to determine the effects of groundwater
90
leaching on Sr movement. Fifty column volumes of groundwater removed 4%
90
of the Sr from the sediment column and 500 column volumes removed 31% of
90
the Sr adsorbed on the sediments.
Magno et al. (1970) investigated the migration of strontium through the
effluent lagooning system of the Nuclear Fuel Services plant in western New
York State. They estimated on the basis of analyses that approximately 90%
90
of the Sr passed through the lagoon system and into nearby surface streams.
90
Further, only 12% of the Sr discharged from the lagoon system was associated
90
with suspended solids. Thus 78% of the Sr discharged from the lagoon system
90
was in solution. The low adsorption of Sr was apparently due to a combina-
tion of pH and competing cations lowering ion exchange loading of solids.
For example, when the plant effluent pH was 11 during one sampling period,
90
75% of the Sr was removed in the lagoon system. At another sampling period,
90
plant effluent was pH 7 and only 10% of the Sr was removed by-the lagoon
system.
Merritt (1967, 1976) reported on the leaching and migration of radio-
activity from nepheline syenite-waste glasses buried below the water table at
90
Chalk River, Ontario. Sampling in 1971 showed that the Sr front in sandy
soil had reached 33 m. Since the leach rate from the glass blocks has been
constant for several years, the plume size has reached a steady state.
Himes and Shufeldt (1970) investigated the effects of soil organics on
90
Sr migration through soils. Thirty 0.002-acre soil microplots were deline-
QO
ated and 21.64 yCi 3USr added to the surface of each. Facilities were in.
90
place to collect runoff water and leachates at 40 cm depth, and Sr was
on
found in the leachate. The decreasing order of effectiveness of Sr removal
90
by organic compounds on the quantity of Sr adsorbed by most soils was phytic
acid > glucuronic acid > citric acid > pectic acid > pyrocatechol > glycine >
90
dextrose, as shown in equilibrium experiments. Adsorption of Sr by soils
high in organic matter treated with HF-HC1 solution to remove clay minerals
90
showed that methylation with diazomethane decreased the total quantity of Sr
adsorbed. Saponification,. acetylation and remethylation also altered the
3-190
-------
90 90
amount of Sr retained on the soil. The Sr contained in plant residues
leached through 30 cm soil columns at a slow but measurable rate with applica-
90
tions of distilled water. Drying of soils containing Sr at 105°C before
90
extraction increased the fraction of water soluble Sr that could be extracted,
Work on the long-term migration of fission products from the natural
reactor site at Oklo (Brookins, 1976) indicated that strontium had partially
migrated from the reactor vicinity in the past 1.8 billion years.
Laboratory Studies ~
Nishita et al. (1956) studied the general extraction of radionuclides by
water leaching, by exchange with IN ammonium acetate and remaining in the soil
90
or clay mineral. In both the soils and clays (bentonite, kaolinite) Sr was
prominent in the water soluble and cation exchangeable fractions. It showed
the least tendency to remain in the soil or clay.
Orcutt et al. (1956) examined hydraulic and ion exchange phenomena as
related to the underground movement of radiostrontium. Ion exchange reduced
the linear velocity of the radionuclide front in relation to the liquid front.
If the solid media are selective for a radionuclide, the radionuclide concen-
tration front becomes sharp. The front becomes wider with distance traveled
if the radionuclide selectivity is unfavorable. The latter condition also
will occur if the waste velocity is too fast to allow instantaneous equilibrium
between exchanger and the radionuclide at the concentration front.
Another article, by Orcutt et al. (1957) includes experimental work on ion
exchange in columns and its theoretical treatment. Several theories are pre-
sented that quantitatively describe hydraulic dispersion, ion exchange equili-
bria and ion exchange kinetics as they apply to the movement of radionuclides
through natural porous media. In the opinion of the authors, the theoretical
models, used in conjunction with field tests, served to reduce the uncertain-
ties of ground disposal operations.
Spitsyn et al. (1960) emphasized again the ease of migration of strontium
in the soil. A well was drilled and the migration of fission products was
+2
followed in the soil at a field site. At the site, up to 50 ml/1 of Ca in
the groundwater caused the initial migration rate of strontium to be about
6 cm/day, practically the same speed as the grcundwater. Drilling 15 months
3-191
-------
later showed that strontium contamination had spread to 6 m distance from the
injection point, and that the radiostrontium migration rate was 1.1 to 1.3 cm/
day compared to 6 cm/day water velocity.
A mathematical model of strontium movement in soils based on the labora-
tory work of Miller and Reitemeier (1957) was proposed by Thornthwaite et al.
(1960). Miller and Reitemeier showed strontium distribution curves for
Norfolk fine sandy loam, Miami silt loam, Huntley clay loam and Fort Collins
silt loam after the application in cycles of up to three solutions (deionized
water, 0.005N CaCU, 0.005N NaCl) at rates of 30 and 300 in. The strontium
in the sandy soil moved farther than in the other three soils.
To describe these distributions, Thornthwaite et al. (1960) set up a
model that relates leaching efficiency (cycles/in.) to soil cation exchange
capacity. The model was used on a New York soil to compute the distributions
90
of Sr from fallout and compare the results to measured values. The compari-
son showed good agreement.
A study of the influence of climatic and hydrologic factors on the move-
ment of strontium through the upper layers of soil was completed by Mather
and Nakamura (1961). Strontium fallout moved through the upper part of the
soil profile at a rate determined by the volume of leaching solution and the
chemical and cation exchange properties of the soil and soil solution. A
mathematical model was devised that allows computation of strontium movement
in the.soil. To use the model, values are needed for the initial strontium
concentration, leaching solution volume and leaching efficiencies (cycles/in.).
The leaching efficiency varied from high values in most regions to low values
in dry regions for soils of the same texture. Leaching efficiency is related
to moisture index. Tables and monograms were prepared to permit determination
of leaching efficiency from readily available climatic data.
Jacobs (1963) reviewed an earlier paper (Jacobs, 1960) and a paper by
Bovard and Grauby (1960) on retention of radiostrontium by soils. Bovard'and
90 90
Grauby made laboratory and field studies of the downward migration of Sr- Y
in undisturbed soils. The recovered solution was analyzed and the exposed
block of soil autoradiographed after the experiment. The experimental results
represented a qualitative description of the nature of the soil and system
3-192
-------
hydrodynamics. The use of a water tracer was suggested by Jacobs to obtain
groundwater velocities. If heterogeneities are not severe, and flow conditions
are fairly constant, the groundwater velocities and a time transformation term
for radionuclide adsorption would provide a quantitative site evaluation.
Miller and Reitemeier (1963) studied leaching of strontium with applica-
tions of 76.2 cm and 762 cm of distilled water, 0.005N NaCl and 0.005N CaCK
solutions to columns containing soils contaminated with strontium-90. The
CaClp solution produced the greatest movement and the deionized water the
least. The maximum distance the strontium moved with 30 in. of water was
1.3 in., and with 300 in. of water the distance moved was 4.3 in. Movement
was greatest in the Norfolk loamy sand.
Some associated K values for the Mol soils are listed in Table 3-85
i
(Baetsle et al., 1964), where K is a measure of the velocity of strontium
through the soil relative to groundwater velocity. Mol sand was almost pure
silica.
TABLE 3.85. K VALUES OF STRONTIUM IN THE SATURATED SUBSOILS
AT MOL, BELGIUM (BAETSLE ET AL., 1964). K IS
DEFINED ON PAGE 2-30.
Soil Type pH 7 pH 3
Mol Sand 5 1
Perturbed Profile 48 1
Mol Sand & Eolian Sand 525
Mol Sand & Lignite 570
Mol Sand 6 1
Mol Sand 71
Mol Sand 3 1
Mol Sand 4 1
Mol Sand 3 1
Mol Sand 6 1
The strontium Kd values and associated relative strontium velocities
through the Houthulst clays at pH 3 are given in Table 3-86. Houthulst is
3-193
-------
located in western Belgium about 40 km from the North Sea. The strontium Kd
values are somewhat higher for the clay but still much less than cesium.
Cesium K averages 1000 while the strontium K averages 100.
TABLE 3-86. STRONTIUM DISTRIBUTION COEFFICIENTS AND RELATIVE MIGRATION
RATES IN HOUTHULST CLAYS AT pH 3 (BAETSLE ET AL., 1964)
Depth, m
0-1
1-2
2-2.5
2.5-3
3-4
4-5
5-6
6-7
7-8
8-9
9-10
0-1
1-2
2-3
3-4
4-5
5-6
6-7
7-8
8-9'
9-10
Core Well
I
Kd, ml/q K Depth, m
7.2
10.6
14.1
9.4
—
11.8
13.9
15.8
19.4
17.3
20.2
__
3
5
18
18
23
29
31
16
29
24
36
48
32
—
40
47
54
66
59
68
Core Well
._
11
18
61
61
78
98
104
54
98
10-11
11-12
12-13
13-14
14-15
15-16
16-17
17-18
18-19
19-20
II
10-11
11-12
12-13
13-14
14-15
15-16
16-17
17-18
18-19
19-20
Kd, ml/g
21.8
27.5
30.7
33.5
33.5
31.8
31.8
34.0
32.6
31.5
30
30
28
29
31
31
29
31
31
31
K
74
93
102
112
112
107
107
114
no
106
101
101
95
98
104
104
98
104
104
104
Polyakov et al. (1969) determined the diffusion coefficients of strontium
-7 -7 2
in three soils. The coefficients varied from 0.6 x 10 to 20 x 10 cm /sec.
These authors pointed out that diffusion coefficients vary with the soil type
3-194
-------
and probably change appreciably within the same soil. Prokhorov (1962) had
shown earlier that the diffusion coefficient in soil varied with the soil/water
content. In the majority of the cases, however, except for the very long-lived
radionuclides and very slow flow velocities, diffusion does not contribute
greatly to radionucl.ide migration.
According to Schulz (1965), the radionuclides showing a degree of mobility
include those of strontium. The amounts and kinds of competing cations have a
marked effect on soil profile distribution as shown by the effects of compara-
gn
ble concentrations of Nad and CaCl2 on redistribution of Sr down the soil
profile. Cad* moves the strontium more readily than Nad, and the more con-
centrated the CaCl2» the more rapid and further the strontium movement.
Nishita and Essington (1967) continued work on leaching strontium from
soil columns. All soils showed that movement by water leaching resulted in
more rapid movement than strontium only by ruthenium.
Carlile and Hajek (1967) reported an example of physical transport of
85
strontium. About 0.1% of the Sr was found in the first effluent from N-Area
soil columns (2 cm diameter by 40 cm in length). Up to 1% random leakage of
strontium was reported until 5 column volumes-of Columbia River water had
passed through-the soil'. Then leakage fell to below detection limits (10 %).
Columbia River water has an annual mean pH of about 7.8, a mean CaC03 hardness
of 69 mg/1 and a mean total dissolved solids of 119 mg/1. Effluent samples
were centrifuged at high speeds, treated with HLO^ to destroy organic material
and recentrifuged. Fifty to 75% of the activity in the leaked effluent was
removed. Hydrogen peroxide treatment did not affect removal of strontium by
centrifugation, indicating that an organic material was not involved. Leach-
ing a soil system with low ionic strength river water containing trace radio-
nuclides had the result of dispersing the soil colloids on which much of the
85
Sr were adsorbed. Column breakthrough was immediate until the dispersed
colloids were flushed from the column. Preceding the contaminated river water
by a preleach of noncontaminated river water flushed the dispersed colloids
and prepared the soil for use. Adding other ions to the system could have
accomplished the same thing, but would have competed with adsorption of
strontium.
3-195
-------
90
Brendakov et al. (1969) followed the average displacement of Sr from a
90
Chestnut soil. The Sr was amended to the soil in the field and sampled at
yearly intervals. Two years after amendment, the average displacement dis-
tance from the surface was 11 mm and 5 years after amendment, 16 mm.
The average displacement was defined as:
. x =
where
pi = fractional % of the radionuclide in a layer (3 mm thickness)
xi = the depth of the layer.
There was little difference in X" for any of the radionuclides used in the
experiment.
Chelation and movement of strontium in a calcareous soil were evaluated
in columns of Mohave sandy loam soil (Fuller and L'Annunziata, 1969). The
89
displacement of Sr from the soil columns was 19, 38, 44, 47, and 81% in the
order HEEDTA < DCyTA < EDTA < EGTA < DTPA. Only 5% of the 89Sr was removed
from the soil with no complexing agents in the leaching solution. The radio-
strontium existed in the leachate as a negative complex. Any natural organic
chelating agents in the soil did not influence radiostrontium movement.
Four soils varying greatly in organic matter content, cation exchange
capacity and texture were studied by Juo and Barber (1970) for their stron-
tium retention characteristics. The adsorption of strontium by soils increased
with increasing pH within the pH range 4 to 8. Saturating cations had some
effect on strontium adsorption, with influence in the order Na>K>Mg>Ca>Ba>H.
As the pH increased, an increasing fraction of the adsorbed strontium became
nonexchangeable with ammonium acetate. The nonexchangeable portion of the
strontium is probably tied up by soil organics and is permanently fixed under
favorable pH conditions. In a higher pH environment, a larger portion of the
strontium in solution was present as soluble strontium chelates or complexes.
Varga and Jacobs (1970) examined how the strontium from dissolved (Ca,
85
Sr) CO, moved through a mineral exchange system as represented by a ver-
85
miculite column. The (Ca, Sr) CO., was precipitated on the top of the column
flC QC
by mixing 50 ml of 0.4M NaOH, 50 ml of 0.7M NaCl, *°Ca, °°Sr, 1 ml 1M NaHC03
3-196
-------
and 1 ml 1M CaCNO^Jot and allowing the mixture to stand 1 hr. ' An HC1 solution
was then used to leach the strontium-contaminated calcite precipitated on top
of the column. The behavior of the released calcium (traced with Ca) and
strontium were very similar. Less than 1% of the Ca and Sr on top of the
column appeared in the effluent prior to acid addition, and no significant
activity was found in the effluent until breakthrough of the acid front
occurred. The release of calcium and strontium from the system was directly
related to the amount of acid added in the leaching process.
Summary
Most of the strontium chemical compounds for which data are available are
very soluble (Figure 3-22). Only in highly alkaline soils could SrCCL control
strontium activity in solutions. Strontium in solution is expected to be
2+
predominantly present as Sr (Figure 3-23) and to be exchangeable on soils.
Laboratory studies show that ion exchange is the principal mechanism of
strontium adsorption by soils and rocks (McHenry, 1955, 1958; Rhodes and
Nelson, 1957; Klechovsky et a!., 1958; Prout, 1958, 1959; Baker and Beetem,
1961; Cohen and Gailledreau, 1961). Strontium Kd values are usually directly
correlated with cation exchange capacity of the soil or rock, but not
Invariably (McHenry, 1958). A direct correlation between solution pH'and
strontium Kd has also been reported (Rhodes, 1957; Prout, 1958, 1959;
Baetsle et al., 1964; Juo and Barber, 1970), which suggests hydrogen ion com-
2+
petition with Sr for exchange sites. Strontium Kd values fall rapidly from
100 to 200 ml/g or more in low ionic strength salt solutions to less than
5 ml/g in high ionic strength salt solutions (Rhodes, 1957; Seme, 1973).
Strontium and calcium most readily replace trace strontium (Kokotov et al.,
1961, 1962; Schulz, 1965) and cause strontium radioisotopes to migrate rapidly
when present in groundwater with the strontium (Spitsyn et al., 1960). Low
pH also results in a relatively high strontium migration velocity in ground-
water (Parsons, 1963; Baetsle et al., 1964) with the strontium and ground-
water velocity identical at pH 3. Calcium competition is the probable cause
of the 90% soluble strontium migration from the lagoons of the Nuclear Fuel
Services plant (Magno et al., 1970), although strontium also migrates on soil
particulates (Carlile and Hajek, 1967). The many strontium Kd values deter-
mined by Berak (1963) demonstrated that secondary minerals such as clays and
3-197
-------
zeolites are much better strontium adsorbers and are more selective for stron-
tium from solutions containing strontium and calcium than are the primary
minerals such as quartz, feldspars and pyroxenes. Another mechanism for
strontium removal from solution is its coprecipitation with calcite (Jenne
and Wahlberg, 1968; Varga and Jacobs, 1970) or adsorption during phosphate or
oxalate replacement of carbonate in calcite (McHenry, 1958; Ames, 1958).
Strontium organic complexes are known that result in a measure of strontium
fixation (Mortensen and Marcusiu, 1963) or enhanced migration rates (Himes
and Shufeldt, 1970).
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Parsons, P. J. 1963. Migration from a Disposal of Radioactive Liquid in
Sands. Health Physics. 9:333-342.
Polyakov, Yu. A., V. F. Gol'tsov, and V. G. Grakovskiy. 1969. The Diffusion
of Strontium-90 in Soils. AEC-tr-7030, pp. 14-25.
Prokhorov, V. M. 1962. Diffusion of Strontium-90 in Soil and Sand.
Radiokhimiya. 4:205-211.
Prout, W. E. 1958. Adsorption of Radioactive Wastes by Savannah River Plant
Soil. Soil Science. 86:13-17.
•Prout, W. E. 1959. Adsorption of Fission Products by Savannah River Plant
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Rhodes, D. W. 1957. The Effect of pH on the Uptake of Radioactive Isotopes
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Rhodes, D. W. and J. L. Nelson. 1957. Disposal of Radioactive Liquid Wastes
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Routson, R. C. 1973. A Review of Studies on Soil-Waste Relationships on the
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Schulz, R. K. 1965. Soil Chemistry of Radionuclides. Health Physics.
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.\
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3-203
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TECHNETIUM
Natural Soil and Rock Distributions
According to Boyd and Larson (1956), technetium occurs in most abundance
in molybdenite (MoS2), a common molybdenum ore, as shown in Table 3-87. The
technetium content of all materials except pitchblende is of secondary origin,
formed by the action of cosmic radiation on molybdenum, ruthenium or niobium.
99 238
The Tc in pitchblende is a product of I) spontaneous fissioning. Not
99
everyone accepts the existence of the natural Tc as having been proven (see
Handbook of Chemistry and Physics, 57th Edition, p. B-50).
TABLE 3-87. TECHNETIUM CONTENT OF SOME NATURAL MATERIALS
(BOYD AND LARSON, 1956)
Material Tc, g/kg Determination Method
Molybdenite, Climax, CO <4 x 10~1 Spectroscopic
Molybdenite, NV <1.5 x 10 7 ,, Spectroscopic
Molydenite, Miami, AZ <1.8 x 10" > 8.3 x 10" , Activation
1.3 x NT 1
Molybdenite, NV 1.3 x 10_g Activation
Osmium-iridium concentrate <1.6 x 10~?, Spectroscopic
Yttrotantalite, West Africa <6 x lO'jl Mass Spec.
Iron-nickel Meteorite <4.5 x 10~g Spectroscopic
KReOa, pure . ' ' <8 x 10~,n QQ Spectroscopic
Pitchblende (--50% uranium) 2.5 x 10"IU *yTc (Ashizawa and Kuroda, 1957)
Brief Chemistry.
There are 16 isotopes and 6 isomers of technetium presently known
(Lavrukhina and Pozdnyakov, 1970). Several technetium isotopes are obtained
by the fissioning of nuclear fuels of uranium in the reactor. The fission
99 101
product technetium isotopes and their yields include: Tc, 6.06%; Tc,
5.6%; 102Tc, 4.'3%; 103Tc, 3.0%; 104Tc, 1.8%; 105Tc, 0.9%; 107Tc, 0.19%
(Katcoff, 1958). According to Boyd (1959), the technetium content of first
QQ C
•cycle Purex waste is 41 mg/1. Only Tc is long half-lived (2.12 x 10 years)
of the several fission product technitium isotopes, and hence a potential
hazard.
Technetium exists in valence states from (VII) to (-1). The most stable
state in aqueous solutions is heptavalent, usually corresponding to the TcOT
(pertechnetate) ion, which other valences encountered chiefly in complex
3-204
-------
+7 °
compounds. The ionic radius of Tc is 0.56 A, while its geochemical rela-
+7 °
tive, Re , is 0.46 A. The corresponding radius of the solvated TcOd ion is
o ^
4.6 A (Smith et al., 1953). In weakly acid, neutral and alkaline solutions,
Tc(IV), the second most stable state of technetium, is oxidized by atmospheric
oxygen to Tc(VII). The solubility of KTc04 in water at 20°C is 21.3 g/1
(Busey and Larson, 1958). Technetium (IV) and Tc(V) form halogen complexes
2
such as (TcClg) , and Tc(II), (III) and (V) form diarsine complexes and
cyanide complexes.
Solid Phase and Solution Equilibria
Baes and Mesmer (1976) reported that technetium (VII) forms strong peracids
(HMOA), and that its oxides are very soluble. No information is available con-
cerning the soluble complexes of technetium or the solubility of its salts
(Pourbaix, 1966). Pertechnetate ion, TcOT, is the most stable species of the
element in aqueous solutions (Boyd, 1959).
Experimental Adsorption Results
Wildung et al. (1975) used a batch equilibrium technique to determine
pertechnetate ion adsorption on soils. The range of the 22 soil type proper-
ties used in the study are given in Table 3-88. The Kd values for pertechneta-te
ranged from 0.007 to 2.8. The Kd values were positively correlated with soil
organic carbon content and negatively correlated with soil pH. The positive
charge on soil organic colloids is probably an important factor in technetium
adsorption by soils. The soil carbon-pertechnetate ion removal correlation
was the only one significant at the 1% confidence level. Cation exchange
capacity-pertechnetate ion removal also was positively correlated but at a 5%
confidence level. Negatively correlated at the 5% confidence level was pH-
pertechnetate ion removal, probably because as pH decreases, positive charge
increases with decreased ionization of acidic groups of soil organic colloids
and increased protonation of basic groups. A regression equation was given
for pertechnetate distribution coefficient determination within the range of
significant independent variables. This was Y-j = 0.08X3 - 0.09X4, where Y-j is
the pertechnetate distribution coefficient (based on the ranges shown in
Table 3-88), X3 is the soil organic carbon content and X, is the soil-solution
pH. The soil silt content (X,) and clay content (Xg) correlations with Kd
3-205
-------
were not statistically significant at the 5% confidence level and did not,
therefore, appear in the regression equation.
TABLE 3-88. RANGE OF PROPERTIES OF 22 SURFACE SOILS USED IN THE PERTECHNETATE
ADSORPTION STUDIES (WILDUNG ET AL., 1975)
Cation Exchange wt%
Capacity, meg/100 g pH C03 Organic C Sand Silt Clay
5.5-90.0 3.6-8.9 0-6.5 0.23-28.8 14.1-73.1 17.6-58.0 3.8-46.6
Gast (1975) studied several Minnesota soils and their abilities to adsorb
pertechnetate ions. The soil characteristics are outlined in Table 3-89.
From 2 to 5 weeks of contact time between soil and solution were required for
equilibrium to be established. Nicollet C and Zimmerman soils, both low in
99
organic carbon content, adsorbed very little Tc, while up to 98% of the
technetium was adsorbed from solution by the remaining soils corresponding to
99
a Tc Kd value of 857.5 ml/g. When the organic carbon was removed by hydrogen
peroxide treatment from two soils originally high in carbon content, the adsorp-
tion of technetium was greatly reduced. Gast also mentioned that anaerobic
conditions .prevailed in the soil-solution environment which may have led to
the precipitation of technetium heptasulfide, rather than technetium,removal
by adsorption processes.
TABLE 3-89. PROPERTIES OF SOILS USED IN THE TECHNETIUM
ADSORPTION STUDIES (GAST, 1975)
Soil
Bearden
Hegne
Ribbing
Nicollet A
Nicoll«t C
Omega
Bergland
Arveson
Uaukegan
Zimmerman
Peat
Sand
9.1
1.9
3.3
24.4
23.9
61.5
14.0
47.1
10.0
69.4
—
wt%
Silt
63.7
37.6
79.5
42.9
48.6
30.7
25.1
24.6
69.2
24.9
~
Clay
27.2
60.5
17.2
29.7
27.5
7.9
60.9
28.3
20.7
5.7
--
pH,
7
8
5
5
8
7
6
7
6
5
7
H?0
.68
.00
.48
.91
.40
.87
.35
.70
.25
.74
.83
Organic
C, wt%
5.39
2.20
2.29
2.39
0.12
1.26
5.67
2.80
2.36
0.80
45.95
wt% % Fe203
16
12
15
1
15
7
.29
.79
—
--
.21
.21
—
.58
—
—
.58
0
0
1
0
1
1
2
0
1
0
0
.09
.29
.40
.86
.35
.20
.39
.21
.03
.47
.40
C.E.C.,
meg/ 100 g
16.9
36,1
11.3
19.3
15.7
6.0
32.3
14.9
15.1
2.7
. 50.9
3-206
-------
Routson et al. (1976, 1977) determined distribution coefficients on tech-
netium with Washington and South Carolina soils. Soil properties are given in
Table 3-90. The organic carbon content was not given. The technetium Kd
values obtained were very low, as seen in Table 3-91, when competing bicarbonate
ions were present in the system. All Kd values were essentially zero.
TABLE 3-90. PROPERTIES OF SOIL SAMPLES
(ROUTSON ET AL., 1977)
wt%
Soil
Washington
South Carolina
CaCO-3, mq/q
0.8
<0.2
Silt
10.1
3.6
Clay
0.5
37.2
CEC, meq/100 q
4.9
2.5
pH, H?0
7.0
5.1
TABLE 3-91.
TECHNETIUM Kd AS A FUNCTION OF NaHCOs CONCENTRATION
FOR A SOUTH CAROLINA SOIL CHARACTERIZED IN TABLE 3-90
(ROUTSON ET AL., 1977)
NaHCOs, M
0.002
0.008
0.020
0.200
Average Kd, ml/q
0.019 ± 0.06
0.0 ± 0.01
' 0.0 ± 0.01
0.010 ± 0.04
Migration Results
Field Studies-
Brown (1967).traced the movement of technetium discharged to the ground
near the chemical separations plant at Hanford via an extensive network of
monitoring wells. Based on analytical results, a map of the contamination
pattern at the surface of the groundwater was given for 1966 at limits defined
99 3
by the 0.01 pCi Tc/cm isoconcentration contour. The contamination pattern
was very similar in shape and size to that for tritium. Technetium-99 was
reported to be moving at essentially the same rate as the groundwater, and was
detectable in the groundwater at below public drinking water limits all the
way to the Columbia River.
3-207
-------
Brookins (1976) pointed out that the technetium formed as a result of
nuclear fission of the deposit of uranium at Oklo, Gabon, had at least par-
tially migrated from the shale location in which it was formed.
Laboratory Studies—
Gast (1975) reported the extractability of technetium from Bergland and
Arveson soils. The characteristics of these soils are given in Table 3-92.
Two-gram samples of air-dried Bergland and Arveson were loaded with 0.06 uCi
QQ QQ
of Tc. Less than 2% of the Tc still remained in the equilibrating solu-
QQ
tion after 1 month of contact. These 2-g samples loaded with 0.06 uCi Tc
were used in the extraction studies. Twenty-five milliliters of extractant
(IN NaOH or IN HCIO^) were added to the 2-g samples in sequence and shaken for
24 hr with the results shown in Table 3-93.
TABLE 3-92. PROPERTIES OF BERGLAND AND ARVESON SOILS USED
IN THE 99fc EXTRACTION WORK (GAST, 1975)
wt%
Bergland
Arveson
Sand
14.0-
47.1
Silt
25,1 _
24.6
Clay
60.9
28.3
Organic I
Carbon, %
5.67
2.80
" %
-
15.58
pH,
HoO
6.35
7.70
CEC,
meq/lOOg
32,3 .
14.9
TABLE 3-93. EXTRACTABILITY OF Tc FROM BERGLAND AND
ARVESON SOILS (GAST, 1975)
Soil Extractant Extraction 99Tc. « Total 99Te.i
Bergland NaOH
Arveson NaOH
Bergland HC104
Arveson
1
2
3
4
1
Z
3
4
1
2
3
4
1
2
3
4
77.4*3.25
8.69*1.97
1.82±0.16
0.83+0.10
83. 68* 1.05
4.77t0.07
2.39*0.11
0.98±0.07
11.95±0.14
6.90*0.31
4.46t0.07
3.52±0.01
7.02*0.01
2.96±0.04
2.16±0.04
1.50±0.01
88.74*1.02
91.82±0.98
26.83±0.60
13.63±0.02
3-208
-------
The concomitant removal of soil organic matter suggested that the
ion may be associated with it. However, Gast showed that large amounts of
p qq
chloride ions (KC1) or HP04 from K2HP04 had no significant effect on Tc
adsorption by Bergland and Arveson soils so that simple anion exchange prob-
qq
ably was not involved in the adsorption. The low Tc adsorption shown by
low organic soils, and the effects of prior hydrogen peroxide removal of
99
Bergland organic material on lowering the Tc adsorption capacity, strongly
suggest that the technetium is adsorbed on the soil organics. The presence
of FLS reported from certain of the adsorption experiments suggests the reduc-
tion of Tc(VII) to Tc(IV) and precipitation of TCgSy as well. The NaOH removes
both adsorbed and precipitated technetium from the soil, while the perchloric
acid is partially consumed by reactions with the organic fraction and is
therefore not as an effective technetium removal agent as NaOH.
Summary
Pertechnetate ion, TcOT, is the most stable species of technetium in
aqueous solutions (Boyd, 1959). A major portion of the ion exchange capacity
of soils and sediments is cation exchange capacity at the usual near-neutral
pH conditions. Therefore a negatively charged TcOT would hardly be exchanged,
and hence show littTe adsorption by soils and rocks (Routson et al., 1976,
1977). Gast (1975) reported high Kd values (800 to 900) for soils with high
organic matter. He also reported the presence of H2S in his equilibrating
samples, and suggested that the high Kd values were due to precipitation of
technetium as Tc^Sy in the presence of H^S (Kotegov, 1968). In conditions
where the soil contains appreciable organic matter, the Tc(VII) may be reduced
to Tc(IV) and adsorbed (Wildung et al., 1975). Tc(IV) can coprecipitate with
ferric hydroxide (Anders, 1960). With the exception of some sedimentary
rocks, most rocks contain very little organic matter. Hence, technetium Kd
values of close to zero would normally be encountered in oxidizing conditions
and in rocks and soils that are relatively low in organic matter. More study
is required to determine the nature of the organic matter-pertechnetate ion
reaction, and the influence of redox conditions on adsorption of technetium
by soils and rocks.
3-209
-------
References
Anders, E. 1960. The Radiochenristry of Technetium. Nuclear Sci. Ser. Nat.
Acad. Sci. - Nat. Res. Council. NAS-NS 3021.
Ashizawa, F. T. and P. K. Kuroda. 1957. The Occurrence of the Short-Lived
Iodine Isotopes in Natural and in Depleted Uranium Salts. J. Inorg. Nucl.
Chem. 5:12-22.
Baes, C. F., Jr. and R. F. Mesmer. 1976. The Hydrolysis of Cations. John
Wiley and Sons, New York.
Boyd, G. E. 1959. Technetium and Promethium. J. Chem. Educ. 36:3-14.
Boyd, G. E. and Q. V. Larson. 1956. Report on the Occurrence of Technetium
on the Earth's Crust. J. Phys. Chem. 60:707-715.
Brookins, D. G. 1976. Shale as a Repository for Radioactive Waste: The
Evidence from Oklo. Environmental Geology. 1:255-259.
Brown, D. J. 1967. Migration Characteristics of Radionuclides Through Sedi-
ments Underlying the Hanford Reservation. ISO-SA-32.
Busey, R. H. and Q. V. Larson. 1958, Chemistry Division Annual Progress
Report for Period Ending June 20, 1958. ORNL-2584.
Gast, R. G. 1975. The Behavior of Technetium-99 in Soils and Plants.
Progress Report. COO-2447-1.
Katcoff, S. 1958. Fission-Product Yield from U, Th, and Pu. Nucleonics.
16:78-86.
Kotegov, K. V., 0. N. Pavlov and V. P. Shvedov. 1968. Technetium. IN^;
Advances in Inorganic Chemistry and Radiochemistry, vol.'IT, pp. 1-90,
Academic Press, NY.
Lavrukhina, A. K. and A. A. Pozdnyakov. 1970. Analytical Chemistry of
Technetium, Promethium, Astatine, and Francium. Translated by R. Kondor.
Ann Arbor-Humphrey Science Publishers.
Pourbaix, M. 1966. Atlas of Electrochemical Equilibria in Aqueous Solutions.
Pergamon Press, Oxford, England.
99 237 241
Routson, R. C., G. Jansen and A. V. Robinson. 1976. Tc, Np, and Am
Sorption on Two Subsoils from Differing Weathering Intensity Areas. BNWL-2000.
Pt. 2, pp. 50-52.
241 237 99
Routson, R. C., G. Jansen and A. V. Robinson. 1976. Am, Np, and Tc
Sorption on Two United States Subsoils from Differing Weathering Intensity
Areas. BNWL-1889.
Smith, W. T., J. W. Cobble, and G. E. Boyd. 1953. Thermodynamitf Properties of
Technetium and Rhenium Compounds. I. Vapor Pressures of Technetium Heptoxide,
Pertechnic Acid and Aqueous Solution of Pertechnic Acid. J. Am. Chem. Soc.
75:5773-5776.
Wildung, R. E., R. C. Routson, R. J. Seme, and T. R. Garland. 1975. Pertech-
netate, Iodide and Methyl Iodide Retention by Surface Soils. BNWL-1950.
Pt. 2, pp. 37-40.
3-210
-------
THORIUM
Natural Soil and Rock Distributions
The content of thorium in rocks and soil is given in Table 3-94. The
thorium in sedimentary rocks may be due either to the selective.adsorption of
thorium on clays or its retention in heavy resistate minerals such as monazite.
Concentrations of thorium in metamorphic rocks are highly variable. The
thorium content of igneous rocks increases from basalts to granites.
TABLE 3-94. THORIUM CONTENT OF COMMON ROCKS AND SOILS
Rock Type
Granites
Intermediate
Basalt and Gabbros
Shales, North America
Bauxites
Bentonites
Limestones
Sandstones
Thorium Average
or Range, ppm
Igneous Rocks
10-20
2-10
0.5-2
Sedimentary Rocks
10-15
49
24
1.1
1.7
Reference
Rogers, 1964
Heier and Carter, 1964
Heier and Rogers, 1963
Adams and Weaver, 1958
Adams and Richardson, 1960
Adams and Weaver, 1958
Adams and Weaver, 1958
Murray and Adams, 1958
Soils
Marble
Slate
Phyllite
Schist
Gneiss
Metamorphic Rocks
0.03
7.5
5.5
7.5
13.1
Vinogradov, 1959
Pliler, 1956
Pliler, 1956
Pliler, 1956
Pliler, 1956
Billings, 1962
Thorium adsorption studies were performed by several investigators.
Holland and Kulp (1954) found red clay, globigerina ooze and green clay
readily adsorbed thorium. They concluded that ion exchange was the adsorption
3-211
-------
mechanism. Adams et al. (1959) suggested that thorium is concentrated by clay
minerals. Up to 50 ppm thorium in the aluminum hydroxide and resistate min-
erals in bauxite was reported by Adams and Richardson (1960).
Brief Chemistry
There are 13 isotopes of thorium with six of them found in nature. Of
the six natural thorium isotopes, five are relatively quantitatively unimpor-
tant members of the 238U, 235U or 232Th decay series. Thcrium-232 is the
major isotope, with a half-life of 1.39 x 1010 years (Ryabchikov and Golbraikh,
1969). Thorium radionuclide data of interest in waste disposal operations
are in given in Table 3-95.
v TABLE 3-95. THORIUM RADIONUCLIDE DATA (RYABCHIKOV
AND GOLBRAIKH, 1969)
Isotope
227Th
228
^bTh
229Th
230Jh
231 Th
232Th
234Th
Half -Life
18.5 days
1.913 years
7340 years
80,000 years
25.5 hours
1.41 x 1010 years
24.1 days
Decay Mode
a
a
a
a
6"
.a
e"
Although other oxidation states of thorium are known in the laboratory,
only Th(IV) is found in nature. Th(IV) is found as Th+ . The atomic radius
of Th+4 is 0.99 A (Ahrens, 1952).
Common insoluble thorium compounds include the hydroxide, fluoride and
phosphate. Soluble compounds include the chloride, nitrate and sulfate.
Thorium in solution is a small, highly charged ion that undergoes extensive
interaction with water and many anions. The solution chemistry of thorium is
largely a-study of its complex ions. Common anions that form strong complexes
with thorium include fluoride, chloride, nitrate, phosphate and sulfate. At
pH values above 3, thorium undergoes hydrolysis in aqueous solutions. During
the sedimentary cycle, thorium usually becomes separated from uranium because
.the uranium tends to mobilize in its U(VI) oxidation state until encountering
a reducing environment to become immobilized U(IV). Thorium does not undergo
a comparable oxidation state change.
'3-212
-------
Solid Phase and Solution Equilibria
4+
Figure 3-24 relates the activity of Th to pH under an assumed weathering
environment in equilibrium with various thorium solid phases. The thermo-
dynamic data for Th02(s) were selected from Baes and Mesmer (1976). The data
for the other compounds were selected from Si 11 en and Martell (1964). Under
the assumptions outlined in Figure 3-24, all thorium compounds except ThF4 can
be arranged in an increasing order of stability throughout the pH range as
follows: Th(HP04)2, Th3(P04)4, Th(OH)4, and
in a pH of approximately >7 and most stable approximately pH <4.
-4
ThF4 would be least stable
fS
S1
(OCP + CaC03AT
pCO,.3.52)
-16-
-18
Figure 3-24. The relative stability of various thorium solids in equilibrium
with Variscite and Gibbsite (V & G) Dicalcium Phosphate Dihydrate
(OCPD) and Octacalcium Phosphate (OCP)
The relative activity of solution species of thorium in equilibrium with
Th02(s) at assumed activities of various ions is plotted in Figure 3-25. The
thermodynamic data for all the hydrolysis species except Th(OH)I were obtained
from Baes and Mesmer (1976). The data for all the other species and Th(OH)"
3-213
-------
were obtained from Sillen and Mar-tell (1964). In general, the total concen-
tration of thorium in solution decreases with an- increase of pH from zero to
5. Above 5, pH does not affect thorium concentration in solution due to the
formation of Th(OH)4< The activity of all positively charged species decreases
with an increase in pH, while the activity of the negatively charged species
increases with the increase in pH. Under the conditions assumed for Fig-
ure 3-25, the total activity of thorium in solution would be expected to be
approximately 10
-6
-9.6
moles/liter above pH 5.
8*
10
Figure 3-25. Activity of various thorium species in soil in equilibrium with
ThOgCs), pCT » pN03- = pS042" = 3.0, pF" = 4.5 and pH2P04~ =
o. u •
- 23
In addition to OH , thorium forms various complexes with SO* , PO, ,
Cl~, N0^~, and F". Various anions in increasing order of their importance to
- - 2
forming complexes are: NO^ , Cl , ^^^ , SO^ , and F . Figure 3-25 shows
that thorium exists as Th only in very acidic solutions (pH < 3). Above
4+
pH 3, Th hydrolyzes very .rapidly and it does not contribute significantly
3-214
-------
to the total thorium concentration. Under the conditions assumed for Fig-
ure 3-25, ThF * would be a dominant solution species at pH < 5, and Th(OH)4°
would be a dominant solution species at pH > 5. If fluoride ion is absent
?+
from the solution, or its concentration is extremely low, Th(OH)2 would
mainly control the thorium concentration in solution at pH < 5.
Experimental Adsorption Results
Schulz (1965) found the thorium in soils to be strongly adsorbed by clay
particles or present as insoluble oxides and hydroxides. Rubtsov (1966, 1972)
found thorium to associate with the fine-grained particles during soil weather-
ing. Katsurayama (1968) determined the distribution coefficient of thorium
but data are not presented in the available abstract.
Nishiwaki et al. (1972) spiked seawater and seawater distilled water mix-
tures with Th and measured the adsorption on a medium sand, very fine sand
and silt-clay. Twenty grams of soil were contacted with 4 liters of spiked
water and mixed until equilibrium was reached. The Kd for thorium increased
as the particle size of the soil decreased. Chlorosity of the water did not
appear to consistently affect the thorium Kd for the fine sand or silt-clay.
The Kd for the medium sand increased as the ch'lorosity of the water decreased.
The chlorosity effect was compounded by a variable pH of the various salt solu-
tions, so that the exact cause of the trend was not determineble. Kd values
for the medium sand, very fine sand and silt-clay were 40 to 130, 310 to 470,
and 2700 to 10,000 ml/g, respectively.
Rancon (1973) measured the thorium Kd for a soil developed on a schist
consisting of quartz and clay with no calcite or organic matter, for a mixed
quartz-clay-calcite-organic matter soil and for illite with 100 mg Th/1 versus
solution pH. For the quartz-clay soil, at pH 6 the Kd was 5 x 10 ml/g, at
pH 4 the Kd was about 1 x 10 ml/g and at pH 2 the Kd was about 5 ml/g. The
mixed quartz-clay-calcite-organic matter soil could not be lowered in pH with-
out removal of soil calcite, but above pH 8, the thorium Kd dropped from
10 ml/g to 100 ml/g at pH 10. Dissolution of humic acids in the soil probably
resulted in thorium complexation and a decreased Kd with rising pH. Illite
behaved similarly to the quartz-clay soil, but the thorium Kd at pH 1 was about
500 ml/g and about 1 x 10 ml/g at pH 6.5. For soils without calcite or organic
material, the thorium Kd decreased as the thorium concentration in solution
3-215
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initially increased. Calcareous soils neutralized even high strength thorium
solutions to precipitate Th(OH)4. The quartz-clay soil and illite, for example,
gave thorium Kd values of 8 ml/g and 120 ml/g, respectively, in a 1 g Th/1
solution, and 60 ml/g and 1000 ml/g, respectively, in a 0.1 g Th/1 solution.
The drop in thorium Kd was caused by saturation of available exchange sites as
a result of increased thorium concentration. There was evidence for the con-
centration dependence of the thorium Kd down to 1 mg Th/1 in the initial con-
tacting solution. In general, three types of soil-thorium adsorption reactions
were found: 1) Th(OH)4 precipitation as a result of soil calcite buffering,
2) strong adsorption on clay-containing soils and dilute thorium (<1 g/1) solu-
tions at a pH above 2, and 3) strong adsorption on organic-containing soils at
the neutral to acid pH range, but diminishing adsorption into the alkaline pH
range.
Bondietti (1974) studied the adsorption of hydrolyzed thorium from waters
at pH 6.5 by calcium saturated reference clays (montmorillonite and kaolinite)
and calcium humate and found 95% and 99.9+% adsorption, respectively. Desorp-
tion studies utilizing calcium citrate removed 10 to 30% of the thorium from the
clays but only 1% from the humate. Stronger complexers (DTPA and EDTA) removed
20 to 30% of the thorium from the humate. A mixed organic-hydroxy complex was
proposed for the reaction of thorium with humic substances.
Migration Results
Field Studies—
i
The thorium content of groundwater was reported by Dementyev and
Syromyatnikov (1965) to be highest in low salinity, low hardness, low pH, high
organic content groundwaters. These characteristics suggest transport of
thorium as colloidal suspensions and anionic complexes involving soil acids.
From a fresh granodiorite containing 9.3 ppm thorium and 2.5 ppm uranium,
the first stages of weathering resulted in apparent removal of 25% of the
thorium and 60% of the uranium (Hansen and-Huntington, 1969). An acid leach
of the fresh rock removed 90% of the thorium and 60% of the uranium indicating
that most of the thorium and uranium are in acid soluble or interstitial mate-
rials. After an initial drop in concentration, the total uranium and thorium
content of the weathered rock increases by at least a factor of 4 in the upper-
most weathered material. Leaching studies showed that thorium was associated
with clays formed during weathering and with accessory minerals such as zircon.
' 3-216
-------
Hansen (1970) reported that when freed from minerals by weathering,
thorium was leached comparatively slowly. From a fresh granodiorite containing
9.3 ppm thorium, the first stages of weathering resulted in apparent removal of
25% of the thorium. An acid Teach of the fresh rock removed 90% of the thorium
indicating that most of the thorium was in acid soluble or interstitial mate-
rials. The thorium content of the weathered rock increased by a factor of 4
due mostly to association with the clays formed during weathering.
Laboratory Studies—
Desai and Ganguly (1970) showed humic acids from a coastal marine sedi-
ment solubilized 100% of the thorium added to an ammonia solution (2.5N).
Thorium in this solution without humic acid was observed to predominantly
precipitate (95%). The humic acid-thorium complex was noncationic. In an
identical experiment, fulvic acid extract was shown to solubilize 59% of
thorium added to an ammonia solution. Again, the solubilized organic-thorium
fraction was noncationic.
Summary
Under alkaline conditions, Th(OH). and ThO^ maintain low activities in
soil solutions (Figures 3-24 and 3^25) and these compounds could form and
govern thorium concentration. Thorium' hydrolyzes readily even in moderately •
acidic environments (Figure 3-25) so that Th would be present only in very
acidic solutions. Laboratory studies also show that thorium tends to precipi-
tate as thorium hydroxide and hydrated thorium oxide in soils (Schulz, 1965;
Rancon, 1973).
An increase in thorium content with increase in CaCO^, phosphate and
humus content of soils and sediments has been reported (Kuznetsov et al., 1968;
Pashneva et al., 1965; Menzel, 1968; Yakobenchuck, 1968; Hansen and Huntington,
1969; Pokidin et al., 1972). However, Tyuryukanova and Kalugina (1971) reported
low thorium concentrations in high humus soils (peats and forest podzols) com-
pared with alluvial soils. Thorium adsorption increases with increase in pH
(Rancon, 1973) and decrease in soil particle size (Hansen and Huntington, 1969;
Hansen, 1970; Nishiwaki et al., 1972; Rubtsov, 1966, 1972; Bondietti, 1974).
Strong humic and fulvic acid complexes with thorium occur in the neutral to
acidic range (Rancon, 1973; Bondietti, 1974) which are noncationic (Desai and
Ganguly, 1970) and mobile. It has been reported also that thorium migrates
primarily in the colloidal form (polymeric) in the natural environment
(Baranov et al., 1956; Lazarev et al., 1961; Kimura et al., 1968).
3-217
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References
Adams, J. A. S., J. K. Osmond, and J. J. W. Rogers. 1959. The Geochemistry
of Thorium and Uranium. IN; Physics and'Chemistry of the Earth. 3.
L. H. Ahrens, F. Press, K. Rankama, and S. K. Runcorn (eds.), pp. 298-348.
Adams, J. A. S. and K. A. Richardson. 1960. Thorium, Uranium, and Zirconium
Concentrations in Bauxite. Econ. Geol. 55:1653.
Adams, J. A. S. and C. E. Weaver. 1958. Thorium-to-Uranium Ratios as Indi-
cators of Sedimentary Processes — An Example of Geochemical Facies. Bull.
Am. Assoc. Petrol. Geol. 42:387.
Ahrens, L. H. 1952. The Use of lonization Potentials. Part I. Ionic Radii
of the Elements. Geochim. et Cosmochim. Acta. 2:155. -
Baes, C. F., Jr. and R. F. Mesmer. 1976. The Hydrolysis of Cations. John
Wiley and Sons, New York.
Baranov, V. I., A. B. Ronoy, and K. G. Kinashova. 1956. On Geochemistry of
Dispersed Thorium and Uranium in Clays and Carbonate Rocks of Russian Platform.
Geokhimiya No. 3:3-8.
Billings, G. K. 1962. A Geochemical Investigation of the Valley Spring Gneiss
and Packsaddle Schist, Llano Uplift, Texas. Texas J. Sci. 9:328.
Bondietti, E. A. 1974. Adsorption of Pu(IV) and Th(IV) by Soil Colloids.
Agronomy Abstracts:
Dementyev, V. S. and N. G. Syromyatnikov. 1965. Mode of Occurrence of Thorium
Isotopes in Ground Waters. Geokhimiya. No. 2:211-218.
Desai, M. V. M. and A. K. Ganguly. 1970. Interaction of Trace Elements with
Organic Constituents in the Marine Environment. BARC-488, p. 102.
Hansen, R. 0. 1970. Radioactivity of a California Terrace Soil. Soil
Science. 110:31-36.
Hansen, R. 0. and G. L. Huntington. 1969. Thorium Movements in Morainal
Soils of the High Sierra, California. Soil Science. 108:257-265.
Heier, K. S. and J. L. Carter. 1964. Uranium, Thorium, and Potassium Contents
in Basic Rocks and Their Bearing on the Nature of the Upper Mantle. IN;
J.A.S. Adams and W. M. Lowder (eds.). The Natural Radiation Environment, p. 75.
Heier, K. S. and J. J. W. Rogers. 1963. Radiometric Determination of Thorium,
Uranium, and Potassium in Basalts and in Two Magmatic Differentiation Series.
Geochim. et. Cosmochim. Acta. 27:137.
Holland, H. D. and J. L. Kulp. 1954. The Mechanism of Removal of Thorium
and Radium from the Oceans. Geochim. et Cosmochim. Acta. 5:214-224.
3-218
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Katsurayama, K. 1968. Accumulation of Radioactivity in Rice Fields. Annu.
Rep. Res. Reactor Inst. Kyoto University. 1:293-97.
Kimura, Y., H. Morishuma, T. Koga, H. Kawai and Y. Honda. 1968. Studies on
the Behavior and Distribution of Radioactive Substances in Coastal and
Estuarine Waters. Kinki Daigaku Genshiryaku Kenkyusho Nonpo. 7:21-31 (in
Japanese).
Kuznetsov, Y. V., Z. N. Simonyak, A. P. Lisitsyn, and M. S. Frenklikh. 1968.
Thorium Isotopes (230ih, 232fh) in the Surface Layer of the Indian Ocean
Sediments. Geochem. Int. 5:169-77.
Lazarev, K. F., D. S. Nikolaev, and S. M. Grashchenko. 1961. Concentration
of Thorium Isotopes in Sea Water. Radiokhimiya. 3:623-35 (in Russian).
Menzel, R. G. 1968. Uranium, Radium, and Thorium Content in Phosphate Rocks
and Their Possible Radiation Hazard. J. Agr. Food Chem. 16:231-34.
Murray, E. G. and J. A. S. Adams. 1958. Amount and Distribution of Thorium,
Uranium, and Potassium in Sandstones. Geochim. et. Cosmochim. Acta. 13:260.
Nishiwaki, Y., Y. Honda, Y. Kimura, H. Morishima, T. Koga, Y. Miyaguchi, and
H. Kawai. 1972. Behavior and Distribution of Radioactive Substances in
Coastal and Estuarine Waters. IN: Radioactive Contamination of the Marine
Environment. IAEA-SM-158/11, pp7 177-193.
Pashneva, G. E., T. P. Slavnina, and V. V. Serebrennikov. 1965. Rare Earth
and Thorium Content in Soils of Tomsk Region. Izv. Sibirsk. Otd. Akad. Nauk
SSSR, No. 4 Ser. Biol.-Med. Naulc. No. 1:48-52 (in Russian). •
Pliler, R. 1956. The Distribution of Thorium and Uranium in Sedimentary
Rocks and the Oxygen Content of the Precambrian Atmosphere. Thesis. Rice
Institute.
Pokidin, V. K., Y. V. Kuznetov, E. A. Prozorovich, and F. A. Asadullaeva.
1972. Radioactivity and Rate of Sediment Formation in the Caspian Sea.
Geokhimiya. No. 7:834-43 (in Russian).
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. IAEA-SM-172/55, pp. 333-346
(in French).
Rogers, J. J. W. 1964. Statistical Test of the Homogeneity of the Radio-
active Components of Granite Rocks. IN: J.A.S. Adams and W." M. Lowder (eds.),
The Natural Radiation Environment, p.~Tl.
Rogers, J. J. W. and J. A. S. Adams. 1974. Thorium Solubilities in Water.
Handbook of Geochemistry, K. H. Wedepohl, Ed., Vol. II-4, p. 90-H-l, Springer-
Verlag, New York.
3-219
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Rubtsov, D. M. 1966. Distribution of Thorium in Various Soils.
Pochvovedenie. No. 3:55-67 (in Russian).
Rubtsov, D. M. 1972. Thorium and Radium Content in the Silt Fraction of the
Podzolic Mountain Soils of Thin Forests. Radioekologicheskie Issledavaniya
Prirodnykh Biogeotsenozakh. IN: Verkhovskaya (ed.) Izdatel'stov Nauka,
pp. 42-53.
Ryabchikov, D. I. and E. K. Golbraikh. 1969. Analytical Chemistry of
Thorium. Translated by A. Aladjem. Ann Arbor-Humphrey Science Publishers.
Schulz, R. K. 1965. Soil Chemistry of Radionuclides. Health Physics.
11:1317-24.
.Si 11 en, L. G. and A. E. Martell. 1964. Stability Constants of Metal-Ion
Complexes. Special Publication No. 17. The Chemical Society, London.
Tyuryukanova, E. B. and V. A. Kalugina. 1971. The Behavior of Thorium in
Soils. Soviet Journal of Ecology. 2:467-469.
Vinogradoy, A. P. 1959. The Geochemistry of Rare and Dispersed Chemical
Elements in Soils. Consultants Bureau, Inc.
Yakobenchuk, V. F. 1968. Radioactivity and Chemical Properties of Sod-Pod-
zolic Soils in the Ukrainian Western Polesie. Visn. Sil's1 Kogosped. Nauki.
11:45-50 (in Ukranian).
3-220
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TRITIUM
Natural Soil and Rock Distributions
Natural tritium concentrations vary geographically from about 10 T.U. for
northern and southern midlatitude precipitation to 1 T.U. for equatorial pre-
cipitation. The background tritium content of deep ocean and groundwater is
essentially zero (Fairbridge, 1972). Hence, the tritium content of rocks,
minerals and deeper sediments also is nearly zero unless contacted by tritium-
contaminated groundwater.
Brief Chemistry
Tritium ( H or T) is an isotope of hydrogen with a half-life of 12.26 years.
Tritium occurs naturally in the upper atmosphere, chiefly by the interaction
of fast neutrons with nitrogen (Kaufman and Libby, 1954). The amount of natu-
ral tritium is about one part in 10 parts of normal hydrogen ( H) (Cotton
and Wilkinson, 1962).
It was a fairly recent discovery that tritium also is a product of uranium
fission (Albenesius, 1959; Albenesius and Ondrejcin, 1960). Several years
prior to the identification of tritium as a uranium fission product, the
Environmental Monitoring Group at. Savannah River plant recognized tritium as
occurring in Purex separations process wastes, and was using it as a ground-
water tracer (Norton and Ross, 1960). Tritium yields of 3.8 to 20.3 curies/
ton uranium were obtained for exposures between 300 and 1600 megawatt-days/
ton uranium.
Tritium oxidizes rapidly to HTO, existing essentially as water, and its
distribution in ground and surface waters is controlled by the operation of
the hydrological cycle. Tritium is measured in tritium units, T.U., where
18
T.U. = T/10 H, or expressed as picocuries/liter. Natural background levels
have risen since weapons testing of nuclear devices from 1 to 10 T.U. to
several hundred T.U. up to the 1963 atmospheric detonation moratorium. Most
of the tritium added to the hydrological cycle by fusion devices was due to
residual tritium H5.7 x 10 Ci H/megaton fusion device, Stead, 1963).
3-221
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Solid Phase and Solution Equilibria
Tritium (H ) is a radioactive isotope of hydrogen. Thus, tritium behavior
in soils would be expected either to be similar to hydrogen or to exist as an
ion, gas and liquid (tritiated water). Except for the slight differences in
vapor pressure, tritiated water behaves the same as ordinary water (Jacobs,
1974). Although no discrete solid phases of tritium are expected in soils,
it could associate itself with soil organic components containing hydrogen or
with some soil minerals as mobile water, water of hydration, or as part of
structural hydroxyl groups (Jacobs, 1974; Stewart, 1967).
Tritium rapidly travels at about the same velocity as the soil water or
groundwater in the form of HTO. Some replacement of nontritiated water on
clays and other hydrated soil constituents occurs, but the reaction is of
marginal value as far as tritium retention is concerned. Tritium can move in
a vapor phase through the soil under certain conditions as well.
Experimental Adsorption Results
Tritium ultimately exists in the soil as a tritiated water molecule.
Theoretically, tritium ions are capable of exchanging for hydrogen and other
ions on soil. However, the. tritium is usually in an aqueous solution before
contacting soils, so that isotopic exchange and replacement of hydroxyl 'water
molecules is the most common mechanism of tritium removal. Little quantita-
tive work on tritium distribution has been reported. Rabinowitz et al. (1973)
have used an electrodialysis technique to speed up tritium hydroxyl exchange
in clays. They reported that tritium hydroxyl exchange was proportional to
cation exchange capacity of the clays used in the study: illite, 35 meq/100 g
and 93 pCi3H/ml; kaolinite No. 4, 10 meq/100 g and 36 pCi3H/ml; kaolinite No. 9,
2 meq/100 g and 12 pCi H/ml.. Most of the field studies indicated that tritium
distribution coefficients were very low because tritium migration velocities
were essentially the same as those of the .accompanying groundwater (Brown and
Haney, 1964; Brown, 1967).
3-222
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Migration Results
Field Studies—
Haney, Brown and Reisenauer (1962) and Brown and Haney (1964) estimated
by studying the rate and direction of groundwater flow that movement of tri-
tium disposed to the ground at the Hanford Purex Plant would require 6 or
7 years travel time to the Columbia River, approximately 17 miles from the
Purex Plant. The half-life of tritium allowed about 70% of the total tritium
discharged to reach the river. The tritium content of Columbia River water
was not expected to rise over 50% higher than the present background because"
of tritium dispersion during travel and the large dilution represented by the
Columbia River (80,000 to 200,000 ft3/sec).
Brown (1967) reviewed the hydrology and geology of the Hanford area, and
showed the 1966 limits of the areas defined by 3000 pCi/cm and 10 pCi/cm
3 3
of tritium based on well water analytical results. The 10 pCi H/cm area was
3 km from the Columbia River in 1966. The concentration of tritium added at
3 3
the Purex Plant disposal sites was 10,000 pCi H/cm . Tritium moved at essen-
tially the same rate as the groundwater.
' Kline and Jordan (1968) reported a field experiment on a clay Puerto
Rican soil where 1 liter of 20 mCi/1 tritiated water was applied to a plot
2
with about 1 m area, and sampled via a lysimeter installed horizontally,
without disturbing the plot surface soil, 18 cm below the soil surface. The
sampling program lasted 210 days during which time 137 cm of rain fell on the
plot. Runoff surface water was collected as well. The soil remained satu-
rated, or close to saturation for the sampling period. Most of the tritium
passed through the profile as a peak or front 16 days after application, and
declined in concentration exponentially during the remainder of the experi-
ment. The surface runoff peaked in 2.9 days, with another change in decline
of specific activity at 35.6 days. Because of the loss rate curve of tritium
in the free soil water, the authors suggested that tritium movement through
soils must be modified to allow for the existence in clay soils of isolated
compartments of immobile water that does not have free interchange with the
more rapidly moving water.
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Jordan et al. (1971) monitored tritium movement through a soil near
Argonne, Illinois. They concluded from the results that a pulse of tritium
moves downward at a rate determined by precipitation entering the top of the
soil column. The tritium peak broadens and flattens from diffusion and will
become immobilized if the upper soil layers dry out. Loss of tritium occurred
both by evapotranspiration and deep drainage. Deviation of the field results
from a model of tn'tiated water movement proposed by Sasscer et al. (1971)
was explained as due to parameters that were difficult to quantify such as
root holes and hydroxyl exchanges.
Ehhalt (1973) pointed out that tritiated molecular hydrogen in the atmo-
sphere was oxidized by soil microorganisms. He also estimated from the tritium
content of H, of 2 x 10 T.U., steady since 1962, that the HT input rate was
2
about 5.4- T atoms/cm sec into the soil from the action of soil microorganisms.
Purtymun (1973) investigated the underground movement of tritium from
solid waste storage shafts in rhyolite tuff at Los Alamos. Levels of 100 pCi
H/ml had moved a distance of 105 ft in 4 years. Core samples of tuff were
collected, water distilled from them and tritium in the water determined.
Asphalt coatings on containers and shaft walls was suggested to control
tritium migration from the shafts.
In contrast, Wheeler and Warren (1975) reported that a later comparison
between asphalted and nonasphalted shafts showed that the asphalt containment
techniques used were ineffective in slowing tritium migration from the shafts.
Double containment and complete encapsulation of the 210-liter disposal drums
in roofing asphalt was suggested to slow tritium movement from the storage
shafts.
Laboratory Studies—
Because soil water and groundwater systems are similar, it would be
expected that tritium oxide would move through these systems at essentially
the same rate as light water. However, in comparing chloride ions arid tritium
breakthrough curves for various soil types, the tritium curve lags behind
that for chloride ions in soils containing clays and silt (Kaufman and Orlob,
1955). This behavior was reported as caused by the tritiated water molecule
entering the clays and replacing nontritiated water adsorbed on clay surfaces
or inter!ayers.
3-224
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Haney (1963, 1964) reported a laboratory study where actual groundwater
containing complexed Ru and tritium was passed through several Hanford
project soils. In silty soils, the ruthenium breakthrough curve was slightly
behind that for tritium, but in sandy soil columns, the breakthrough curves
were essentially identical. Several retention mechanisms of tritium on clays
are possible. For example: 1) exchange of tritium ions for hydrogen ions in
exchange positions, 2) tritiated water molecules exchanged for hydroxyl water
or water of crystallization, 3) exchange of tritium ions for exchangeable
cations other than hydrogen, and 4) replacement of aluminum in lattice sites
by tritium, with the aluminum ions moving into exchangeable cation positions.
In the last case, the tritium may be more or less "fixed" in its octahedral
position. This "fixation" has been considered to be minimal by some investi-
gations (Higgins, 1959; Halevy, 1964; Corey and Morton, 1968; Teller et al.,
1968).
It has been hypothesized by others (Corey and Fenimore, 1968) that the
chloride ions are repelled from negatively charged soil particles into the
central region of the soil pore where groundwater velocity is greatest (anion
exclusion). This would be most effective in high cation exchange capacity
soils which do not retain anions. Also, Corey and Fenimore (1968) showed that
chloride ions lag far behind tritium in acid kaolinitic soil due to anion
exchange of the chloride ions.
Corey and Morton (1968) reported no appreciable tritium "fixation" upon
investigation of the movement of tritiated and deuterated water through
acidic, kaolinitic soil. The work of other investigators, however, indicates
that hydrogen isotopes may be selectively fixed (Rosenqvist, 1963; Koranda,
1965; Clayton et al., 1966; Savin, 1967; Stewart, 1967). The work of
Rabinowitz (1969) and Rabinowitz et al. (1973) with forced exchange of tri-
tiated water on clays also supports the view that part of the tritium can be
"fixed" on clays. The isotopic exchange and fixation was forced with a mild
form of electrodialysis (2 V/cm) resulting in an increase in the reaction
rate of about five times the rate without electrodialysis. Aluminum and
other basing cations were determined after each experiment. The presence of
+3
exchangeable Al correlated with tritium loss from the water resulting from
tritium "fixation". A mass balance on the tritium also was determined for each
3-225
-------
experiment. The analytical data for tritium facilitated the differentiation of
isotopic exchange and the "fixation" of tritium by kaolinite and ill He.
Summary
Tritium substitutes readily for the hydrogen in water and thus becomes a
part of the hydrological cycle (Jacobs, 1974). Migration of tritium through
the groundwater takes place at the same velocity as the groundwater through
sandy soils (Haney 1963, 1964; Haney et al., 1962; Brown and Haney, 1964;
Brown, 1967). Some investigators have reported the selective fixation of
tritiated water on clays and other hydrated minerals (Rosenqvist, 1963;
Koranda, 1965; Clayton et al., 1966; Savin, 1967; Stewart, 1967; Rabinowitz,
1969; Rabinowitz et al., 1973) while other investigators have considered tri-
tium fixation to be minimal (Higgins, 1959; Halevy, 1964; Corey and Horton,
1968; Teller et al., 1968). However, all of the field studies have indicated
that tritium movement is synonomous with water movement (Brown, 1967).
References
Albenesius, E. L. 1959. Tritium as a Product of Fission. Physics Review
Letters. 3:274.
Albenesius, £. L. and R. S. Ondrejcin. 1960. Nuclear Fission Produces
Tritium. Nucleonics. 18:100.
Brown, D. J. 1967. Migration Characteristics of Radionuclides Through
Sediments Underlying the Hanford.Reservation. ISO-SA-32.
Brown, D. J. and W. A. Haney. 1964. Chemical Effluents Technology Waste
Disposal Investigations July-December, 1973 - The Movements of Contaminated
Ground Water from the 200 Areas to the Columbia River. HW-80909.
Clayton, R. N., I. Friedman, D. L. Graaf, T. K. Mayeda, W. F. Meents and
N. F. Shimp. 1966. The Origin of Saline Formation Waters. I. Isotopic
Composition. J. Geophys. Res., 71, 3869-3882.
Corey, J. C. and J. W. Fenimore. 1968. Tracing Groundwater with Chloride
Ions and Tritium Through Acid Kaolinitic Soil. Int. J. Appl. Rad. Isotopes.
19:741-746.
Corey, J. C. and J. H. Horton. 1968. Movement of Water Tagged with 2H, 3H
and '°0 Through Acidic Kaolinitic Soil. Soil Sci. Soc. Am., Proc. 32:471-475.
Cotton, F. A. and G. Wilkinson. 1962. Advanced Inorganic Chemistry. Inter-
science Publishers.
Ehhalt, D. H. 1973. On the Uptake of Tritium by Soil Water and Ground Water.
Water Resources Research. 9:1073-1074.
3-226
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Fairbridge, R. W. (ed.). 1972. The Encyclopedia of Geochemistry and Environ-
mental Sciences. Van Nostrand Reinhold Company.
Halevy, E. 1964. The Exchangeability of Hydroxyl Groups in Kaolinite.
Geochim. et Cosmochim. Acta. 28:1139-1145.
Haney, W. A. 1963. Fission Product Tritium in Fuel Reprocessing Wastes.
HW-SA-3303.
Haney, W. A. 1964. Consequences of Activity Release. Nuclear Safety.
5:399-403.
Haney, W. A., D. J. Brown, and A. E. Reisenauer. 1962. Fission Product
Tritium in Separations Wastes and in the Ground-Water. HW-74536.
Higgins, G. H. 1959. Evaluation of the Ground-Water Contamination Hazard
from Underground Nuclear Explosions. J. Geophys. Res. 64:1509-1520.
Horton, J. H. and D. I. Ross. 1960. Use of Tritium from Spent Uranium Fuel
Elements as a Ground-Water Tracer. Soil Science. 90:267-271.
Jacobs, D. C. 1974. Impacts in Groundwater of Effluents Arising in the
Nuclear Industry. NP-20456, pp. vp., Paper 16.
Jordan, C. F., J. R. Kline, and D. S. Sasscer. 1971. Tritium Movement in an
Old-Field Ecosystem Determined Experimentally. IN: Radionuclides in Eco-
systems. CONF-710501-P1, pp. 199-203.
Kaufman, S. and W. F. Libby. 1954. The-Natural Distribution of Tritium.
Phys. Rev. 93:1337.
Kaufman, W. J. and G. T. Orlob. 1955. Measuring Ground Water Movement with
Radioactive and Chemical Tracers. Paper presented at the Annual Meeting of
the California Section of AWWA, Sacramento.
Kline, J. R. and C. F. Jordan. 1968. Tritium Movement in Soil of Tropical
Rain Forest. Science. 160:550-551.
Koranda, J. J. 1965. Preliminary Studies of the Persistence of Tritium and
14c.in the Pacific Proving Ground. Health Phys., 11, 1445-1457.
Purtymun, W. D. 1973. Underground Movement of Tritium from Solid-Waste
Storage Shafts. LA-5286-MS.
Rabinowitz, D. D. 1969. Forced Exchange of Tritiated Water with Natural
Clay. MS Thesis, N. Mex. Inst. of Min. and Tech., Socorro, NM.
Rabinowitz, D. D., C. R. Holmes and G. W. Gross. 1973. Forced Exchange of
Tritiated Water with Clays. IN; Trftium, Messenger Graphics Pub., Phoenix
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Rosenqvist, I. T. 1963. Studies in Position and Mobility of the H Atoms in
Hydrous Micas. IN_: Clays and Clay Minerals, Monograph No. 13, p. 11-135,
Macmi11 an Co., NY.
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URANIUM
Natural Soil and Rock Distributions
The range in abundance of uranium in natural rocks is given in Table 3-96.
Vinogradov (1959) reported 1 ppm uranium as the average content of uranium in
soils. The oxidation of organic matter in black shales tends to precipitate
U(IV). Consequently, the black shales usually contain more uranium than red,
green or gray shales.
Brief Chemistry
There are 14 known isotopes of uranium from U to U and one uranium
235
isomer ( U) (Fairbridge, 1972). The half-lives of the uranium isotopes vary
from a few minutes to over 4 billion years. Only three of the isotopes occur
235 238
naturally, and two of these ( U and U) are parents of series that end in
lead isotopes. The nuclear data on the three natural uranium isotopes are
given in Table 3-97. The uranium isotopes of interest in waste disposal are
given in Table 3-98. The common oxidation states of uranium are U(III), U(IV),
U(V) and U(VI) (Udaltsova, 1963). U(III) is easily oxidized in air to U(IV).
3-228
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TABLE 3-96. ABUNDANCES OF URANIUM IN NATURAL MATERIALS
(ADAMS ET AL., 1959; CLARK ET AL., 1966)
Material
U Concentration Range, ppm
Igneous Rocks
Dunites and Peridotites
Gabbro and Diabase
Intermediate (diorite and
quartz diorite)
Sialic (granite, syenite,
monaozite)
0.001-
0.3-
0.1-1
•0.8
•3.4
-11.0
0.15-21.0
Sedimentary Rocks
Black shales
Red, gray and green shales
Orthoquartzite
Limestone and dolomite
Bentonite
Bauxite
Halite
Anhydrite
Metamorphic Rocks
Marble
Slate
Phyllite
Schist
Gneiss
Amphioolite
Granulite
0.11-
1.2-
1.0-
1.8-
4.5-
2.6-
3.2-
•0.24
•6.1
•2.7
•2.9
•15.0
4.1
7.0
Median 0.5
Median 1.7
Median 3.9
3.0-25.0
1.2-12.0
0.2-0.6
0.1-9.0
1.0-21.0
3.0-27.0
0.01-0.02
0.25-0.43
Median 8.0
Median 3.2
Median 0.45
Median 2.2
Median 5.0
Median 8.0
Median 0.013
Median 0.37
The U(IV) state is fairly stable in aqueous solutions if they are very acidic.
Uranium (V) disproportionates. to U(IV) and U(VI):
+ 6H20 (Udaltsova, 1963). U(VI), as UO*2 (Uranyl) at pH < 2.5, is the most
40"
H
3-229
-------
TABLE 3-97. NUCLEAR PROPERTIES OF NATURAL URANIUM
ISOTOPES (FAIRBRIDGE, 1972)
Half-Life, Decay Series
Isotopes % Abundance yr Parent
234 • R
"4U 0.0056 2.48 x 10b
235U 0.72 7.13 x 108 4N + 3
238U 99.27 4.51 x 109 4N + 2
TABLE 3-98. URANIUM RADIONUCLIDE DATA
(WEAST, 1976)
Isotope Half-Life Decay Mode
232
"ni 73.6 years a, SF
U 162,000 years a
27A
"\l 247,000 years a
235U 7.13 x 108 years a, SF
- 236U 2.39 x 107 years ct, SF
237U 6.75 days . 6"
238U -' 4.51 x 109 years a, SF
stable state of uranium in aerated aqueous solutions (Seaborg and Katz, 1954).
At higher pH values, hydrolyzed uranyl ions predominate. Approximately
103 uranium minerals have been confirmed. Uraninite (ideally U02) is the
primary ore mineral of uranium, but other secondary minerals include carbonates,
sulfates, molybdates, phosphates, vanadates, silicates and multiple oxides.
Uranyl ions readily form many complexes with anions ordinarily found in soil-
water environments such as carbonate, sulfate and fluoride.
Solid Phase and Solution Equilibria
The thermodynamic data for 1) U03> Na2U04 and U02C03 were obtained from
Garrels and Christ (1965), 2) U02(OH)2> U02(OH)2H20 and Na4U02(C03)3 were
obtained from Sillen and Martell (1964), and 3) the remaining species reported
in Figure 3-26 were obtained from Palei (1970). The relative stability of
several uranium solid phases, in terms of the uranyl ion activity produced by
each, is shown in Figure 3-26. Since all the compounds shown in Figure 3-26
3-230
-------
2+
are of U(VI), and are plotted as a function of U02 , their curves would not
move with changes in oxidation-reduction conditions. U(IV) compounds, such
as UCL, are very soluble in an oxidizing environment and fall outside the
boundary of Figure 3-26. With an increase in reducing conditions, U(IV) com-
pounds would become more stable. At a pO« of >71, the U02 curve would fall
just below the UOJW.PO. curve. U(VI) compounds are stable in an oxidizing
environment and U(IV) compounds are stable in a reducing environment. Consis-
tent with the thermodynamic data is the observation that 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.
-2
f
-8
-10
-12
10
PH
Figure 3-26.
The relative stability of various uranium solids in an oxidizing
soil environment [pOgtg) = 0.68 atm], PC02(g) = 3-52 atm, pK+ =
pNa* = PNH4"1" = 3.0 and phosphate levels in equilibrium with "
Variscite and Gibbsite.
The thermodynamic data for UO^F and UO^PO* were selected from Palei
(1970) and selected data from Si 11 en and Martell (1964) were used for the
remaining species in Figure 3-27. All the species except U02 shown in Fig-
ure 3-27 are of U(VI). Since equilibrium is assumed with Na2U04, a U(VI) com-
pound, a change in oxidation-reduction conditions will affect the position of
the curves. In an oxidizing environment (p02 = 0.68 atm), the U(IV) species
are very low in concentration (log a^ <38) and fall outside of the boundaries
of Figure 3-27.
3-231
-------
CO
s*
Figure 3-27. Activity of various uranium species in equilibrium with
in an oxidizing soil environment [p02(q) = 0.68 atm], pCOpfn)
- 3.52 atm, pCl" = pS042- = 3.0, pF~ = 4T5 and pH2P04 = 5.0.
Figure 3-27 shows that U (VI) species will control the solution concen-
2+
tration in an oxidizing environment. UO? is the predominant solution species
up to a pH of approximately 6. The predominant solution species over pH
ranges of from 6 to 8 and >8 are UOgtOHK" and U02(CO,), , respectively.
Experimental Adsorption Results
Uranium adsorption studies have been performed by several investigators.
Szalay (1954, 1957) showed that uranium adsorption by decomposing plant debris,
peat, lignite and brown coal is quite high. He determined that humic substances
3-232
-------
in these materials were responsible for the adsorption which was described as
an ion exchange-like sorption. Adsorption isotherms for the humic acid were
measured.
Goldsztaub and Wey (1955) determined that 7.5 g of uranium was adsorbed
from a 1% uranyl nitrate solution per 100 g of calcined montmorillonite and
2.0 g of uranium per 100 g of calcined kaolinite.
Manskaya et al. (1956) described the adsorption of uranium on fulvic
acids as a function of pH. The curve of percent adsorption versus pH showed
a maximum at pH 6 of about 90% uranium removal. At pH 4 and pH 7, uranium
adsorption was down to 30%.
Starik et al. (1958) showed a similar adsorption curve of pH versus
uranium adsorption on ferric hydroxide. The maximum of the adsorption curve
was at pH 5 with about 50% uranium adsorption, and rapidly decreased above
and below pH 5.
Rozhkova et al. (1959) showed similar uranium adsorption curves versus
pH for lignite and humic acids. The curve maxima occurred between pH 5 and
6; Dementyev and Syromyatnikov (1968) showed that these adsorption curve
maxima result because the pH 6 region is a-boundary between anionic and
cationic uranium forms and corresponds to: [UO^2] + LUOgOH"1"] ~ [UO^CO^2],
an equality between cation and anion uranium forms in solution.
Horrath (1960) measured the enrichment factor for the adsorption of
uranium by peat and obtained an average of 200 to 350. Although it is not
possible to determine what methods or exactly what the enrichment factor means
from the English translated abstract, it is assumed that the data represent a
uranium Kd of 200 to 350 by weight or volume.
Kovalevskii (1967) found the uranium content of noncultivated soils in
western Siberia increased with the clay content of the soils. Clay soils con-
tained at least three times as much uranium as sands. Yakobenchuck (1968)
correlated the total uranium content in Russian sodpodzilic soils from the
Ukraine with other soil constituents. Uranium showed correlation with the
oxidizes of silicon, iron, and aluminum suggesting coprecipitation or
inclusion.
3-233
-------
Masuda and Yamamoto (1971) studied the adsorption of uranium (1 to
100 ppm U) dissolved in water onto volcanic ash, alluvial, and sandy soils.
The uranium was almost completely adsorbed on each of the soils. Uranium
desorption with salt solutions was extremely difficult especially for the
volcanic ash. Similar studies by Yamamoto et al. (1973) on the three soils
_2
using uranium (1 to 50 ppm U) and carbonated waters (4 to 109 ppm COg )
showed approximately 100% adsorption and less than 2% desorption.
Rubtsov (1972) determined the uranium content in forest podzolic mountain
soils and found a relatively high level of uranium in the <1 u particle size
fraction of the podzolic k* horizon. In general, for the soils studied 58%
of the total uranium was found in the <1 u soil fractions. Ritchie, Hawks,
and McHenry (1972) found the uranium content of sediments from the Little
Tall ahatchie River to increase with decreasing particle size.
Rancon (1973) studied the adsorption of uranium using four soils described
as follows: 1) .a river sediment containing a mixture of quartz, clay, calcite
and organic matter, 2) a river peat, 3) a sediment from Cadarache containing
a mixture of quartz, clay and calcite with no organic matter, and 4) a soil
developed on an altered schist from near laHague containing a mixture of
quartz and clay but no calcite or organic matter. The"first two soils were
equilibrated with their river waters containing 10 ppm uranium and the last
two soils were equilibrated with their respective groundwaters also containing
10 ppm uranium. The resulting uranium distribution coefficients are shown in
Table 3-99, which also includes the Kd values on pure quartz, calcite and
illite. The clay minerals in Soils 1, 3 and 4 were not identified or the
soils further characterized. Rancon also examined the effects of initial
uranium concentration on Kd values. Both the uranium concentration and solu-
tion pH changed as uranium was added to the solution. At 0.1 mg U/l, the pH
was 7.6, for example, and at 1.0 g U/l, the pH was 3.5. Because the pH
changes are a function of uranium concentration changes, the results are not
easily interpreted. In addition, the Kd concept, is invalid above the trace
uranium concentration KI.O mg U/l). Uranium adsorption data at 1 ppm versus
Kd also are presented. For Soil 4, three peaks were observed: Kd 300 ml/g
at about pH 5.5, Kd 2000 ml/g at pH 10 and Kd 270 ml/g at pH 12.
3-234
-------
TABLE 3-99. URANIUM Kd VALUES (RANCON, 1973)
Soil Kd, ml/g
1 - River Sediment 39
(clay, CaC03, OM)
2 - River Peat 33
3 - Sediment (clay, CaC03) 16
4 - Altered Schist (clay) 270
Quartz 0
Calcite 7
Illite 139
Rancon believed that the adsorption maxima represented by the three
peaks also represent electrokinetic potential maxima. Quartz was charac-
terized as inert, calcite was a poor uranium adsorber and clays were the
best adsorbers of uranium from solution. Acid, organic-rich soils show
much higher uranium Kd values than the alkaline peat (Soil 2) of this study.
Migration Results
Field Studies—
A study of granitic rock weathering by Harriss and Adams (1966) included
autoradiographs of fresh and weathered samples of several granitic rocks.
There was a definite increase in the density of concentrated radioactive mate-
rials with weathering. However, analyses indicated a small loss for uranium.
The increased density, therefore, must be due to losses of other materials
(alkalies and alkaline earths) during weathering. From a fresh granodiorite
containing 2.5 ppm uranium, the initial weathering resulted in losses of 60%
of the uranium. An acid leach of fresh rock also removed 60% of the uranium
indicating that most of the uranium was in acid soluble or interstitial mate-
rials. After an initial drop in concentration, the total uranium content of
the weathered rock increased by at least a factor of 4 in the uppermost
material.
Laboratory Studies—
Schulz (1965) suggested that uranium may be present in the soil as the
2+
divalent uranyl ion, UO^. , and will be mobile in soils if present as the uranyl
ion.
3-235
-------
Masuda and Yamamoto (1971) examined the desorption of uranium from
alluvial, sand and volcanic ash soils. The cation exchange capacities were
13.7, 7.7 and 33.0 meq/100 g, respectively, for the alluvial sand and vol-
canic ash soils. Strong salt solutions and distilled water were used as
leachates. Loads of more than 2000 yg U/g of soil were required before desorp-
tion by distilled water was 1% of the uranium on the soils. The uranium was
amended to the soils as uranyl nitrate before the desorption work. Desorption
of uranium was higher with 0.5M (NH4)2S04, 1.34M KC1 and 1.44M <2HP04 salt
solutions, but reached 50% removal only for the alluvial soil with a high
adsorbed uranium content in 1.44M KpHP04 solution. The volcanic ash soil did
not attain 5% uranium desorbed in any of the salt solutions.
The effects of carbonate ions on uranium desorption from soils was
examined by Yamamoto et al. (1973). They used an alluvial soil, a volcanic
ash soil and a sandy soil containing up to 500 ug U/g air-dried soil in their
desorption experiments. The desorption of uranium declined as a power func-
tion of the amount of uranium on the soil. Desorption results are shown in
Table 3-100. Ten grams of soil were magnetically stirred with 100 ml of potas-
sium carbonate solution for 30 min and stood overnight before filtering and
a fluorometric uranium analysis. As can be seen in Table 3-92, uranium desorp-
tion was very low in the presence of low to moderate environmental carbonate
concentrations.
TABLE 3-100. DESORPTION OF URANIUM FROM SOILS WITH DISTILLED WATER
AND CARBONATE SOLUTIONS (YAMAMOTO, 1973)
Soil
Alluvial
Sandy
Volcanic Ash
U Content,
yg/g
7.1
485.0
10.1
488.2
8.3
500.0
Distilled
Water, %
- 0.31
0.22
0.14
0.12
0.18
0.09
Carbonate
4.3mg
C03-2/l
0.62
0.25
0.27
0.13
0.57
0.10
Solution, %
43.4mg
CO,-2/T
— o — ! —
1.20
0.41
0.90
0.46
1.15
0.20
Dall'Aglio et al. (1974) discussed some of the geochemical processes
responsible for precipitating secondary uranium minerals. There are a series
3-236
-------
of processes capable of bringing about a very effective separation and con-
centration of uranium. The most important process is the attainment of high
+2
U09 activity in circulating waters because of low concentration or depletion
-2
of COg which is the most effective uranyl complexing agent. The micro-
organisms involved were soil microflora and bacteria of mine-waters and
granites. Some species were identified. Batch cultures were used to study
uranium insolubilization involving biodegradation of uranium complexing
organic compounds. The authors suggested that the redeposition of uranium by
bacteria may be the origin of some uranium deposits.
Uranium solubilization and insolubilization from granites by heterotrophic
bacteria were investigated by Magne et al . (1974). Microbial activity
increased the solubilization of uranium from 2 to 97 times by biosyntheses of
complexing and chelating compounds.
Grandstaff (1976) arrived at a rate expression for the effects of surface
area and uraninite composition, oxygen content of the solution, carbonate
content, organics content, pH and system temperature on uraninite dissolution.
R, the rate of uraninite (U02) dissolution = "d^an^ s 1020'25 (SS)(RF)"]
(1Q3. 38-10.8 NOC} (aZcQ) (D>Q>) (aj exp (_7045/T)day-l , where SS is the
specific surface area (cm2/g), RF is an organic retardation factor, NOC is. the
mole fraction of nonuranium cations in the uraninite, D.O. is the dissolved oxy-
gen content of the water (ppm), CC^ is the total dissolved carbonate and T is
the absolute temperature. The rate expression was used to predict UOg dis-
solution rates under varying conditions in the absence of organics in the
contacting water with good results. Dissolved inorganic species in artificial
seawater other than hydrogen ion activity (pH), total carbonate and dissolved .
oxygen did not appreciably affect the U02 dissolution rate. The organic
retardation factor must be determined experimentally for each organic-
containing solution by comparing the calculated U02 dissolution rate without
organics i-n the environment with the observed dissolution rate, or RF = R
calculated/R observed. Retardation values of up to 420 were obtained. The
magnitude of the affects of the several factors in the expression on the
oxidation rate of U(IV) to U(VI) could be of assistance in understanding
uranium mobility in soil and rock environments.
3-237
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Summary
The common oxidation states of uranium are U(III), U(IV), U(V) and U(VI)
(Udaltsova, 1963). However, in the geologic environment U(IV) and U(VI) are
the most important oxidation states. In oxidizing environments U(VI) compounds
such as KU02V04 (Garrels and Christ, 1965), U02NH4P04, Na2U04 and U02KP04
(Figure 3-26) are stable and can precipitate. U(IV) would precipitate as U02
in a reducing environment. U(VI) solution species govern uranium concentra-
tions and movement in oxidizing environments (Figure 3-27). Uranium retention
by soils and rocks in alkaline conditions is poor because the predominant
uranium species at pH > 6 in oxidizing environments (Figure 3-27) are either
neutral or negatively charged. An increase in CCL pressure in soil solutions
reduces uranium adsorption and can increase uranium concentration. The cation
exchange properties of soils could contribute to the adsorption of uranium in
2+
the neutral to acidic pH range due to the presence of UCL . Oxidation-
reduction conditions and pH would be important parameters of uranium mobiliza-
tion and immobilization.
The above theoretically based predictions are substantiated by experi-
mental results. Uranium has been reported to be solubilized and highly mobile
in .carbonate-containing waters (Brown and Keller, 1952; Naumov, 1961;
Ermolaev et al., 1965; Legin et al., 1966; Haglund, 1968, 1969). Soluble
uranium [U(VI}] can:
1. precipitate in the presence of phosphorus as evidenced by a direct cor-
relation of uranium and phosphate content in soils and rocks (Bell, 1960;
Sakanoue, 1960; Habashi, 1962; Kuznetsov et al., 1968; Menzel, 1968;
Mihalik, 1968), and
2. be adsorbed by the soil organic component and/or reduced to U(IV) fol-
lowed by precipitation (Breger et al., 1955; Kolodny, 1969; Kolodny and
Kaplan, 1970; Baturin, 1971; OalTAglio, 1971; Dorta and Rona, 1971;
Gabelman, 1971; Baturin and Kochenov, 1973; Mo et al., 1973).
An increase in uranium content with a decrease in soil or sediment par-
ticle size was reported by several workers (Kovalevskii, 1967; Mizuno and
Mochizuki, 1970; Ritchie et al., 1972).
3-238
-------
The uranyl ion can be adsorbed on clay minerals (Goldsztaub and Wey, 1955;
Kovalevskii, 1967; Rubtsov, 1972; Ritchie et al., 1972; Rancon, 1973) and
other adsorbent materials (Masuda and Yamamoto, 1971; Yamamoto et al., 1973),
but also is inclined to form complexes with anions, such as carbonate, that
are commonly found in the soil solution (Figure 3-27). Uranyl salts also
2+
have been shown to substitute for Ca during replacement of calcite by
2+ 2+
apatite (Ames, 1960), and Ca competes with UO^ for available sites during
ion exchange reactions on inorganics such as calcite (Rancon, 1973) resulting
in low uranium Kd values. However, uranyl ion adsorption on organic mate-
rials and humic substances is quite high (Szalay, 1954, 1957; Manskaya et al.,
1956; Rozhkova et al., 1959; Rancon, 1973) especially at acid pH values. The
most important parameters of uranium migration and retention are system Eh
and pH.
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Studies on the Weathering of Granitic Rocks. Am. J. Science. 264:146-173.
Horrath, E. 1960. Investigations of Uranium Adsorption to Peat in^Natural
Waters Containing U-Traces. Magyar Tudomanyos Akad. Atommag Kutato Intezete,
KOzlemenyek. 2:177-183 (in Hungarian).
Kolodny, Y. 1969. Studies in "Geochemistry of Uranium and Phosphorites.
University of California. Thesis.
Kolodny, Y. and I. R. Kaplan. 1970. Deposition of Uranium in the Sediment
and Interstitial Water of an Anoxic Fjord. CONF-700965-1.
3-240
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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. 0. V. Makew. Buryat. Khizhn.
Izd. Ulan-Ude, USSR.
Kuznetsov, Y. V., Z. N. Simonyak, A. P. Lisitsyn, and M. S. Frenklikh. 1968.
Thorium Isotopes (230Th, 232TH) in the Surface Layer of the Indian Ocean
Sediments. Geochem. Int. 5:169-177.
Latimer, W. M. 1952. The Oxidation States of the Elements and Their Potentials
in Aqueous Solutions. Second Edition. Prentice-Hall, Inc., Englewood Cliffs,
NO.
Legin, V. K., Y. V. Kuznetsov and K. F. Lazarev. 1966. Uranium Occurrence
in Marine Sediments. Geokhimiya No. 5:606-608 (in Russian).
Magne, R., J. R. Berthelin, and Y. Dommergues. 1974. Solubilization and
Insolubilization of Uranium from Granites by Heterotrophic Bacteria. IN;
Formation of Uranium Ore Deposits. IAEA-SM-183/18, p. 73.
Manskaya, S. M., G. V. Drozdora, and M. P. Yemel'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. 0. Radiat.
Res. 12:94-99.
Menzel, R. G. 1968. Uranium, Radium, and- Thorium Content in Phosphate Rocks
and Their Possible Radiation Ha.zard. J. Agr. Food Chem. 16:231-34.
Mihalik, P. 1968. Uranium Compounds in the Dominion Reefs, and Their Associa-
tion with Phosphorus Compounds. NIM-415.
Mizuno, A. and T. Mochizuki. 1970. Distribution of Uranium in the Sediments
from San'in Off-Shore, Southwestern Japan. Chishitsu Chosasho Geppo.
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Mo, T., A. D. Suttle, and W. M. Sackett. 1973. Uranium Concentrations in
Marine Sediments. Geochim. et Cosmochim. Acta. 37:35-51.
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3-241
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3-242
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ZIRCONIUM
Natural Soil and Rock Distributions
The zirconium distribution and content in igneous and sedimentary rocks
are shown in Table 3-101. It should be noted that much of the zirconium con-
tained in sedimentary rocks is present as resistate material such as zircon
(ZrSiO^), which is very slightly soluble in any aqueous solution. The average
content of worldwide soils is given by Vinogradov (1959) as 300 ppm zirconium.
TABLE 3-101. THE AVERAGE ZIRCONIUM CONTENT OF ROCKS
(TUREKIAN AND WEDEPOHL, 1961)
Rock Type Zr, ppm
Igneous Rocks
Ultramafic 45
Basaltic 140
Granitic, high Ca 140
Granitic, low Ca 175
Syenite 500
Sedimentary Rocks
Shale ' 160
Sandstone 220
Limestone-dolomite 19
Brief Chemistry
Eighteen isotopes and two isomers of zirconium are known. The naturally
nc
occurring isotopes are listed in Table 3-102. All are stable except Zr which
has a half-life listed as >3.6 x 10 years (Weast, 1976). For all practical
96
purposes, Zr also is a stable isotope.
95
The immediate concern in radioactive waste disposal is for Zr with a
relatively short half-life of 65 days. Also obtained in much smaller yields
QQ c
as a fission product is the longer-lived Zr (half-life 1.5 x 10 years)
95
which is of long-term interest (Schneider and Platt, 1974). Zr decays by
- 95 - 95
6 emission to Nb which also emits a 6 and decays to stable Mo.
3-243
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TABLE 102. NATURAL ISOTOPES OF ZIRCONIUM AND THEIR
ABUNDANCES (WEAST, 1976)
Isotope Abundance, %
90Zr 51.46
91Zr 11.23
Q2
y^Zr 17.11
94
*^Zr 17.40
96Zr 2.80
The most stable and common form of zirconium in rocks is as the simple
orthosilicate zircon, ZrSiO*. The only oxidation state of zirconium of any
+4
importance in aqueous solutions is Zr(IV), though usually not as Zr cations,
o
Due to the high charge and a small atomic radius (0.72 A in six-fold coordi-
o
nation, 0.84 A in octahedral coordination, Shannon and Prewitt, 1969), zir-
conium ions are readily hydrolyzed in aqueous solutions. When an acid solu-
tion is neutralized with a base such as NH.OH, hydrous zirconium oxide is
precipitated rather than zirconium hydroxide (Blumenthal, 1958). Zirconium
compounds, such as ZrCl., react violently with water, producing high acidity
solutions. A zirconyl salt (ZrOCl^'SHpO) is produced with liberation of
hydrogen ions. A solution 0.05M in ZrOCl9-8H«0, for example, has a pH of 1.
+4 4-n +
The equilibrium established is Zr + nH20 -»• Zr(OH)p + nH and depends on
the solution acidity.
According to Connick and McVey (1949), in 2.0 to 0.2M perchloric acid at
25°C, zirconium exists as Zr+ and Zr(OH) ions. At concentrations of per-
+2
chloric acid of less than 0.2M, the zirconium species was probably Zr(OH)2 .
At or above 2.0M HC104, zirconium exists exclusively as Zr ions (Larsen and
Wang, 1954). When zirconium hydroxides are acidified with HC1, oxo cations
+2
(ZrO ) first appear, and not until the solution is 2N in HC1 or higher, do
+4
appreciable amounts of Zr ions appear (Lister and McDonald, 1951).
Zirconium forms very stable complexes with halogens of the type M-ZrGg
where M is a univalent metal and G is a halogen. Zirconium also forms a
series of complexes with fluoride phosphate and sulfate ions and many of the
organic ligands. Electrolysis of a solution of ZrfSO^^HgO, for example,
produces hydrogen at the cathode and zirconium sulfate at the anode. Hence,
3-244
-------
the compound was H2 [ZrO(S04)2], with zirconium in the complex anion
[ZrO(SOA)9] . Since most of the waste solutions containing zirconium are
+4
near-neutral solutions, no Zr ion should exist in solution. There are only
charged to neutral polymers and complexes of zirconium existing in the pH
range of most radioactive waste solutions.
Solid Phase and Solution Equilibria
The zirconium minerals in increasing order of stability are: Zr(OH)4,
Zr02, and ZrSiO^, as shown in Figure 3-28. The thermodynamic data for Zr(OH)4,
Zr02 and ZrSiCL were obtained from Sillen and Martell (1964), Baes and Mesmer
(1976) and Wagman et al. (1971), respectively. The most stable mineral
throughout the pH range is zircon (ZrSiO.). This is consistent with the
empirical observations of many geologists and soil scientists who find that
zircon (ZrSiO*) is ubiquitous in soils and is very stable even compared to
other common soil minerals.
-8
-10
-12
-14
S1
-16
-18
-20
-22
6
PH
Figure 3-28. The stability of zirconium solids at 25°C and in
equilibrium with soil silica (pHUSiCty = 3.1).
3-245
-------
The activity of various solution species with the change in pH and in
equilibrium with zircon (ZrSiO^) and soil Si02 is given in Figure 3-29. The
thermodynamic data for all the hydrolysis species were selected from Baes and
Mesmer (1976). Other solution species data were selected from Si 11 en and
MarteH (1964). Zirconium forms complexes with the common soil am'ons (F~,
- 2- -
Cl", SOA , NO,"). The significance of complexation of zirconium with these
* ** 2-
anions decreases in the following order: F , SO, , Cl , NO, . Although
3+
zirconium forms complexes with common soil anions, ZrF complex would be the
only one which could contribute significantly in acidic environments (pK 3.5)
to the total zirconium in solution. It can be seen from the diagram that the
uncomplexed zirconium (Zr ) would exist in significant amounts only in very
acidic solutions (pH <1). The zirconium ions are hydrolyzed in solutions with
pH values of greater than 1.0. The most dominant zirconium solution species
3+ °
with the conditions assumed in Figure 3-29 are ZrF , Zr(OH)4 , and Zr(OH)5"
in pH ranges of <3.5, 3.5-6.25, and >6.25, respectively.
s
-2
-16
-20
-24
-23
ZrtOHI.
Z 3 4 5 6
PH • '
D
Figure 3-29. The activity of various zirconium species in equilibrium with
zircon [ZrSiO^s)] and soil solution silica (pfyS^ = 3.1) at
pCT = pS042~ = 2.5, pN03~ = 3.0 and pF" = 4.5.
If a waste solution containing zirconium is assumed to be in equilibrium
with zircon and the conditions assumed in Figure 3-29, then the total activity
-14
of zirconium in solution would be lower than 10 over a pH range of 3.5 to 8.
3-246
-------
The Zr ion becomes increasingly hydrolyzed above pH 1 (Figure 3-29).
4-f
The Zr ' content of the soil solution falls rapidly, as does the content of
fluoride complex ZrF , in favor of hydrolyzed species at higher pH values.
Only at relatively low pH values of less than 4 do positively charged species
of zirconium predominate. Because most soils and rocks buffer waters contact-
ing them to pH 5 or higher, ion exchange probably plays a relatively minor
role in zirconium retention by soils and rocks. Because of the formation of
negatively charged complexes in solutions of pH >6.5 (Figure 3-29), zirconium
adsorption would be expected to decrease with an increase in pH above 6.5.
Experimental Adsorption Results
Rhodes (1957) determined the effect of pH on the distribution coefficient
of zirconium-niobium on a Hanford soil. The soil was. made up of 2% CaC03 with
>2 mm diameter material, 6%; 2 to 0.2 mm diameter material, 67%; 0.02 to
0.2 mm diameter material, 19%; 0.02 to 0.002 mm diameter material, 6%; and
<0.002 mm diameter material, 2%. The cation exchange capacity was 5 meq/100 g
and the water paste pH was 8.6. The bulk of the soil consisted of quartz,
feldspars, augite, olivine and mica. The clay fraction contained mainly mont-
morillonite. The Kd values for zirconium as a function of pH are given in
Table 3-103.
TABLE 3-103. ZIRCONIUM-NIOBIUM Kd VALUES AS A FUNCTION OF pH
(RHODES, 1957)
fitL.
0
i
2.6
4.1
6.4
7.7
Kd, ml/g
82
1028
1340
>1980
>1980
>1980
PH
8,3
9.6
10.2
11.0
12.2
' 14.0
Kd, ml/q
282
90
104
180
>1980
>1980
Rhodes considered that the dips in zirconium Kd for pH 8 to 12 were due to
a charge alteration of the zirconium polymer or colloid within this pH range.
Dlouhy (1967) gave Kd values for zirconium on Casaccia soil and Casaccia
tuff from CNEN, Italy. The soil and tuff were not characterized. The range of
zirconium Kd values given for the soil was 130 to 150 ml/g and for the tuff,
260 to 350 ml/g.
3-247
-------
Benson (1960) reported a visible, white precipitate in systems containing
-2 95
3 x 10 M zirconium. The uptake of trace Zr was not appreciably different
from that noted in the presence of carrier indicating that precipitation also
occurred at the trace concentration. Zirconium was not appreciably affected
by the presence of other cations unless the pH was below 2. Benson also
reported that carbonate, oxalate and citrate ions inhibited zirconium uptake
on soil over a wide pH range, probably because of the formation of complex
anionic species.
Sorathesn et al. (1960) stated that 95Zr behaved like a colloid at pH 6
to 9. They reported the zirconium Kd on illite at 7 days equilibration time
to be 24,470 ml/g.
-3 -2
Prout (1959) reported that tests made at pH 4 with 10 M C^O^ showed
adsorption of zirconium-niobium on soil to be inhibited, but at pH 8, the
oxalate did not seriously affect the removal of zirconium-niobium. Schulz
(1965) reported that zirconium and niobium were two of the elements that were
adsorbed by soil in a very immobile form.
Arnold and Grouse (1971).determined the adsorption of fission products on
copper ore produced by nuclear fracturing. The copper leach' solution was
pH 1.2 to 1.8 with FLSO, and at 85°C. The copper ore was in 2.5 kg, 4-1-n.
columns with 2000 g of radioactive rubble from a nuclear shot added between
500 g layers of minus 1/2-in. copper ore. Sulfuric acid was added during the
run to assure that pH values stayed below 2. Almost none of the Zr- Mb
was found in the column effluent because it was strongly adsorbed by the ore.
95 95
When a second run'was made with the fission products, including Zr- Mb
added in soluble form, the adsorption results were the same. Zirconium Kd
values of 50 to 60 ml/g were obtained on batch equilibria from pH 1.5 to
pH 2.8.
Migration Results
Field Studies—
Spitsyn et al. (1958) used an alkaline solution (4 to 8 g NaOH/1, 200 g
NaN03/l) and an acidic solution [6 to 8 g HN03/1, 200 g A1(N03)3/1] in labora-
tory and field studies of radionuclide migration rates. Zirconium was
3-248
-------
adsorbed and retained by a sand in field tests near the point of injection of
the waste solutions for both the acidic and alkaline solutions. Little move-
ment was indicated even under acidic conditions.
Magno et al. (1970) investigated the radionuclides from the Nuclear Fuel
Services plant in western New York that were migrating through the plant
95
effluent lagooning system. They estimated that greater than 90% of the Zr
discharged from the plant was deposited by sedimentation in the lagoon system.
Brookins (1976) reported that fission product zirconium was retained in
the shale very near to areas of generation in the natural reactor at Oklo.
Laboratory Studies-
Fallout material on soil from an underground nuclear explosion was used
by Essington et al. (1965) to determine the leaching effects of water, HC1 and
chelating agents. The soil was suspended in 25 ml of solution and equilibrated
-4
for up to 106 days. The 10 molar chelating agents in solution included
sodium-diethylenetriaminepentaacetate, sodium-cyclohexane-1,2-diaminetetra-
acetate, sodium-ethylenediamine di(o-hydroxyphenylacetate) and 0.1M HC1.
Zr- Nb was identified in the original soil. Chelating agents, HC1 and
water equilibrations did not appreciably increase Zr- Nb solubility. In
column studies with Yolo soil (pH 7.9), Sorrento soil (pH 7.8) and Hanford
soil (pH 6.6), water and calcium-diethylenetriaminepentaacetate (10" M) leaches
of 7 days produced no detectable Zr- Nb in the effluent solution. The domi-
nant clay minerals in the Hanford, Sorrento and Yolo soils were illite, kao-
linite and montmorillonite.
According to Schulz (1965), the zirconium in soils is either very strongly
adsorbed by the clay particles or is present as an insoluble oxide or hydroxide.
In either case, the zirconium in the soil is immobile.
Eichholz et al. (1967) investigated the partitioning of dissolved radio-
nuclides between suspended sediment particles and aqueous solutions. Several
natural water samples were characterized and used in the adsorption studies as
suspended solids sources (see Table 3-30). A fission product mixture was
added to the natural water samples, the system equilibrated and the water
95 95
recovered and filtered. The adsorption of Zr and daughter Nb are shown
3-249
-------
in Table 3-104. A large portion of the zirconium was associated with the par-
ticulate matter in a highly concentrated form.
TABLE 3-104. 'ADSORPTION OF 95Zr AMD 95Nb ON SUSPENDED SOLIDS
IN NATURAL WATERS (EICHHOLZ ET AL., 1967)
95 95
Suspended Zr Adsorption, Nb Adsorption,
Source Solids, mg/l, % %
Colorado River 229 61.1 33.3
Camp McCoy 12 50.0 4.7
Bayou Anacoco 24 80.5. 23.8
Lodgepole Creek 965 81.3
Chattahoochee River 131 88.4 18.9
Billy's Lake 8 53.9 8.9
Adsorption does not appear to be a direct function of suspended solids
content but has to include a consideration of dissolved solids content, pH and
solids composition before the adsorption results begin to make any sense. The
95
high Zr adsorption on Chattahoochee River suspended solids was due to the
low dissolved solids content (31 ppm), which suggests that a positively charged
zirconium species was involved.
95 95
Harrison (1969), using nuclear debris contaminated with Zr- Nb as a
source, approached equilibrium from the direction of starting with the radio-
nuclide on the solid. The leaching produced Zr- Nb Kd values of 5,200 to
17,000 ml/g in distilled water as a function of particle size range (<4000 to
<62 urn), 790 to 2700 ml/g in synthetic seawater for the same particle size
range and 270 to 1200 ml/g in 1M ammonium acetate for that particle size range.
An implication of this work is that cation exchange is involved because the
95 95
Zr- Nb Kd values are sensitive to competing ion concentrations.
Summary
Trace concentrations of zirconium are strongly adsorbed on soils at pH 1
or higher with a dip in Kd values .between pH 8 to 12 (Rhodes, 1957), probably
due to the prevalence of uncharged or anionic zirconium species over this pH
range (Figure 2-14). There is some evidence for affects caused by competing
cations at near-neutral pH, suggesting at least partial removal by an ion
3-250
-------
exchange mechanism of polymerized or colloidal species (Eichholz et a!., 1967;
Harrison, 1969). Zirconium is well-adsorbed by the soil over the normal pH
range of 4 to 8 (Spitsyn et al., 1958; Schulz, 1965; Magno et al., 1970).
Carbonate, oxalate and citrate complexes of zirconium have been reported to
migrate rapidly through soil columns (Bensen, 1960).
References
Arnold, W. D. and D. J. Grouse. 1971. Radioactive Contamination of Copper
Recovered from Ore Fractured with Nuclear Explosions. ORNL-4677.
Baes, C. F., Jr. and R. F. Mesmer. 1976. The Hydrolysis of Cations. John
Wiley and Sons, New York.
Benson, D. W. 1960. Review of Soil Chemistry Research at Hanford. HW-67201.
Blumenthal, W. B. 1958. The Chemical Behavior of Zirconium. Van Nostrand
Company, Inc.
Brookins, D. G. 1976. Shale as a Repository for Radioactive Waste: The
Evidence from Oklo. Environmental Geology. 1:255-259.
Connick, R. E. and W. H. McVey. 1949. The Aqueous Chemistry of Zirconium.
J. Am. Chem. Soc. 71:3182-3192.
Dlouhy, Z. 1967. Movement of Radionuclides in the Aerated Zone.- IN:
Disposal of Radioactive Wastes into the Ground. IAEA-SM-93/18.
Eichholz, G. G., T. F. Craft, and A. N. Galli. 1967. Trace Element Fractiona-
tion by Suspended Matter in Water. Geochim. et Cosmochim. Acta. 31:737-745.
Essington, E. H., H. Nishita and A. J. Steen. 1965. Release and Movement of
Radionuclides in Soils Contaminated with Fallout Material from an Underground
Thermonuclear Detonation. Health Physics. 11:689-698.
Harrison, F. L. 1969. Radioactive Debris from Underground Nuclear Explosions:
1. Physical and Chemical Characteristics, 2. Biological Availability of the
Radionuclides to Aquatic Animals. UCRL-5Q596.
Magno, P., T. Reavy, and J. Apidianakis. 1970. Liquid Waste Effluents from
a Nuclear Fuel Reprocessing Plant. BRH-NERHL-70r2.
Larsen, E. M. and P. Wang. 1954. The Ion Exchange of Zirconium and Hafnium
in Perchloric Acid with Amberlite IR-120. J. Am. Chem. Soc. 76:6223-6229.
Lister, B. and L. McDonald, 1951. Some Aspects of the Solution Chemistry of
Zirconium. AERE-C/R-801.
Prout, W. E. 1959. Adsorption of Fission Products by Savannah River Plant
Soil. OP-394.
3-251
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Rhodes, D. W.' 1957. The Effect of pH on the Uptake of Radioactive Isotopes
from Solution by a Soil. Soil Science Society America, Proceedings.
21:389-392.
Schneider, K. J. and A. M. Platt. 1974. High-level Radioactive Waste
Management Alternatives. BNWL-1900, Vol. !.
Schulz, R. K. 1965. Soil Chemistry of Radionuclides. Health Physics.
11:1317-1324.
Shannon, R. D. and C. T. Prewitt. 1969. Effective Ionic Radii in Oxides and
Fluorides. Acta Cryst. B. 25:925.
Sillen, L. G. and A. E. Martell. 1964. Stability Constants of Metal-Ion
Complexes. Special Publication No. 17. The Chemical Society, London.
Sorathesn, A., G. Bruscia, T. Tamura, and E. G. Struxness. 1960. Mineral
and Sediment Affinity for Radionuclides. CF-60-6-93.
Spitsyn, V. I., V. D. Balukova, A. F. Naumova, V. V. Gromov, F. M. Spiridonov,
E. M. Vetrov, and G. I. Gravfov. 1958. A Study of the Migration of Radio-
elements in Soils. IN: Proc. Second Annual U.N. Conf. Peaceful Uses of
Atomic Energy. 18:439-448.
Turekian, K. K. and K. H. Wedepohl. 1961. Distribution of the Elements in
Some Major Units of the Earth's Crust. Bull. Geol. Soc. Am. 72:175-192.
•
Vinogradav, A. P. '1959. The Geochemistry of Rare and Dispersed Chemical
Elements in Soils. Consultants Bureau.
Wagman, D. D., W. H. Evans, V. B. Parker, I. Halow, S. M. Bailey, R. H. Schumm,
and K. L. Churney. 1971. Selected Values of Chemical Thermodynamic Properties.
Tables for Elements 54 Through 61 in the Standard Order of Arrangement. NBS
Technical Note 270-5.
Weast, R. C. 1976. Handbook of Chemistry and Physics. The Chemical Rubber
Co.
3-252
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SECTION 4
CONCLUSIONS AND EVALUATION
The available data on interaction of various elements with the geologic
media were collected and discussed in Section 3. At the end of discussions on
each element a summary is presented. The information presented in Section 3
will not be repeated here. This section is concerned with general conclusions
based upon Section 3. For more details the reader is directed to the previous
sections.
The solution ions exist in dynamic equilibrium with soils, sediments, and
rocks. Many factors that are inherent to the solid matrix and the solution
can influence the concentration of the element and its species in solution.
These factors include pH, Eh, CEC, type and amount of soil minerals, solid
phases of the element, complexing ligands, and competing ions. In order to
determine the distribution of the element between solution and solid and its
reactions with the solid matrix, the identity and relative magnitude of the
various factors are required. Such information would help predict the solu-
tion concentrations of the element and help extrapolate the results to other
situations. However, the data presented in Section 3 indicate that there is
a general lack of systematic evaluation of various factors that determine
element-solid matrix interactions, and no information at present is available
to determine the magnitude of the various factors. At best, the available
data suggest trends of the influence of some of the factors that control solu-
tion concentrations and interaction with solid matrices. This type of informa-
tion would not be very useful for precisely predicting the general fate of
radionuclides in the environment. Additional research required for general
understanding of the behavior of radionuclides is outlined at the end of this
section.
4-1
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FACTORS INFLUENCING ADSORPTION OF RADIONUCLIDES ON GEOLOGIC MEDIA
Some of the factors that have been reported to influence the adsorption
of elements by geologic media are listed for selected elements in Table 4-1.
An "X" indicates that a given factor was reported to influence adsorption of
that element (present as a radionuclide). The results reported in Table 4-1
are qualitative only. In most cases, the authors did not definitively show
that a specific factor was operative, but hypothesized that it was. Since
these data cannot be used for quantitative predictions, they are mainly pre-
sented for use as guidelines in future experiments. The probable principal
adsorption mechanisms were deduced from an examination of the factors reported
for each element. For example, colloid formation is a good indication of pre-
cipitation, as pH also may be. An adsorption sensitivity to soil cation
exchange capacity, competing ions and system pH usually indicates an ion
exchange adsorption mechanism. There is, admittedly, a certain amount of
value judgment that went into the summation process because in some cases the
literatur-e was ambiguous and vague concerning adsorption-influencing factors,
and in other cases, the experimental work has yet to be done. Neptunium and
technetium literature data could not be used to determine a probable adsorp-
tion mechanism.. Adsorption of'iodine by organic matter is listed as the
adsorption mechanism, but various authors gave conflicting results. How-
ever, the majority of the workers reported organic matter to be the principal
adsorption media for iodine.
In addition to concentration, the nature of the solution species would
have, a tremendous effect on the radionuclide interaction with the geologic
media and on the mobility of the radionuclides. Very few data are available
in the literature delineating species that have been identified in natural
environments. Thermodynamic data were used (Section 3) to develop solution
species diagrams to predict the predominant solution species. The limitations
and drawbacks of this approach are discussed in Section 2. The predicted
nature of the predominant solution species of various elements are reported in
•*
Table 4-2. The environmental conditions assumed for Table 4-2 data include a
pH range of 4 to 9, p02 range of 0.68 to 80 atmospheres, pC02 of 1.52 to 3.52
atmospheres, pd" = pN03" = pSOJ£ = 3.0, pF" 4.5, and pH^O^ 5.0. It should
be pointed out that a change in environmental conditions, in addition to the
4-2
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TABLE 4-1. FACTORS REPORTED TO EFFECT ADSORPTION OF RADIOELEMENTS OVER THE
pH RANGE OF 4 TO 9
Factors
Soil
Element pll Eh CEClO
Am
Sb
Ce
Cs
Co
Cm
Eu
I
Np
Pu
Pm
Ra
Ru
Sr
Tc
Th
3H
U
Zr
X
X
X
X
X
X
XX X
X
X
X
X X
X X
XX X
X
Competing Selectively Inorganic
. Ions Adsorbed On Llgands
X
Iron Oxides
Zeolites,
Micas
Illlte,
Iron Oxide
X
OM
X .
X Zeolites,
Bar He
OM
X • Calclte,
Zeolites
OM
H20
OM
X
X
X
X
X
X
X
X
X
X
X
X
X
Complex Ions
Organic
Constituents
X
X
X
X
X
X
X
X
X
X
X
X
Colloid Probable Adsorp-
Formatlon tlon Mechanlsmsl?)
IE
X PPT
X TE. PPT
IE
X IE, PPT
X PPT
IE, PPT
OM
X UNK
X IE, PPT
IE, PPT
IE
PPT
IE
UNK
X IE, PPT
NONE
PPT, IE
X PPT
1. CEC = Cation Exchange Capacity
2. IE = Ion exchange, OM = Organic matter adsorption, PPM =? Precipitation, UNK = Unknown.
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TABLE 4-2. PREDOMINANT SOLUTION SPECIES OF ELEMENTS
(a)
El ements
Am
Sb
Ce
Cs
Co
Cm
Eu
I
Np
Pu
Pm
Ra
Ru
Sr
Tc
Th
U
Zr
Little Affected by In An In A
Oxidation-Reduction Oxidizing Environment Reducing Environment
Am3*, AmS04*. AM(OH)2*
HSb02°, Sb(OH)3°, SbOF\ SbO*
Sb(OH)4-
Ce3*, CeS04*
Cs*
Co2*, Co(OH)2*
Cm3*, CmOH2*, Cm(OH)2*
Eu3*, EuS04*. Eu2P20?2*
r, io3" r
Np02*. Np02HP04', NpOH3*, Np4*
- Np02HC03
Pu022*, Pu02(C03)(OH)|', PuOH2*,.Pu3*
Pn,3* .
Ra2*
Ru(OH)2*, Ru04", Ru042" Ru04"
Sr2*
Tc04' Tc2*
ThF3*, Th(OH)3*
H*. 3H-0-H
UO|*. U02F*. U02(OH2)8, U02,*, UOH3*, U02*.
UO,(CO,)i~ UO,(CO,),4"
6 J J ' fc O J
Zr(OH)4°, Zr(OH)s-,
ZrF3*
(a) In a pH 4 to 9, pOa 0.68 to 80, pCOe 1.52 to 3.52, pCl" = pNOs" =
pS042" =3.0, pF~ 4.5 and pH2P04" 5.0 environment without organic
ligands.
4-4
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ligands not considered in the above assumptions, may change the nature of the
solution species. Depending upon the given chemical environment, several
species of an element can exist in solution. Cesium, cobalt, promethium,
radium, strontium, and tritium would be expected to exist predominantly as
uncomplexed solution species as Cs , Co , Pm , Ra , Sr and H+. Certain
elements such as plutonium, neptunium and uranium can exist in acidic solu-
+ 2+
tions in an oxidizing environment as uncomplexed solution ions (PuO~ , PuCL ,
+ 2+
NpO-, UOg ). Some of the elements such as curium, thorium and zirconium
hydrolyze very readily even in acidic environments, so they mainly exist as
hydrolyzed species [CmOH2*, Cm(OH)t, Th(OH)t, Zr(OH)!, Zr(OH)"]. Some of the
n n <- J *T D
common soil ligands (COg » S0^~) form strong complexes with americium, curium,
europium, neptunium, plutonium and uranium so that these ligands would
influence the concentration of these elements in solutions. Oxidation-
reduction conditions would predominantly influence the nature of the solution
species of antimony, neptunium, plutonium and uranium.
Based upon the nature of the predominant solution species, qualitative
predictions regarding the adsorption and movement of various elements can be
made. Soils and sediments mainly show cation exchange.capacities (since
materials carry a large net negative charge) and to a limited extent,
anion exchange capacities. Thus, most cations migrate through the soil or
rock column at speeds slower than the groundwater. Relative to each other,
the trivalent cations generally move the slowest, the divalent cations at
intermediate velocities and the monovalent cations most rapidly. Tritium is
unique in that it readily substitutes for hydrogen in water and migrates,
therefore, at the same velocity as water. Complicating factors include a
higher selectivity of a soil or rock component for a given cation or a more
successful hydrogen ion competition with one cation relative to another. The
relative mobilities of strontium and cesium at low pH is a good example of
the latter effect. Strontium is much more mobile at low pH than is cesium.
The simple anions tend to migrate through soils and rocks with little
reaction because usually a pH of less than 4 is required to activate a signifi-
cant soil anion exchange capacity. However, both I and Tc which exist pre-
dominantly as anions (I", TcOT) undergo other reactions with organic ligands
that can greatly slow their migration.
4-5
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The migration and retention of inorganic complex species (mononuclear and
polynuclear) would also be dependent upon the charge and size of the species.
Positively charged species would move slower than the negatively charged
species. The highly charged metal ions tend to polymerize or form colloidal
size, charged precipitates. The trivalent metal ions (Am, Cm, Sb, Ru) and
quadrivalent or higher charged metal ions (Pu, Th, U, Zr) are known to form
polymers. Initially, at lower pH, the polymers are positively charged and at
pH 8 become increasingly negatively charged. The result is good adsorption
and very slow movement at acidic to neutral pH, and greatly lessened adsorp-
tion and rapid migration from pH 8 upward.
The behavior of organic complexed species of elements is difficult to
predict because of the lack of knowledge regarding the exact nature of the
organic ligands, a wide variation in amounts and types of organic ligands,
and the size and solubility of these organics. These organic materials can
either retain the element or complex it in a form that migrates readily. The
source of these organic ligands can be the radioactive waste solutions con-
taining the synthetic organic ligands such as TBP and/or the organic ligands
produced by soil flora.
The metal ions that take part in replacement reactions depend upon the
size and charge of the species involved and the ability of the final mineral
product to accept that species into its growing structure. Apatite, for
example, accepts into its structure a wide range of metal ions such as
strontium, radium, uranium and probably americium and cerium. Cobalt replaces
calcium in calcite, and radium and strontium replace barium in barite. These
metal ions are buried in the replacement product crystal structure and are
effectively retained..unless a changing chemical environment causes instability
and dissolution of the crystal.
ADDITIONAL RESEARCH NEEDS
The areas that require more work to better understand radionuclide inter-
actions with soil and rock media include:
1. Determination of radionuclide adsorption mechanisms.
4-6
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Additional work is required in the area of radionuclide adsorption
mechanisms (ion exchange, precipitation, coprecipitation). There are many Kd
values that include several soil adsorption mechanisms and elemental species
operating simultaneously. Definitive experimental data that would allow
deduction of probable soil or rock adsorption mechanisms for different elements
are lacking in the literature. In addition, the environment in which a given
adsorption mechanism predominates should be delineated as well.
2. Factors that influence radionuclide-geologic media interactions.
The influence of different factors (such as pH, Eh, complexing ligands,
competing ions, CEC, type and amount of soil minerals, solid phases of ele-
ment) on the magnitude and extent of adsorption of radionuclides by the
geologic media need to be evaluated.
3. An experimental .consideration of the effects of redox potential on radio-
nuclide migration.
At depth (below the soil humic zone), where geologic storage has been
proposed, the environment is considerably different from that contributing to
radionuclide migration on the land surface. There has been minimal examina-
tion of the effects on radionuclide migration of the relatively low oxygen
partial pressures that will be encountered as a result of geologic waste
storage at depth. Many of the igneous rocks comprising a proposed waste
repository contain phases such as magnetite, amphiboles or pyroxenes with
ferrous iron as a constituent that tends to "buffer" the system at a rela-
tively low oxygen partial pressure until it is all oxidized to ferric iron.
The equilibrium partial pressure of oxygen for the reaction, 2Fe304 (magnetite)
+ 1/2 02 g » 3Fe203 (hematite) at 25°C and one atmosphere pressure is 10"68'2,
a reducing environment. Of the 19 elements studied a large change in redox
potential would be expected to affect the oxidation states of antimony,
neptunium, plutonium, ruthenium, technetium and uranium, henoa their solution
species and ultimately their migration rate through the surrounding rock.
-------
4. Characterization and influence of organic ligands on radionuclide
migration.
The existing thermodynamic data on species are incomplete in many cases,
and of dubious quality in other cases. The thermodynamic data should be con-
firmed by experimental evidence on radionuclide adsorption and migration.
Thermodynamic data on the radionuclide complexes with natural soil and water
organic components are essentially nonexistent. Note in Table 4-1 that radio-
nuclide reactions with organic material were reported for 12 of the 19 radio-
nuclides reviewed. Hence, what may prove to be a most important influence on
radionuclide adsorption and migration is one of the least understood. The
sources of organic materials in the environment capable of forming complexed
species with radionuclides include the many industrial wastes from various
manufacturing processes disposed to the atmosphere and surface waters, the
organics included in municipal sewage treatment effluents also disposed to
surface waters, the various pesticides and herbicides used during farming
operations contained in the atmosphere and irrigation runoff, humic substances
and decaying plant and animal remains already present in the soil and rocks
along with the associated living soil microflora. With such varied sources
and types of organic ligands, a large quantity of work remains to be done in
even identifying the metal ligand species involved in radionuclide migration
on the land surface.
5. Thermodynamic data determinations for several solids and solution species.
Thermodynamic data are badly needed for several solids and complex solu-
tion species of different elements. For example, the ruthenium complexes
RuNO(N02)2OH(H20)2 and RuNON02(N03)2(H20)2 were reported as very stable and
commonly found in nuclear waste solutions by several authors. Yet no firm
thermodynamic data for these two complexes were located in the literature.
When the data are determined, they may drastically change the order and
speciation for ruthenium solution species reported in Section 3.
4-8
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6. The generation of Kd data that can be used in a comparative mode.
As mentioned earlier, an attempt was made to evaluate and select the
best Kd values from the literature in writing Section 3. However, the experi-
mental conditions, including soil, rock and solution characterization, were
only sketchily given, or not given at all in most cases. Hence, the Kd values
in Section 3 are the best available, but evaluation of these data by any
objective methodology would result in discarding a large number of them. The
requirements for determination of a Kd value that could be used in a compara-
tive mode were listed in Section 2. A program to determine such Kd values is
now underway by the Office of Waste Isolation of the Department of Energy.
7. Maintenance of a bank of such Kd data.
A bank of the Kd data should be maintained and kept current. The data
could be computer-stored and made available to all personnel concerned with
waste management upon request. The present system has resulted in using
unevaluated Kd values that are often not applicable to the case in point.
If Kd values for biotite, quartz and orthoclase are known, a computed
resultant Kd value for a granite, for example, should be possible if com-.
parable surface areas are equated and if the biotite, quartz, and orthoclase
content of the granite are known. If the granite is composed of 25% quartz,
65% orthoclase and 10% biotite, and surface areas of minerals and rocks are
similar, then (0.25) (quartz Kd) + (0.65) (orthoclase Kd) + (0.10) (biotite Kd)
should equal the granite Kd value. There is no evidence in the literature
that the above hypothesis has ever been tested. Until Kd determinations are
placed on a comparable experimental basis, there would be little point in
attempting to test and apply it.
8. Determination of radionuclide adsorption and desorption reaction kinetics
with soils and rocks.
Several authors that have equilibrated solutions containing radionuclides
with soils or rocks for several weeks have reported that a considerable period
of time was required to attain equilibrium. Kinetic studies 1) could indicate
radionuclide adsorption mechanisms, 2) could help in separating adsorption
rates due to different mechanisms, and 3) would help in increasing the confi-
dence in extrapolating the predicted results into the future.
4-9
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ACKNOWLEDGMENTS
The authors would like to acknowledge the following persons who particv
pated in the review of this report:
Mr. H. J. Morton
Mr. W. K. Summers
Dr. K. A. Kraus
Dr. D. Isherwood
Mr. G. Grisak
Mr. G. L. Meyer
Dr. D. C. Hoffman •
Mr. J. P. Hamric
Other support was given by the EPA Office of Research and Development and
the Office of Energy, Minerals and Industry.
The authors also wish to thank the following Battelle staff members
who assisted in the report preparation: S. I. Thoreson, D. J. Kennedy,
M. V. Heid, and S. R. Gano.
Ack-1
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