defining abiotic and biotic pathways of metal...

DEFINING ABIOTIC AND BIOTIC PATHWAYS

OF METAL REDOX TRANSFORMATION IN NATURAL SEDIMENTS

A DISSERTATION

SUBMITTED TO THE DEPARTMENT OF GEOLOGICAL AND

ENVIRONMENTAL SCIENCES

AND THE COMMITTEE ON GRADUATE STUDIES

OF STANFORD UNIVERSITY

IN PARTIAL FULFILLMENT OF THE REQUIREMENTS

FOR THE DEGREE OF

DOCTOR OF PHILOSOPHY

Matthew Allen Ginder-Vogel

September 2006

ii

© Copyright by Matthew Allen Ginder-Vogel

All Rights Reserved

iii

I certify that I have read this dissertation and that, in my opinion, it is fully adequate in

scope and quality as a dissertation for the degree of Doctor of Philosophy.

(Scott Fendorf) Principal Adviser



(Gordon E. Brown, Jr.)



(Craig S. Criddle)

Approved for the University Committee on Graduate Studies.

iv

ABSTRACT

Uranium and chromium contamination is widespread and extensive, which makes

understanding processes that affect the environmental mobility of these elements

imperative. The hexavalent oxidation state of these two elements is most susceptible

to environmental transport and biological uptake; however, in their reduced states,

Cr(III) and U(IV), their solubility, and hence their mobility, are limited. Since the

mobility of these contaminants is largely determined by each contaminant’s valence, it

is therefore vital to identify the environmental processes affecting the oxidation state.

Once identified, these processes will further our understanding of contaminant natural

attenuation and allow predictive modeling of their fate and transport; additionally,

they may be exploited for in situ contaminant immobilization.

The research presented in Chapters 2 through 4 examines the role of abiotic

(chemical) and biotic (bacterial) controls on uranium and chromium mobility in

environmental systems. The oxidation of biologically precipitated uraninite by Fe(III)

(hydr)oxides may limit uranium sequestration under mildly reducing conditions. This

oxidation pathway is examined in detail in the second and third chapters. The impact

of geochemical conditions on the energetic favorability and extent of uraninite

oxidation is investigated using a combination of thermodynamic calculations and

static batch reactions in Chapter 2. This reaction is further investigated in Chapter 3,

where I demonstrate that the rate of oxidation is controlled by the uraninite dissolution

rate. Mechanisms responsible for the reduction and retention of chromate within arid

sediments are identified in Chapter 4. The research presented in this thesis combines a

fundamental approach to investigating the molecular-scale interaction of heavy metal

v

and radionuclide contaminants with field-based observations in order to identify

the operative biogeochemical pathways controlling contaminant mobility.

vi

ACKNOWLEDGMENTS

I am grateful for the excellent mentorship I received from my advisor, Scott Fendorf.

I thank my committee members, Scott Fendorf, Chris Francis, Gordon E. Brown, Jr.,

and Craig Criddle, for their time and encouragement. I am also appreciative of the

time Kevin Arrigo gave me to chair my committee. The members of the Soil and

Environmental Biogeochemistry group from 2001-2006, Guangchao Li, Matt Lappe,

Thomas Borch, Chris Oze, Colleen Hansel, Matthew Polizzotto, Ben Kocar, Yoko

Masue, Brandy Stewart, Kate Tufano, Celine Pallud, and Karen Murray, were

wonderful colleagues upon whom I could rely for advice and collaboration. I also

thank the members of Gordon Brown’s Surface and Aqueous Geochemistry Group for

their collegiality. My time at Stanford was made much more enjoyable by a large

group of friends including Tricia Fiore, Jeff Allwardt, Tasha Reddy, Steve Loheide,

Mike Beman, Juyoung Ha, Bob Jones, and Gary Ernst. The support of Sam Webb,

John Bargar, Joe Rogers, Steve Sutton, and Matthew Newville at the Stanford

Synchrotron Radiation Laboratory and Advanced Photon Source is greatly

appreciated. I am thankful to my family members and friends for their consistent

support, and would like to especially thank my wife, Katie, for her superb editorial

input. Lastly, a little chirp goes out to my orange tabby cat, Poppy, who kept me

company during long nights of writing and revising.

vii

TABLE OF CONTENTS

LIST OF TABLES ........................................................................................................ ix

LIST OF FIGURES........................................................................................................ x

CHAPTER 1

INTRODUCTION

1.1 Introduction .................................................................................................. 2

1.2 In Situ Biological Uranium Reduction ......................................................... 6

1.3. Scope of Thesis and Chapter Summaries .................................................. 10

1.4. Literature Cited.......................................................................................... 14

Appendix 1A: Supporting Information for Chapter 1 ...................................... 24

CHAPTER 2

THERMODYNAMIC CONSTRAINTS ON THE OXIDATION OF BIOGENIC UO2

BY FE(III) (HYDR)OXIDES

2.1 Introduction ................................................................................................ 35

2.2 Materials and Methods ............................................................................... 38

2.3 Results ........................................................................................................ 42

2.4 Discussion................................................................................................... 45

2.5 Implications for Uranium Natural Attenuation and Bioremediation.......... 48

2.6 Acknowledgements .................................................................................... 50


Appendix 2A: Supporting Information for Chapter 2 ...................................... 64

viii

CHAPTER 3

KINETIC AND MINERALOGICAL CONSTRAINTS ON THE OXIDATION OF

BIOGENIC URANINITE BY FERRIHYDRITE

3.1 Introduction ................................................................................................ 74

3.2 Materials and Methods ............................................................................... 77

3.3 Results ........................................................................................................ 83

3.4 Discussion................................................................................................... 87

3.5 Implications for Biogeochemical Uranium Cycling................................... 93

3.6 Acknowledgements .................................................................................... 95


Appendix 3A: Supporting Information for Chapter 3 .................................... 110

CHAPTER 4

KINETIC AND MINERALOGICAL CONSTRAINTS ON THE OXIDATION OF

BIOGENIC URANINITE BY FERRIHYDRITE

5.1 Introduction .............................................................................................. 119

5.2 Materials and Methods ............................................................................. 122

5.3 Results ...................................................................................................... 127

5.4 Discussion................................................................................................. 131

5.5 Environmental Implications ..................................................................... 134

5.6 Acknowledgements .................................................................................. 136

5.7. Literature Cited........................................................................................ 137

Appendix 5A: Supporting Information for Chapter 5 .................................... 147

CHAPTER 6

CONCLUSIONS

6.1 Conclusions .............................................................................................. 154

6.2 Literature Cited......................................................................................... 158

ix

LIST OF TABLES

CHAPTER 1

Table 1.1. Sediment Sample Characteristics ................................................................ 17

Table 1.2. Extractable Uranium.................................................................................... 18

CHAPTER 2

Table 2.1. Gibb’s Free Energy of Reactions ................................................................ 55

Table 2.1. Conditions at the conclusion of each UO2 oxidation experiment. .............. 56

APPENDIX 2

Table 2A.1. Gibb’s Free Energy of formation ............................................................. 65

Table 2A.2. References and relevant geochemical conditions used in Figure 2.5....... 66

CHAPTER 3

Table 3.1. Reaction conditions and observed first order rate coefficients ................. 100

CHAPTER 4

Table 4.1. Experimental Conditions for Column Experiments .................................. 140

Table 4.2. EXAFS Fitting Parameters ....................................................................... 141

x

LIST OF FIGURES

CHAPTER 1

Figure 1.1. Diagram of field-site wells......................................................................... 20

Figure 1.2. Uranium LIII –edge XANES spectra of sediment retrieved on February 24,

2005 .............................................................................................................................. 21

Figure 1.3. Representative Fe K-edge derivative XANES........................................... 22

Figure 1.4. Micro-scale uranium and iron distribution and speciation......................... 23

APPENDIX 1

Figure 1A.1. XRD Patterns of Solid-Phase Sediment Samples ................................... 32

CHAPTER 2

Figure 2.1. Representative Fe(III)/Fe(II) and U(VI)/U(IV) redox couples .................. 58

Figure 2.2. Free energy of reaction for UO2 (biogenic) oxidation as a function of pH ..... 60

Figure 2.3. Free energy of reaction for UO2 (biogenic) oxidation as a function of

bicarbonate concentration............................................................................................. 61

Figure 2.4. Effect of carbonate concentration on the oxidation of UO2 (biogenic) by

ferrihydrite.................................................................................................................... 62

Figure 2.5. Thermodynamic viability of UO2 (biogenic) oxidation by ferrihydrite for

conditions reported for various field sites and soil-column experiments..................... 63

xi

APPENDIX 2

Figure 2A.1. Soluble uranium speciation ..................................................................... 68

Figure 2A.2. Extractable Fe(II) and soluble U(VI) at the conclusion of the oxidation

reactions........................................................................................................................ 69

Figure 2A.3. Representative Fe(III)/Fe(II) and soluble U(VI)/U(IV) redox................ 70

CHAPTER 3

Figure 3.1. First order rate plots. ................................................................................ 102

Figure 3.2. Pseudo first-order rate comparisons......................................................... 104

Figure 3.3. Transformation of ferrihydrite by Fe(II) generated during uraninite

oxidation. .................................................................................................................... 106

Figure 3.4. Comparison of predicted uraninite dissolution rates to uraninite oxidation

rates............................................................................................................................. 108

Figure 3.5. Predicted U(IV) speciation in 3 mM KHCO3 in equilibrium with

UO2(am). ....................................................................................................................... 109

APPENDIX 3

Figure 3A.1. Representative Fe(III)/Fe(II) and U(VI)/U(IV) redox couples. ............ 111

Figure 3A.2. Comparison of extractable Fe(II) and dissolved U(VI) concentrations as a

function of uraninite concentration. ........................................................................... 112

Figure 3A.3. Fe K-edge EXAFS spectra and linear combination fits ........................ 113

xii

Figure 3A.4. Extractable Fe(II) produced during uraninite oxidation........................115

Figure 3A.5. Uranium LIII XANES spectra showing uranium oxidation................... 116

CHAPTER 4

Figure 4.1. Effluent metal concentration. ................................................................... 142

Figure 4.2. Chromium K-edge XANES, EXAFS, and Fourier-transformed EXAFS

spectra......................................................................................................................... 143

Figure 4.3. XRD patterns obtained from areas of high chromium concentration ...... 144

Figure 4.4. Microanalysis of mica grains ................................................................... 145

Figure 4.5 Effect of sediment pre-treatment on Cr(VI) breakthrough ....................... 146

APPENDIX 4

Figure 4A.1. Mass of Cr(III) retained by acid treated column................................... 148

Figure 4A.2. Effluent Fe(II) and Cr(VI) concentrations for 0.5 M HCl treated columns

at various influent chromium concentrations. ............................................................ 149

Figure 4A.3. Experimental Fe-EXAFS and linear combination fits .......................... 150

Figure 4A.4 X-ray photoelectron spectra of magnetic separates ............................... 151

Figure 4A.5 Batch Experiments ................................................................................. 152

xiii

CHAPTER 5

Figure 6.1. Conceptual model of uranium biogeochemical cycling.......................... 160

Figure 6.2. Representative redox potentials for predicting the stability of uraninite in

environmental systems. .............................................................................................. 161

1

CHAPTER 1

Introduction to “Defining Abiotic and Biotic Pathways

of Metal Redox Transformation in Natural Sediments”

Matthew Ginder-Vogel

Department of Geological and Environmental Sciences, Stanford University, Stanford, CA 94305

2

1.1 INTRODUCTION Despite centuries of study, fundamental biogeochemical processes controlling

the dynamics of contaminants within complex environmental media remain poorly

defined. In addressing the fate of environmental contaminants, one must consider the

complete system, rather than solely considering a simplified fraction of the soil.

However, in natural environments, elemental cycling may be mediated by a variety of

processes, inclusive of adsorption/desorption phenomena and oxidation/reduction

reactions (abiotic or biotic). The focus of my research has therefore been to identify

and characterize biotic and abiotic pathways of metal transformation in soils and

sediments.

Heavy metals and radionuclides introduced into the environment as a

consequence of the industrial revolution and nuclear age pose a unique long-term

threat to environmental quality. These contaminants have migrated into surface and

groundwater systems that serve human and ecological needs, and unlike organic

compounds, they cannot be degraded into innocuous compounds. Among these

contaminants, chromium and uranium are of particular concern and have been

identified as priority pollutants in soils and groundwater at most U.S. Department of

Energy (DOE) sites because of the large inventory present, the health risk associated

with each element, and their mobility relative to waste sources (1). Uranium and

chromium contamination is widespread and extensive, which makes understanding

processes that affect the environmental mobility of these elements imperative. The

hexavalent oxidation state of these two elements is most susceptible to environmental

transport and biological uptake; however, in their reduced states, Cr(III) and U(IV),

3

their solubility, and hence their mobility, are limited. Since the mobility of these

contaminants is largely determined by each contaminant’s valence, it is therefore vital

to identify the environmental processes affecting the oxidation state. Once identified,

these processes will further our understanding of contaminant natural attenuation and

allow predictive modeling of their fate and transport; additionally, they may be

exploited for in situ contaminant immobilization.

While the hexavalent forms of uranium and chromium are both quite soluble,

their speciation in groundwater shows marked differences. Uranium(VI) forms

several sparingly soluble complexes with phosphate (2,3) and readily forms inner-

sphere complexes with many transition metal hydroxides (4-7); however, its solubility

is particularly increased by complexation with carbonate, a common groundwater

ligand, and thus U(VI) remains subject to migration within oxic systems. Hexavalent

Cr exists primarily in groundwater systems as the oxyanion CrO42- (chromate) or its

protonated counterparts, which exhibit high water solubility over much of the

environmental pH range, are strong oxidants, and are known mutagens, teratogens,

and carcinogens (8).

The reduction of U(VI) and Cr(VI) may proceed through several geochemical

and biological processes. The abiotic (chemical) reduction of U(VI) is essentially

limited to aqueous sulfide (at low pH and low bicarbonate concentration) (9), sulfide

minerals (10), and surface bound Fe(II) (11). In contrast, hexavalent chromium can be

chemically reduced to trivalent chromium by aqueous and adsorbed Fe(II) (12,13),

organic matter (14), Fe(II)-bearing minerals (15-17), and sulfide compounds (18,19).

In addition, anaerobic bacteria can mediate the transformation of these redox sensitive

4

elements, either directly, by use as terminal electron acceptors in the absence of O2, or

indirectly, through reaction with reduced products of microbial metabolism (e.g. H2S,

Fe(II) and reduced organic species). For example, numerous common dissimilatory

metal (DMRB) and sulfate reducing bacteria (SRB), including Shewanella, Geobacter,

and Desulfovibrio species, couple the oxidation of organic matter and H2 to the

reduction of U(VI), resulting in the precipitation of uraninite (U(IV)O2) (20,21) a

sparingly soluble phase. In contrast to biological uranium reduction, biological

chromate reduction generally proceeds at a slower rate than chemical reduction, by

either sulfide, at pH < 5.5, or by Fe(II), at pH > 5.5 (22). Thus, in anaerobic

environments, it appears likely that chemical reduction pathways will be the major

avenue by which chromate is reduced, while uranium reduction is likely biologically

mediated; however, in arid, aerobic environments with limited organic carbon,

chromate reductants are essentially restricted to Fe(II)-bearing mineral phases,

primary in origin.

One must also consider potential U(IV) and Cr(III) oxidation pathways in

addition to reduction pathways, in order to predict their long-term stability. The only

naturally occurring oxidants of Cr(III) at neutral pH are Mn-oxides (23); in contrast,

biologically precipitated uraninite can be oxidized and remobilized by several

common environmental constituents. Molecular oxygen and Mn oxides rapidly and

extensively oxidize UO2 (24). Nitrate, a common co-contaminant with uranium (1),

not only impedes biological uranium reduction (25-27), but it also induces U(IV)

oxidation through the production of reactive intermediates (i.e. NO2-, NO, and N2O) in

denitrification (28) and through direct respiration (29). Recent work confirms that

5

Fe(III) (hydr)oxide minerals also play a role in the oxidation of biogenic U(IV)

(28,30-35).

The structure, speciation, and subsequent environmental mobility of chromium

and uranium will be dictated by the kinetics and thermodynamics of an intricate

network of abiotic and biotic reactions. The operative biogeochemical pathway may

be a result of complementary or opposing reaction mechanisms, which may not be

reflected in the broader redox signature and biogeochemical profile. Therefore, the

research presented in this thesis has been devoted to identifying and characterizing

biotic and abiotic pathways of uranium and chromium transformation in soils and

sediments.

6

1.2 IN SITU BIOLOGICAL URANIUM REMEDIATION

Large quantities of uranium have been introduced into the environment,

through waste disposal practices that were considered adequate at the time but are now

know to be short-sighted. For instance, trillions of gallons of acidic, uranium-bearing

waste was generated at the Y-12 Facility, Oak Ridge, TN and discharged, for 31 years,

into the unlined S-3 Ponds, resulting in uranium contamination of groundwater. The

aquifer underlying the Y-12 facility contains weathered saprolite overlying

interbedded shale and limestone. Groundwater flow at the field site occurs primarily

in thin zones of interconnected fracture networks with high hydraulic conductivity

(36). Differential weathering of the parent rock has resulted in these fracture

networks becoming coated with phyllosilicate minerals and transition metal

(oxyhydr)oxides, including biotite, illite, vermiculite, kaolinite, and goethite

A field site has been established at Area 3 of the U.S. Department of Energy

Environmental Sciences Research Program (ERSP) Field Research Center (FRC) in

Oak Ridge, TN, USA, to investigate the potential of in situ biological uranium

reduction and immobilization as a potential uranium remediation strategy. The field

site is located adjacent to the former S-3 ponds, which received trillions of liters of

acidic plating wastes from nuclear weapon manufacturing process. The details of

field site establishment and the initial phases of operation are described in detail by

Wu et al. (37) and Wu et al. (38). In brief, the ambient groundwater contains 84 – 210

µM uranium, is acidic (pH 3.4 – 4.5) and buffered by up to 20 mM aluminum, and

also contains up to 160 mM nitrate, 10 mM sulfate, and 25 mM calcium. Solid phase

uranium concentrations range up to 800 mg kg-1. To establish and maintain hydraulic

7

control, we installed a multiple well recirculation system, with an inner treatment loop

nested within an outer isolation loop (Figure 1.1). A region of the subsurface was

prepared for stimulation of biological activity by pumping groundwater to the surface

for treatment to remove nitrate, calcium and aluminum. The pH of the treated water

was then adjusted to approximately 6.0, and with additional flushing, the pH of the

groundwater increased to 5.5 - 6.0, and the nitrate concentrations fell to 0.5 – 1.0 mM.

These conditions were judged adequate for biostimulation, and ethanol (1 mM) was

intermittently added to the inner treatment loop. The ethanol addition initially

stimulated denitrification of matrix entrapped nitrate, and after two months, aqueous

uranium levels fell from 5 to 1 µM, and sulfate reduction ensued. Microbial analyses

confirmed the presence of denitrifying, sulfate-reducing, and iron-reducing bacteria

(38). Periodic additions of ethanol over the past four years have kept aqueous

uranium concentrations below the United States Environmental Protection Agency’s

(EPA) drinking water standard of 0.126 µM (39).

Prior to stimulation of biological uranium reduction at the Area 3 field site,

uranium was present solely in the hexavalent oxidation state (38). A decrease in the

dissolved U(VI) concentration upon the addition of ethanol was first noted after ~200

d of field site operation, and the presence of U(IV) was confirmed in the treatment

loop injection well (FW104) (38). Periodic ethanol injections continued for another

300 d, resulting in the accumulation of reduced uranium not only in the treatment loop

injection well, but also in a sampling well (FW101-2), and in the treatment loop

extraction well (FW024) (Figure 1.2, Table 1.1). Despite the presence of high

8

proportions of U(IV), uraninite was not observed in these samples (day 535) (data not

shown).

After several years of field site operation (774 d), we quantified the extent of

biologically reduced uranium throughout the subsurface system by examining wells

which had responded to past tracer tests (Table 1.1) (37). Uranium(IV) was detected

in all wells that have a good hydraulic connection to the treatment-loop injection well

(37). Reduced uranium was coincident with reducing conditions, as determined by the

presence of dissolved Fe(II) and sulfide, although, at this point in time, the quantity of

U(IV) did not exceed 51% in any of the sediments investigated (Table 1.1).

Furthermore, uraninite was observed in XRD patterns (Figure 1A.1) of sediments that

contained both high total uranium concentrations and the highest proportions of U(IV)

(Table 1.1). In addition to uranium reduction, Fe(II) production is noted in the

aqueous phase (39), and FeS is the predominant iron species in sediments retrieved

from the treatment loop injection well; however, Fe(III) predominates in two sampling

wells, with only minor amounts of FeS (Figure 1.3). Analysis of the micro-scale

distribution of uranium and iron reveals that both of these elements are

heterogeneously distributed, and that areas of high U(VI) concentration appear to be

correlated with areas of high Fe concentration (Figure 1.4). Interestingly, uraninite

diffraction peaks were not observed in areas of high or low uranium concentration

(data not shown).

Continued accumulation of reduced uranium was observed in two samples

retrieved from wells FW101-2 and FW102-3 on July 21, 2006 (day 1063). These

samples now contain 70% U(IV) (Table 1.2), rather than the ~50% U(IV) observed at

9

earlier time points. Subsequent extraction with anaerobic salt and bicarbonate

solutions results in the release of only < 1%, and ~5%, respectively, of the solid-phase

uranium (Table 1.2). These results indicate that minimal sorbed and easily dissolved

U(VI) is left in the sediments after biostimulation.

The long-term stability of biologically reduced uranium will be determined by

the complex interplay of soil and sediment mineralogy, aqueous geochemistry,

microbial activity, and potential U(IV) oxidants. Many of these factors have been

studied under laboratory conditions; however, the impact of these factors on uranium

cycling in natural, subsurface environments is still poorly understood. Factors that

may result in the reoxidation of biologically precipitated uraninite are explored in

Chapters 2 and 3. The pilot-scale uranium bioremediation system located at Area 3 of

the FRC provides a controlled subsurface environment in which these factors can be

further investigated. Microbial activity has produced high proportions of U(IV)

throughout the subsurface system remediation system and the precipitation of

uraninite. Moreover, uranium residing within the subsurface after nearly 3 y of

treatment, as both U(IV) and remnant U(VI), appears resistant to factors promoting

migration.

10

1.3 SCOPE OF THESIS AND CHAPTER SUMMARIES

The research presented in Chapters 2 through 4 examines the role of abiotic

(chemical) and biotic (bacterial) controls on uranium and chromium mobility in

environmental systems. The oxidation of biologically precipitated uraninite by Fe(III)

(hydr)oxides may limit uranium sequestration under mildly reducing conditions. This

oxidation pathway is examined in detail in the second and third chapters. The impact

of geochemical conditions on the energetic favorability and extent of uraninite

oxidation is investigated using a combination of thermodynamic calculations and

static batch reactions in Chapter 2. This reaction is further investigated in Chapter 3,

where I demonstrate that the rate of oxidation is controlled by the uraninite dissolution

rate. In Mechanisms responsible for the reduction and retention of chromate within

arid sediments are identified in Chapter 4. The research presented in this thesis

combines a fundamental approach to investigating the molecular-scale interaction of

heavy metal and radionuclide contaminants with field-based observations in order to

identify the operative biogeochemical pathways controlling contaminant mobility.

Chapter 2: Thermodynamic Constraints on the Oxidation of Biogenic UO2 by Fe(III)

(hydr)oxides

Uranium mobility in the environment is partially controlled by its oxidation

state, where it exists as either U(VI) or U(IV). In aerobic environments, uranium is

generally found in the hexavalent form, is quite soluble, and readily forms complexes

with carbonate and calcium. Under anaerobic conditions, common metal respiring

11

bacteria can reduce soluble U(VI) species to sparingly soluble UO2 (uraninite);

stimulation of these bacteria, in fact, is being explored as an in-situ uranium

remediation technique. However, the stability of biologically precipitated uraninite

within soils and sediments is not well characterized. Here we demonstrate that

uraninite oxidation by Fe(III) (hydr)oxides is thermodynamically favorable under

limited geochemical conditions. Our analysis reveals that goethite and hematite have

a limited capacity to oxidize UO2 (biogenic), while ferrihydrite can lead to UO2 (biogenic)

oxidation. The extent of UO2 (biogenic) oxidation by ferrihydrite increases with

increasing bicarbonate and calcium concentration but decreases with elevated Fe(II)(aq)

and U(VI)(aq) concentrations. Thus, our results demonstrate that the oxidation of UO2

(biogenic) by Fe(III) (hydr)oxides may transpire under mildly reducing conditions when

ferrihydrite is present. MGV performed all laboratory work and wrote the manuscript

while CS and SF provided editorial input.

Chapter 3: Kinetic and Mineralogical Constraints on the Oxidation of Biogenic

Uraninite by Ferrihydrite

The oxidation state of uranium plays a major role in determining uranium

mobility in the environment. Uranium is generally found in the hexavalent form, is

quite soluble, and readily forms complexes with calcium and carbonate in aerobic

environments. However, under anaerobic conditions, common metal respiring

bacteria can enzymatically reduce U(VI) to U(IV), resulting in the formation of

sparingly soluble UO2 (uraninite). Uranium(VI) reduction, therefore, has a prominent

role in uranium natural attenuation and is being explored as a potential uranium

12

remediation technique. The stability of biologically precipitated uraninite is critical

for determining the long-term fate of uranium and is not well characterized within

soils and sediments. Here, we demonstrate that uraninite oxidation by ferrihydrite, a

disordered Fe(III) (hydr)oxide, proceeds through a soluble U(IV) intermediate and

results in the concomitant production of Fe(II) and dissolved U(VI). Uraninite

oxidation rates are accelerated under conditions that increase its solubility, which

include high bicarbonate concentration and pH values deviating from neutrality.

Additionally, Fe(II) produced during uranium oxidation catalyzes the transformation

of ferrihydrite into goethite and lepidocrocite, which, combined with elevated Fe(II)(aq)

and U(VI)(aq), may ultimately limit UO2 oxidation over longer time frames. Thus, our

results demonstrate that UO2 oxidation by Fe(III) (hydr)oxides is controlled by the rate

of uraninite dissolution and that this process may limit uranium sequestration under

mildly reducing conditions. MGV performed all laboratory work and wrote the

manuscript while SF provided editorial input.

Chapter 4: Chromate Reduction and Retention Processes within Arid Subsurface

Environments

A number of chromate reductants, inclusive of Fe(II) (aq), magnetite, green

rust, ilmenite, and bacteria, have been extensively studied in model systems, but few

studies have investigated reduction processes within natural sediments, particularly

those from aerated systems. Accordingly, I examined chromate reduction by arid

sediments from the Hanford, WA site. Iron(III) (hydr)oxide coatings limit mineral

reactivity in this arid environment; high pH (pH = 14), representative of conditions

13

near high-level nuclear waste tanks, results in Fe(II) solubilization and concurrent

Cr(VI) reduction. The X-ray spectroscopic analysis of solids from columns that were

either acid pre-treated or included 10 M NaOH reveals that reduced chromium,

Cr(III), occurs in association with antigorite and lizardite, in addition to magnetite and

Fe(II)-bearing clay minerals. Additionally, in a column containing 10 M NaOH in the

feed solution, Cr(III) and Cr(VI) are found associated with portlandite, suggesting a

secondary mechanism for chromium retention at high pH. The data and analysis

presented herein provide novel information on reactions of chromate with ferrous-

bearing primary minerals within natural sediments. MGV wrote the manuscript and

performed all experiments except batch reactions which were performed by MM, all

other authors provided editorial input.

14

1.4 LITERATURE CITED

(1) Riley, R. G.; Zachara, J. M.; Wobber, F. J. "Chemical contaminants on DOE lands and selection of contaminant mixtures for subsurface science research," U.S. Department of Energy, 1992. (2) Langmuir, D. Uranium solution-mineral equilibria at low temperature with applications to sedimentary ore deposits. Geochim. Cosmochim. Acta 1978, 42, 547-569. (3) Sandino, A.; Bruno, J. The solubility of (UO2)3(PO4)2·4H2O(s) and the formation of U(VI) phospate complexes: Their influence in uranium speciation in natural waters. Geochim. Cosmochim. Acta 1992, 56, 4135-4145. (4) Moyes, L. N.; Parkman, R. H.; Charnock, J. M.; Vaughan, D. J.; Livens, F. R.; Hughes, C. R.; Braithwaite, A. Uranium Uptake from Aqueous Solution by Interaction with Goethite, Lepidocrocite, Muscovite, and Mackinawite: An X-ray Absorption Spectroscopy Study. Environ. Sci. Technol. 2000, 34, 1062-1068. (5) Bostick, B. B.; Fendorf, S.; Barnett, M. O.; Jardine, P. M.; Brooks, S. C. Uranyl Surface Complexes Formed on Subsurface Media from DOE Facilities. Soil Sci. Soc. Am. J. 2002, 66, 99-108. (6) Barnett, M. O.; Jardine, P. M.; Brooks, S. C.; Selim, H. M. Adsorption and transport of uranium(VI) in subsurface media. Soil Sci. Soc. Am. J. 2000, 64, 908-917. (7) Barnes, C. E.; Cochran, J. K. Uranium geochemistry in esturaine sediments: Controls on removal and release processes. Geochim. Cosmochim. Acta 1993, 57, 555-569. (8) Fendorf, S. Surface reactions of chromium in soils and waters. Geoderma 1995, 67, 55-71. (9) Hua, B.; Xu, H.; Terry, J.; Deng, B. Kinetics of uranium(VI) reduction by hydrogen sulfide in anoxic aqueous systems. Environ. Sci. Technol. 2006, 40, 4666-4671. (10) Wersin, P.; Hochella, M. F.; Persson, P.; Redden, G.; Leckie, J. O.; Harris, D. W. Interaction between aqueous uranium(VI) and sulfide minerals : Spectroscopic evidence for sorption and reduction. Geochim. Cosmochim. Acta 1994, 58, 2829-2843. (11) Liger, E.; Charlet, L.; Cappellen, P. V. Surface catalysis of uranium(VI) reduction by iron(II). Geochim. Cosmochim. Acta 1999, 63, 2939-2955. (12) Fendorf, S.; Li, G. Kinetics of chromate reduction by ferrous iron. Environ. Sci. Technol. 1996, 30, 1614-1617. (13) Buerge, I. J.; Hug, S. J. Influence of mineral surfaces on Cr(VI) reduction by iron(II). Environ. Sci. Technol. 1999, 33, 4285-4291. (14) Deng, B.; Stone, A. T. Surface-catalyzed Cr(VI) reduction: Reactivity comparisons of different organic reductants and different oxide surfaces. Environ. Sci. Technol. 1996, 30, 2484-2494.

15

(15) Peterson, M. L.; White, A. F.; Brown, G. E.; Parks, G. A. Surface passivation of magnetite by reaction with aqueous Cr(VI): XAFS and TEM results. Environ. Sci. Technol. 1997, 31, 1573-1576. (16) White, A. F.; Peterson, M. L. Reduction of aqueous transition metals on the surfaces of Fe(II)-containing oxides. Geochim. Cosmochim. 1996, 60, 3799-3814. (17) Bond, D. L.; Fendorf, S. Kinetics and structural constraints of chromate reduction by green rusts. Environ. Sci. Technol. 2003, 27, 2750-2757. (18) Eary, L. E.; Rai, D. Kinetics of chromate reduction by ferrous ions derived from hematite and biotite at 25 oC. Am. J. Sci. 1989, 289, 180-213. (19) Patterson, R. R.; Fendorf, S.; Fendorf, M. Reduction of hexavalent chromium by amorphous iron sulfides. Environ. Sci. Technol. 1997, 31, 2039-2044. (20) Lovley, D. R.; Phillips, E. J. P. Bioremediation of uranium contamination with enzymatic uranium reduction. Environ. Sci. Technol. 1992, 26, 2228-2234. (21) Gorby, Y. A.; Lovley, D. R. Enzymatic Uranium Precipitation. Environ. Sci. Technol. 1992, 26, 205-207. (22) Fendorf, S.; Hansel, C. M.; Wielinga, B. W. In Geochemistry of Soil Radionuclides; Zhang, P.-C., Brady, P. V., Eds.; Soil Sci. Soc. Am: Madison, WI, 2002; Vol. Special Publication No. 59. (23) Fendorf, S. E.; Zasoski, R. J. Chromium(III) oxidation by delta-MnO2. Environ. Sci. Technol. 1992, 26, 79-85. (24) Fredrickson, J. K.; Zachara, J. M.; Kennedy, D. W.; Liu, C. G.; Duff, M. C.; Hunter, D. B.; Dohnalkova, A. Influence of Mn oxides on the reduction of uranium(VI) by the metal-reducing bacterium Shewanella putrefaciens. Geochim. Cosmochim. Acta 2002, 66, 3247-3262. (25) Finneran, K. T.; Housewright, M., E.; Lovley, D. R. Multiple influences of nitrate on uranium solubility during bioremediation of uranium-contaminated subsurface sediments. Environ. Microbiol. 2002, 4, 510-516. (26) Senko, J. M.; Istok, J. D.; Suflita, J. M.; Krumholz, L. R. In-situ evidence for uranium immobilization and remobilization. Environmental Science & Technology 2002, 36, 1491-1496. (27) Istok, J. D.; Senko, J. M.; Krumholz, L. R.; Watson, D.; Bogle, M.; Peacock, A. D.; Chang, Y.-J.; White, D. C. In situ bioreduction of technetium and uranium in a nitrate contaminated aquifer. Environ. Sci. Technol. 2004, 38, 468-475. (28) Senko, J. M.; Mohamed, Y.; Dewers, T.; Krumholz, L. R. Role for Fe(III) minerals in nitrate-dependent microbial U(IV) oxidation. Environ. Sci. Technol. 2005, 39, 2529-2536. (29) Beller, H. R. Anaerobic, nitrate-dependent oxidation of U(IV) oxide minerals by the chemolithoautotrophic bacterium Thiobacillus dentrificans. Appl. Environ. Microb. 2005, 71, 2170-2174. (30) Nevin, K. P.; Lovley, D. R. Potential for Nonenzymatic Reduction of Fe(III) via Electron Shuttling in Subsurface Sediments. Environ. Sci. Technol. 2000, 34, 2472-2478.

16

(31) Wan, J.; Tokunaga, T. K.; Brodie, E.; Wang, Z.; Zheng, Z.; Herman, D.; Hazen, T.; Firestone, M. K.; Sutton, S. R. Reoxidation of bioreduced uranium under reducing conditions. Environ. Sci. Technol. 2005, 39, 6162-6169. (32) Sani, R. K.; Peyton, B. M.; Amonette, J. E.; Geesey, G. G. Reduction of uranium(VI) under sulfate-reducing conditions in the presence of Fe(III)-(hydr)oxides. Geochim. Cosmochim. 2004, 68, 2639-2648. (33) Sani, R. K.; Peyton, B. M.; Dohnalkova, A.; Amonette, J. E. Reoxidation of reduced uranium with iron(III) (hydr)oxides under sulfate-reducing conditions. Environ. Sci. Technol. 2005, 39, 2059-2066. (34) Senko, J. M.; Suflita, J. M.; Krumholz, L. R. Geochemical controls on microbial nitrate-dependent U(IV) oxidation. Geomicrobiology Journal 2005, 22, 371-378. (35) Ginder-Vogel, M.; Criddle, C.; Fendorf, S. Thermodynamic constraints on the oxidation of biogenic UO2 by Fe(III) (hydr)oxides. Environ. Sci. Technol. 2006, 40, 3544-3550. (36) Solomon, D.; Moore, G.; Toran, L.; RB, D.; McMaster, W. "Status Report: A hydrologic framework for the Oak Ridge Reservation," Oak Ridge Nationa Laboratory, 1992. (37) Wu, W.-M.; Carley, J.; Fienen, M.; Mehlhorn, T.; Lowe, K.; Nyman, J.; Luo, J.; Gentile, M.; Rajan, R.; Wagner, D.; Hickey, R.; Gu, B.; Watson, D. B.; Cirpka, O.; Kitanidis, P.; Jardine, P. M.; Criddle, C. Pilot-scale in situ bioremediation of uranium in a highly contaminated aquifer. 1. Conditioning of a treatment zone. Environ. Sci. Technol. 2006, 40, 3978-3985. (38) Wu, W.-M.; Carley, J.; Gentry, T.; Ginder-Vogel, M.; Fienen, M.; Mehlhorn, T.; Yan, H.; Caroll, S.; Pace, M.; Nyman, J.; Luo, J.; Gentile, M.; Fields, M. W.; Hickey, R.; Watson, D. B.; Cirpka, O.; Zhou, J.; Fendorf, S.; Kitanidis, P.; Jardine, P. M.; Criddle, C. Pilot-scale in situ bioremediation of uranium in a highly contaminated aquifer. 2. Geochemical control of U(VI) bioavailability and evidence of U(VI) reduction. Environ. Sci. Technol. 2006, 40, 3986-3995. (39) Wu, W.-M.; Carley, J.; Luo, J.; Ginder-Vogel, M.; Cardanans, E.; Leigh, M. B.; Hwang, C.; Kelly, S. D.; Ruan, C.; Wu, L.; Gentry, T.; Lowe, K.; Mehlhorn, T.; Carroll, S. L.; Fields, M. W.; Gu, B.; Watson, D.; Kemner, K. M.; Marsh, T. L.; Tiedje, J. M.; Zhou, J.; Fendorf, S.; Kitanidis, P.; Jardine, P. M.; Criddle, C. In situ bioreduction of uranium(VI): Impact of dissolved oxygen. Environ. Sci. Technol. 2006, Submitted.

17

Table 1.1. Biological activity, redox conditions, uranium content and uranium oxidation state for solid samples retrieved on February 25 and October 5, 2005. The methods used to collect the data are described in Appendix 1A.

Well Well Description Biological Activity

Reducing Conditions observed

October 5, 2005

(day 774)

U content (g/kg)

February 24, 2005 (day 535)

U content (g/kg)

October 5, 2005

% U(IV) Feb. 24,

2005

% U(IV) Oct 5, 2005

FW024 Isolation Loop Injection Low No -1 0.4 - 0 FW104 Treatment Loop Injection Very High Yes 4.32 10.3 51 61 FW 026 Treatment Loop Extraction Medium Slightly 1.14 1.2 28 0 FW 103 Isolation Loop Extraction Low No - 0.7 - 0 FW 105 Isolation Loop Extraction Low No - 1.0 - 0

FW 100-1 Sampling Well – Isolation Loop Low No - 0.2 - NA2 FW 100-2 Sampling Well – Isolation Loop Medium No - 1.0 - 0 FW 100-3 Sampling Well – Isolation Loop Medium No - 1.1 - 0 FW 100-4 Sampling Well – Isolation Loop Low No - 1.5 - NA FW 101-1 Sampling Well – Treatment Loop Low No - 0.05 - NA FW 101-2 Sampling Well – Treatment Loop High Yes 0.91 1.2 35 54 FW 101-3 Sampling Well – Treatment Loop High Yes - 1.8 - 51 FW 101-4 Sampling Well – Treatment Loop Low No - 0.6 - NA FW 102-1 Sampling Well – Treatment Loop Low No - 0.04 - NA FW 102-2 Sampling Well – Treatment Loop High Yes - 0.5 - 17 FW 102-3 Sampling Well – Treatment Loop High Yes - 0.9 - 30 FW 102-4 Sampling Well – Treatment Loop Low No - 0.5 - NA 1. No sample collected 2. Sample collected but not analyzed

18

Table 1.2. Uranium content, oxidation state, and extractable uranium for samples collected on July 21, 2006 (day 1063). The methods used to collect the data are described in Appendix 1A.

Well Strong Acid Extractable Fe

(g/kg)

U content (g/kg)

KCl Extractable U3 (% total)

HCO3 Extractable U2

(% total)

% U(IV)3

FW101-2 10 0.33 0.25 4.5 70 FW102-3 15 1.64 0.90 6.5 70

1. 100 mM KCL at pH 9 2. 100 mM KHCO3 at pH 9 3. As determined by XANES spectroscopy (Appendix 1)

19

Figure Captions Figure 1.1. Plan view (A) and schematic diagram (B) of the subsurface remediation system. The streamlines with arrows (solid lines) indicate flow direction, while the dark lines are streamlines that separate the outer and inner treatment zones (after Wu et al, 2006). The methods used to collect the data are described in Appendix 1A. Figure 1.2. Uranium LIII

–edge XANES spectra of sediment retrieved on February 24, 2005 from the isolation injection well (FW104), a monitoring well (FW101-2) and the isolation extraction well (FW026). The uranium concentration for each sample is listed in Table 1.1. The methods used to collect the data are described in Appendix 1A. Figure 1.3. Representative Fe K-edge derivative XANES of samples of samples retrieved on October 5, 2005. Samples from all other wells (Table 1.1) contained only Fe(III). The gray lines indicate the maximum for each of the three standards. The methods used to collect the data are described in Appendix 1A. Figure 1.4. Micro-scale uranium (A) and iron (B) distribution in samples obtained from sampling well FW101-2 on February 24, 2005. Warmer colors indicate more uranium/iron fluorescent X-rays. Background (blue) for both maps is ~5,000 uranium or iron counts per second (cps) with U hotspots containing 30,000 cps and Fe hot spots containing 12,000 cps. Micro-XANES spectra collected from points 1 and 2 (C). The methods used to collect the data are described in Appendix 1A.

20

Figure 1.1. Plan view (A) and schematic diagram (B) of the subsurface remediation system. The streamlines with arrows (solid lines) indicate flow direction, while the dark lines are streamlines that separate the outer and inner treatment zones (after Wu et al, 2006).

21

Energy (eV)17140 17150 17160 17170 17180 17190 17200

Nor

mal

ized

Flo

ures

cenc

e

0.0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

1.6

1.8

Uraninite [U(IV)]Injection Well - 51% U(IV)MLS Well - 35% U(IV)Extraction Well - 28% U(IV)Uranyl Nitrate [U(VI)]

Figure 1.2. Uranium LIII

–edge XANES spectra of sediment retrieved on February 24, 2005 from the isolation injection well (FW104), a monitoring well (FW101-2) and the isolation extraction well (FW026). The uranium concentration for each sample is listed in Table 1.1. The methods used to collect the data are described in Appendix 1A.

22

Energy (eV)

7100 7125 7150

Nor

mal

ized

Firs

t der

ivat

ive

Inte

nsity

Ferrihydrite - Fe(OH)3

Siderite - FeCO3

FeS

FW104

FW101-2

FW101-3

FW102-2

FW102-3

FW026

Figure 1.3. Representative Fe K-edge derivative XANES spectra of samples retrieved on October 5, 2005. Samples from all other wells (Table 1.1) contained only Fe(III). The gray lines indicate the maximum for each of the three standards. The methods used to collect the data are described in Appendix 1A.

23

Energy (eV)

17140 17150 17160 17170 17180 17190 17200

Nor

mal

ized

Flu

ores

cenc

e

0.0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

1.6

1.8

2.0

Uraninite [U(IV)]Uranyl Nitrate [U(VI)]Point 1Point 2

C

Figure 1.4. Micro-scale uranium (A) and iron (B) distribution in samples obtained from sampling well FW101-2 on February 24, 2005. Warmer colors indicate more uranium/iron fluorescent X-rays. Background (blue) for both maps is ~5,000 uranium or iron counts per second (cps) with U hotspots containing 30,000 cps and Fe hot spots containing 12,000 cps. Micro-XANES spectra collected from points 1 and 2 (C). The methods used to collect the data are described in Appendix 1A.

24

APPENDIX 1A

Supporting Information for Chapter 1

25

1A.1 MATERIALS AND METHODS

1A.1.1 Field Site

A field site has been established at Area 3 of the U.S. Department of Energy

Environmental Sciences Research Program (ERSP) Field Research Center (FRC) in

Oak Ridge, TN, USA, to investigate the potential of in situ biological uranium

reduction and immobilization as a potential uranium remediation strategy. The field

site is located adjacent to the former S-3 ponds, which received trillions of liters of

acidic plating wastes from nuclear weapon manufacturing process. The details of field

site establishment and the initial phases of operation are described in detail by Wu et

al. (12) and Wu et al. (13). In brief, the ambient groundwater contains 84 – 210 µM

uranium, is acidic (pH 3.4 – 4.5) and buffered by up to 20 mM aluminum, and also

contains up to 160 mM nitrate, 10 mM sulfate, and 25 mM calcium. Solid phase

uranium concentrations range up to 800 mg kg-1. To establish and maintain hydraulic

control, we installed a multiple well recirculation system, with an inner treatment loop

nested within an outer isolation loop (Figure 4.1). A region of the subsurface was

prepared for stimulation of biological activity by pumping groundwater to the surface

for treatment to remove nitrate, calcium and aluminum. The pH of the treated water

was then adjusted to approximately 6.0, and with additional flushing, the pH of the

groundwater increased to 5.5 - 6.0, and the nitrate concentrations fell to 0.5 – 1.0 mM.

These conditions were judged adequate for biostimulation, and ethanol (1 mM) was

intermittently added to the inner treatment loop. The ethanol addition initially

stimulated denitrification of matrix entrapped nitrate, and after two months, aqueous

uranium levels fell from 5 to 1 µM, and sulfate reduction ensued. Microbial analyses

26

confirmed the presence of denitrifying, sulfate-reducing, and iron-reducing bacteria

(13). Periodic additions of ethanol over the past four years have kept aqueous uranium

concentrations below the United States Environmental Protection Agency’s (EPA)

drinking water standard of 0.126 µM (30).

1A.1.2 Analytical Methods

Groundwater and sediment samples were collected as described by Wu et al.

(12) and Wu et al. (13). On February 24, 2005 (day 535 of field site operation), solid

samples were collected from the treatment (FW104) and isolation (FW024) loop

injection wells, the treatment (FW026) and isolation (FW103) loop extraction wells,

and one sampling well (FW101-2). On October 5, 2005 (day 774), samples were

retrieved from all wells (Figure 4.1), while on July 21, 2006 (day 1063), samples were

collected from sampling wells FW102-3 and FW101-2. All solid samples were

analyzed within one month of collection. Uranium content in solids was determined

by extraction with concentrated nitric acid (31), followed by kinetic phosphorescence

(KPA) analysis. Loosely bound and bicarbonate extractable U(VI) were determined

for the samples collected on July 21, 2006. Sediment (0.25 g) was extracted with

anaerobic solutions of 100 mM KCl and 100 mM KHCO3 for 48 h; dissolved U(VI)

was then measured spectroflourometically. Uranium samples were diluted 1:30 in

10% phosphoric acid, and the fluorescence of the uranyl-phosphate complex was

measured at 515.4 nm. All measurements were referenced to the fluorescence of the

background matrix. Strong acid extractable iron was determined by extraction of

27

sediment (0.25 g) for 48 h; iron was then quantified using inductively coupled plasma-

optical emission spectrometry (ICP-OES).

1A.1.3 X-ray Absorption Spectroscopy and Diffraction

X-ray absorption near-edge structure (XANES) spectroscopy was used to

determine the relative oxidation state of uranium and iron. Sediment samples were

centrifuged and mounted as a wet paste in Teflon sample holders and sealed with

Kapton polyamide tape in an anaerobic glovebox, to prevent sample oxidation while

minimizing X-ray absorption. Fluorescence data were collected on bearmline 11-2 at

the Stanford Synchrotron Radiation Laboratory (SSRL), using either a 30 element Ge

semiconductor detector (U) or a Lytle fluorescence chamber (Fe). Incident and

transmitted X-ray intensities were measured with in-line ionization chambers. The

energy range studied was -200 to +800 eV around the U LIIIα-edge of U (17,166 eV)

and the Fe K-edge (7,111 eV). All samples were internally referenced to either a

U(VI)-nitrate or Fe-metal standard and placed between the second and third in-line

ionization chambers. Two to four individual spectra were averaged for each sample.

Synchrotron micro-X-ray fluorescence (µ−XRF) mapping, micro-X-ray

absorption spectroscopy (µ−XAS), and micro-X-ray diffraction (µ-XRD)

measurements were performed on GSE-CARS beamline 13-ID-C at the Advanced

Photon Source (APS, Argonne, IL). The incident X-ray beam was focused to a size of

2 x 2 µm using two Si mirrors in a Kirkpatrick-Baez geometry. Sediments were

mounted between Lexan plates and attached to a x-y translation stage, the incident

beam intensity (Io) was measured with an in-line ionization chamber, and fluorescence

28

yield was measured, using a multi-element Ge solid-state detector and normalized by

I0. X-ray fluorescence spectra were recorded on selected regions of the samples, on

the basis of elemental associations obtained from µ−XRF maps. Microfocused X-ray

diffraction (µ-XRD) patterns were collected on select areas in transmission geometry,

using monochromatic radiation (20 keV) and a MAR 345 image plate. The resulting

images were processed, using the same techniques described below, for bulk X-ray

diffraction measurements.

XANES spectra were processed, using the SixPACK (32) interface to IFEFFIT

(33). XANES data were background-subtracted and normalized to a unit-edge step.

The relative amount of reduced uranium in each sample was determined by

comparison of the half-height edge position of each sample to a standard curve

obtained from samples with varying known mole ratios of U(VI)/U(VI); this method

has an uncertainty of ± 10% (34). The presence of Fe(II) and Fe(II)-sulfides was

determined by comparison of the first derivative of the Fe-XANES spectra with

ferrihydrite, siderite, and amorphous FeS standards.

X-ray diffraction (XRD) was used to identify crystalline phases in sediment

samples retrieved from the FRC. XRD patterns were collected on beamline 11-3 at

SSRL in transmission geometry, using monochromatic radiation (12,732 eV) and a

MAR 345 image plate. The resulting images were processed using FIT2D (35). The

sample-to-detector distance and geometric corrections were calculated from the

pattern of LaB6. After these corrections were applied, the 2D images were integrated

radially, to yield 1D powder diffraction patterns, which could then be analyzed using

standard techniques. Peak identification and background correction, including

29

removal of the scattering from the lexan window, were performed in JADE 6.5

(Materials Data, Inc., Livermore, CA). Samples were mounted in the anaerobic

chamber, between Lexan windows sealed with double-sided tape to limit sample

oxidation during analysis.

1A.2 RESULTS

1A.2.1 Sediment Characterization

The mineralogy of sediment samples retrieved from wells at Area 3 of the

FRC, as examined with synchrotron based X-ray diffraction, is remarkably

homogeneous, spatially and temporally. The only major crystalline minerals detected

are quartz and muscovite (Figures 1A.1). Uraninite is noted in wells FW104 and FW

102-3 on October 5, 2005 (Figure 1A.1). Neither crystalline iron(III) (hydr)oxide nor

U(VI) phases can be detected in the samples analyzed.

XANES spectroscopy was used to examine the relative oxidation state of

uranium in all sediment samples collected, in order to determine if the stimulation of

metal-reducing bacterial activity is resulting in reduction of uranium(VI) to

uranium(IV). After 535 d of periodic ethanol injection, reduced uranium is detected

throughout the inner circulation loop (Figure 1.2, Table 1.1). Periodic ethanol

injection was continued for another 198 d, after which the molecular oxygen was

excluded from the entire system for 41 d (ending on day 774), through treatment of the

injection water with sodium sulfite, in order to investigate the stability of the

subsurface remediation system without the addition of electron donor (30). Dissolved

uranium in treatment zone wells with high biological activity remained below the

30

EPA’s drinking water limit of 0.126 µM during this time period (30). XANES

analysis of sediment retrieved at the end of this period reveals that the sediment in

wells with high levels of biological activity contain up to 61% U(IV) (Table 1.1).

Continued ethanol injection resulted in the accumulation of more U(IV) with

approximately 70% of the uranium resided in the U(IV) oxidation state on July 21,

2006 (Table 1.2). Extraction of the samples with anaerobic pH 9 KCl results in the

release of less than 1% of solid-phase uranium (36). However, extraction with 100

mM KHCO3 at pH 9 released 4.5 and 6.5% of the solid-phase uranium for wells

FW101-2 and FW102-3 respectively (Table 1.1).

In addition to the oxidation state of uranium, iron biogeochemistry is a key

factor in determining the mobility of uranium in subsurface environments (22);

therefore, we also analyzed the oxidation state of iron in the samples obtained on

October 5, 2005. Solid-phase Fe(II)-sulfide is detectable only in samples obtained

from the inner loop injection well (FW104) and two of the sampling wells (FW101-2

and FW101-3) (Figure 1A.1). In all other wells, solid-phase Fe is present solely as

Fe(III), within the detection limits of this analytical technique (~10%) (37).

1A.2.2 Micro-scale Analysis

Synchrotron micro-X-ray fluorescence (µ-XRF) mapping of the uranium

distribution reveals that uranium is present throughout the sediments obtained from

sampling well FW102-3 (Figure 1.4). However, there are several spots, ranging from

50 to 100 µm in diameter, with uranium counts that are approximately six times more

than the background (Figure 1.4). Although iron is also distributed throughout the

31

sample, iron concentration increases by a factor of ca. 1.5 in the areas with high

uranium counts (Figure 1.4). Uranium XANES analysis of these spots reveals that at

least 90% of the uranium is present in the hexavalent oxidation state (Figure 1.4).

Micro-X-ray diffraction (µ-XRD) patterns, collected from areas of high uranium

concentration in aged samples, reveal the presence of quartz and muscovite.

Interestingly, uraninite diffraction peaks were not observed, either in areas of high or

low uranium concentration (data not shown), likely due to the microcrystalline nature

of biologically precipitated uraninite (38).

32

FW024

FW026

FW104

FW100-1

FW100-2

FW100-3

FW101-1

FW100-4

FW101-2

FW101-3

FW101-4

2θ10 12 14 16 18 20

Nor

mal

ized

Inte

nsity

FW102-1

FW102-2

FW102-3

FW102-4

FW103

FW105Q

QMM M M M

MM

Uraninite (111)

Figure 1A.1. XRD patterns for solid-phase samples retrieved October 5, 2005. Minerals identified include quartz (Q) and muscovite (M). The gray line indicates the position of the uraninite(111) reflection, which is the most intense uraninite feature. The uranium content for each sample is listed in Table 1.1.

33

CHAPTER 2

Thermodynamic Constraints on the Oxidation of Biogenic UO2 by Fe(III) (hydr)oxides

Matthew Ginder-Vogel1, Craig Criddle2, and Scott Fendorf1

1. Department of Geological and Environmental Sciences, Stanford University, Stanford, CA 94305

2. Department of Civil and Environmental Engineering, Stanford University, Stanford, CA 94305

Published as: Ginder-Vogel, M.; Criddle, C.S.; Fendorf, S. Thermodynamic constraints on the oxidation of biogenic UO2 by Fe(III) (hydr)oxides. Environ. Sci. Technol. 2006, 40, 3544-3550.

34

ABSTRACT

Uranium mobility in the environment is partially controlled by its oxidation

state, where it exists as either U(VI) or U(IV). In aerobic environments, uranium is

generally found in the hexavalent form, is quite soluble, and readily forms complexes

with carbonate and calcium. Under anaerobic conditions, common metal respiring

bacteria can reduce soluble U(VI) species to sparingly soluble UO2 (uraninite);

stimulation of these bacteria, in fact, is being explored as an in-situ uranium

remediation technique. However, the stability of biologically precipitated uraninite

within soils and sediments is not well characterized. Here we demonstrate that

uraninite oxidation by Fe(III) (hydr)oxides is thermodynamically favorable under

limited geochemical conditions. Our analysis reveal that goethite and hematite have a

limited capacity to oxidize UO2(biogenic) while ferrihydrite can lead to UO2(biogenic)

oxidation. The extent of UO2(biogenic) oxidation by ferrihydrite increases with increasing

bicarbonate and calcium concentration, but decreases with elevated Fe(II)(aq) and

U(VI)(aq) concentrations. Thus, our results demonstrate that the oxidation of

UO2(biogenic) by Fe(III) (hydr)oxides may transpire under mildly reducing conditions

when ferrihydrite is present.

35

2.1 INTRODUCTION

A consequence of the nuclear age has been the release of radionuclides, toxic

heavy metals, and organic co-contaminants into the environment, posing a unique,

long-term environmental problem (1,2). Among these contaminants, uranium is of

particular concern because of its carcinogenicity, long half-life, widespread

distribution, and mobility (1). The sites involved with the United States’ nuclear

weapons production program, for example, are highly polluted with uranium. Uranium

contamination spreads far beyond the U.S. Department of Energy’s weapons complex,

however, and contains more than 24 uranium mine tailings sites in eleven states—

including Navajo and Hopi tribal lands (3). Unfortunately, uranium has migrated into

groundwater and contaminated more than 10 billion gallons of freshwater (1,4).

Uranium is generally found in the hexavalent form within oxic groundwater.

Uranium(VI) is quite soluble, and its solubility is particularly enhanced upon

complexation with carbonate, a common groundwater ligand (5). However, U(VI)

forms several sparingly soluble complexes with phosphate (6,7) and readily forms

inner-sphere complexes with many transition metal hydroxides(8-11). Nevertheless,

U(VI) tends to be relatively soluble and thus subject to migration within groundwater.

Conversely, U(IV) is sparingly soluble, even in the presence of common groundwater

ligands (12,13), and thus tends to be relatively immobile.

In part, the mobility of uranium in the environment will be controlled by its

oxidation state. Abiotic chemical reduction of uranium(VI) is essentially limited to

sulfide minerals (14) and surface bound Fe(II) (15). However, numerous, common

36

dissimilatory metal and sulfate reducing bacteria (DMRB and SRB), including

Shewanella, Geobacter, and Desulfovibrio species, couple the oxidation of organic

matter and H2 to the reduction of U(VI), resulting in the precipitation of (UO2

(biogenic uraninite) (16-18), a sparingly soluble phase. Thus, microbial U(VI)

reduction may play an important role in the cycling and, particularly, natural

attenuation of uranium. Additionally, current research is exploring the potential of in

situ uranium remediation by stimulating native metal reducing bacterial populations

(19,20).

The stability of biogenic uraninite is critical for determining the viability of

uranium remediation via in-situ biological uranium precipitation and for discerning

natural uranium cycling. While uraninite has an estimated solubility product of 10-8 to

10-12 (12,13), it can be rapidly oxidized and remobilized by a variety of common

environmental constituents (21,22). In addition to molecular oxygen, for example, Mn

oxides rapidly and extensively oxidize UO2 in the absence of biological activity (21).

Nitrate, a common co-contaminant with uranium (1), not only impedes biological

uranium reduction (23-25), but induces U(IV) oxidation through the production of

reactive intermediates (i.e. NO2-, NO, and N2O) in denitrification (22).

Under common groundwater conditions, the redox couples for U(IV)/U(VI)

and Fe(III) oxide/Fe(II) occur at similar potentials (Figure 2.1), and therefore, small

changes in aqueous chemistry can result in oscillations between thermodynamic

viability and non-viability of UO2 oxidation by Fe(III) (hydr)oxides (Figure 2.1).

Senko et al. (22) observed oxidation of biogenic U(IV) and concomitant release of

Fe(II) in the presence of Fe(III) (hydr)oxide minerals, which were produced by nitrate

37

dependent oxidation of Fe(II). Additionally, Fe(III) (hydr)oxide minerals have been

implicated in U(IV) oxidation under sulfate-reducing (26,27) and methanogenic

conditions (28). However, the evidence for abiotic U(IV) oxidation by Fe(III)

(hydr)oxides is not conclusive. Interpretation of abiotic U(IV) oxidation experiments

may be convoluted by several factors, including, but not limited to, high carbonate

concentrations, actively metabolizing microbes, Mn(IV) doping in Fe(III) oxides, and

remnant microbial and nitrate reduction intermediates (NO2-, N2O, and NO). In order

to determine conditions under which UO2(biogenic) may be oxidized by common iron

(hydr)oxides, we experimentally examine the abiotic oxidation of biogenic UO2 by

various environmentally relevant ferric-iron phases and present a detailed

thermodynamic analysis of Fe(III) (hydr)oxide promoted U(IV) oxidation.

38

2.2 MATERIALS AND METHODS

2.2.1 Biogenic Uraninite

All chemicals were analytical grade or better. Cell suspensions were prepared

by growing Shewanella putrefaciens strain CN32 (CN32) aerobically on Tryptic Soy

Broth at 30 oC to late-log phase. Cells were harvested by centrifugation (4000 g, 5

minutes), washed twice in 100 mL of anaerobic bicarbonate buffer (24 mM KHCO3,

pH 7), and resuspended in bicarbonate buffer. Uranium reduction was initiated by

inoculating 1 L of anaerobic U(VI) reduction media (pH = 7, 4 mM uranium acetate,

30 mM KHCO3, 10 mM PIPES, 0.03 mM NH4Cl, 40 mM lactate, and 10 mL Wolfe’s

vitamins) with 100 mL of bacterial suspension (108 cells mL-1). Media was then

stirred continuously in an anaerobic glove box with an atmosphere of 95% N2 and 5%

H2 (Coy Laboratory Products). After 4 d, the solids were collected and incubated in

10 % NaOH for 3 d to digest cell material, washed three times in 24 mM KHCO3, and

then washed twice in degassed, deionized (DI) water. X-ray diffraction (XRD)

patterns were identical to those previously reported for biogenic UO2 (16).

2.2.2 Ferric (Hydr)oxide Preparation

Two-line ferrihydrite was synthesized by rapidly titrating a ferric chloride

solution with NaOH to a pH of 7.5. Goethite was prepared by slow oxidation of a

ferrous chloride/sodium bicarbonate solution (29). Hematite was made by slow

titration of a ferric nitrate solution in boiling DI water (29). The Fe(III) (hydr)oxide

flocs were washed by centrifugation twice with 1% HCl and three times with distilled

39

H2O. Oxide mineralogy was confirmed with XRD. The BET surface areas of the

ferrihydrite, goethite, and hematite were 219, 80, and 60 m2 g-1, respectively.

2.2.3 Oxidation Experiments

Two sets of oxidation experiments were performed, one set to asses UO2

oxidation by various iron (hydr)oxides and a second set to asses the impact of

carbonate concentration on biogenic uraninite oxidation. All solutions used in the

study were prepared using distilled deionized water that had been treated to remove

dissolved O2 by boiling the solution while purging with nitrogen gas that had been

passed over hot Cu-metal filings. The solution was cooled for 12 h while being purged

with N2, then immediately transferred to the anaerobic glovebox. Incubations to

assess biogenic UO2 oxidation by various iron (hydr)oxides were performed in 10 mL

of bicarbonate buffer (3 mM, pH 7.0) under a headspace of N2. The impact of

carbonate concentration on UO2 oxidation by ferrihydrite was assessed using 10 mL

buffer with potassium bicarbonate concentrations of either 3, 10, 50, or 100 mM.

Biogenic U(IV) and the appropriate iron (hydr)oxide were added from wet slurries to a

final concentration of 120 µmoles L-1 U and 450 µmoles L-1 Fe. Control reactions

containing only biogenic UO2 or Fe(III) (hydr)oxide were also performed. All

oxidation experiments were performed within serum vials sealed with thick butyl

stoppers and wrapped in aluminum foil. The headspace in each vial was 5 mL; less

than 0.1% of dissolved bicarbonate was lost to the headspace. Vials were shaken on a

rotary shaker at 20 rpm for the duration of the experiment.

40

2.2.4 Sampling and Analytical Techniques

Given the propensity for the reaction of ferrous-iron with Fe(III) (hydr)oxides,

particularly ferrihydrite (30), and the tendency of uranyl ions to absorb on iron

(hydr)oxides (31,32), we measured the extent of UO2 oxidation by changes in

ferrozine-extractable Fe(II). Measuring Fe(II) after acidic dissolution of the reaction

product does not provide an accurate picture of the extent of UO2 oxidation because

the reaction becomes more energetically favorable at low pH (see Figure 2.2). After 24

h, incubations were destructively sampled and pH was measured—no change was

observed in the experiments. For aqueous U(VI) and Fe(II), a 5 mL aliquot was

withdrawn and filtered using a 0.2 µM polyethylene filter. Total soluble Fe and U

were analyzed in the filtered samples by inductively coupled plasma optical emission

spectrometry (ICP-OES). Soluble Fe(II) was analyzed using the ferrozine method

(33), while soluble U(VI) was measured using spectrofluorometry. Uranium samples

were diluted 1:30 in 10% phosphoric acid, and the fluorescence of uranyl-phosphate

complexes was measured at 515.4 nm in emission acquisition mode. All

measurements were referenced to the fluorescence of the background matrix. Solid-

phase associated Fe(II) was measured by extracting 1 mL of the reaction slurry with 3

mL of ferrozine reagent (33) for 1 min, which was then filtered with a 0.2 µm

polyethylene filter and analyzed for Fe(II) using a UV-vis spectrophotometer.

X-ray absorption spectroscopy was used to confirm the oxidation state of

biogenic UO2, using a previously described method (19,34). X-ray absorption spectra

41

were collected on beamline 11-2 at the Stanford Synchrotron Radiation Laboratory.

The oxidation state of the biogenic UO2 used in all experiments was approximately

96% U(IV) and 4% U(VI) . Subsequent measurement of the U(VI) content in the

biogenic UO2 by extraction with 100 mM KHCO3 confirmed the XANES analysis.

2.2.5 Thermodynamic Calculations

Except for the calcium-uranyl-carbonate species, Gibb’s free energies of

formation for all uranium species were obtained from Guillaumont et al. (35).

Amorphous UO2 (UO2 (am)) was chosen as the representative U(IV) species for all

thermodynamic calculations because freshly bioreduced U(IV) is generally fine-

grained and poorly crystalline (21,36,37). While UO2(am) is not thermodynamically

well-defined, it provides a more realistic prediction of reactivity over shorter time-

scales than crystalline UO2 (uraninite) (28). The Gibb’s free energy of formation for

Ca2UO2(CO3)3 and CaUO2(CO3)32- were calculated from stability constants provided

in Bernhard et al. (38). The values for hematite (Fe2O3) and goethite (α-FeOOH) were

obtained from Cornell and Schwertmann (39), while the value for 2-line ferrihydrite

(Fe(OH)3) was obtained from Majzlan et al. (40). All values were checked for internal

consistency and are listed in the supporting information (Table 2A.1). The Gibb’s free

energy of reaction for specific conditions was calculated using standard convention at

298 K. Aqueous speciation of the uranyl-cation under various geochemical conditions

was calculated using Visual Minteq with thermodynamic values noted in Table 2A.1

(41).

42

2.3 RESULTS

2.3.1 Thermodynamic Calculations

The Gibb’s free energy of reaction for UO2 (am) oxidation by ferrihydrite,

goethite, and hematite was calculated as a function of pH. The geochemical conditions

used in these calculations are representative of reducing groundwater having a

composition of 1 µM U(VI), 3 mM HCO3-, 1 mM Ca2+, and 10 µM Fe(II). Under

these conditions only ferrihydrite is capable of oxidizing UO2(am) at pH > 6 (Figure

2.2). However, at more acidic pH values, UO2(am) oxidation by goethite (pH < 6 ) and

hematite (pH < 5.5) becomes thermodynamically favored (Figure 2.2).

Aqueous U(VI) speciation plays a critical role in determining the

thermodynamic favorability of oxidation/reduction (Figure 2.1). Independent of iron

(hydr)oxide type, at circum-neutral pH oxidation of UO2 to Ca2UO2(CO3)3 is the most

thermodynamically favored reaction (Figures 2.1 and 2.2). In the absence of Ca2+, the

most thermodynamically favored oxidation reactions proceed with UO2(CO3)22- and

UO2(CO3)34- as products (Figures 2.1 and 2.2). At lower pH, UO2 oxidation to UO2

2+

and UO2CO3 become the most energetically favored reactions (Figure 2.2). The

energetic favorability of the oxidation reactions also increases with decreasing

aqueous Fe(II) and U(VI) concentrations for all of the specific reactions evaluated

(Figure 2.1).

Due to the stability of aqueous uranyl carbonate complexes, bicarbonate

concentration also plays an important role in determining the energetic favorability of

UO2 oxidation by Fe(III) (hydr)oxides. The oxidation reaction becomes more

energetically favorable with increasing carbonate concentration (Figure 2.3). An

43

exponential relation exists between the free energy of oxidation and carbonate

concentration, with 50% of the total free energy change that occurs between 1 and 100

mM occurring at < 10 mM carbonate (Figure 2.3). In addition to directly changing the

energetic favorability of the oxidation reaction, changes in the pH, carbonate, and

calcium concentration also affect the energetic favorability of the oxidation reaction

by changing the dominant aqueous U(VI) species. For instance, in the presence of

calcium, Ca2UO2(CO3)3 is the dominant species above pH 6 at all carbonate

concentrations. However, in the absence of calcium, uranyl carbonate species

dominate above pH 4.5 at all carbonate concentrations (Figure 2A.1).

2.3.2 UO2 Oxidation by Ferric (Hydr)oxides

The ability of ferrihydrite, goethite, and hematite to oxidize UO2 in the absence

of calcium was investigated in a series of batch experiments performed in 3 mM

HCO3- at pH 7. Production of soluble U(VI) is most extensive in the ferrihydrite

system (0.70 µM), with lesser amounts detected in goethite (0.50 µM) and hematite

(0.38 µM) systems (Table 2.2, Figure 2A.2). Dissolved U(VI) may result from either

oxidation of U(IV) or liberation of U(VI) in the biogenic UO2, of which 4% of U

remains in the hexavalent state. Reduction of Fe(III) and concurrent production of

Fe(II) therefore serves to determine the contribution of each uranium oxidative-

dissolution pathway. Extensive Fe(III) reduction, resulting in 23.3 µM extractable

Fe(II), occurs after UO2 exposure to ferrihydrite (Table 2.2, Figure 2A.2). Minimal

Fe(II) evolves in the presence of goethite (2.8 µM extractable Fe(II)), and none is

present with hematite as the oxidant (Table 2.2, Figure 2A.2). In all systems total

44

dissolve uranium and dissolved U(VI) were equal indicating minimal dissolved

U(IV).

The effect of carbonate concentration on UO2 stability in the presence of

ferrihydrite was investigated at four bicarbonate concentrations (3, 10, 50 and 100

mM). Soluble U(VI) concentrations increase with higher carbonate concentration in

the presence and absence (control) of ferrihydrite (Figure 2.4, Table 2.2), reflecting

the 4% U(VI) in the biogenic UO2. Extractable and soluble Fe(II) concentrations also

increase with increasing carbonate concentration (Figure 2.4, Table 2.2), while neither

extractable nor dissolved Fe(II) is detected in UO2 or Fe(III) (hydr)oxides free control

samples.

The free energy of reaction was calculated for each of the major soluble U(VI)

species, using the reactions presented in Table 2.1 together with the HCO3-, Fe(II)(aq),

and U(VI)(aq) concentrations at the conclusion of the oxidation experiments (Table

2.2). At all carbonate concentrations, other than 50 and 100 mM, the Gibb’s free

energy of reaction is ~0, indicating that most energetically favored oxidation reaction

has reached a near equilibrium state (Table 2.2).

45

2.4 DISCUSSION

The long-term stability of biologically reduced uranium in natural subsurface

environments will be dependent on both the prevalence of materials capable of

oxidizing UO2 (biogenic) and the geochemical characteristics of the groundwater. In the

case of UO2 oxidation by Fe(III) (hydr)oxides, the thermodynamic viability of the

reaction is dependent not only on Fe(III) (hydr)oxide mineralogy, but also on the

aqueous concentrations of Ca2+, Fe(II), U(VI), and HCO3- (Figures 2.1 and 2.3, and

Table 2.1).

The presence of Fe(II) limits UO2 (biogenic) oxidation through two mechanisms:

Fe(III) (hydr)oxide transformation and a decrease in the concentration gradient.

Indeed, high (> 0.100 mM) Fe(II) limits UO2 oxidation by Fe(III) (hydr)oxides

generated through microbial nitrate reduction (42). Iron(II)-induced transformations of

poorly crystalline Fe(III) (hydr)oxides, such as ferrihydrite, lead to the formation of a

complex assembly of minerals that, depending on initial chemistry and Fe(II)

concentration, is often dominated by magnetite or goethite (Figure 2.1) (43-45), both

of which are less favorable oxidants than ferrihydrite (39). Additionally, an increase in

the Fe(II) concentrations diminishes the thermodynamic favorability of UO2 oxidation

(Figure 2.1). Further, oxidation of UO2 by ferrihydrite after 24 h approaches

equilibrium in 3 mM and 10 mM bicarbonate (Table 2.2), while the oxidation of UO2

in 50 and 100 mM HCO3- remains energetically favorable for extended time periods

and increased Fe(II) concentrations (Table 2.2). Oxidation of UO2 by goethite is not

energetically favorable after 24 h reaction period; the presence of extractable solid-

phase Fe(II) indicates that a small amount of UO2 was oxidized by goethite (Table 2.2,

46

Figure 2A.2). The oxidation of UO2 is thermodynamically unfavorable for the

geochemical conditions listed in Table 2.2 at aqueous Fe(II) concentrations > 100 nM

for goethite and > 10 nM for hematite. Although dissolved U(IV) was not observed in

this study, it should be noted that the oxidation of U(OH)4, and U(CO3)44- to dissolved

U(VI) species is thermodynamically more favorable than the oxidation of UO2(am)

(Figure 2A.3) and may be an important U(IV) oxidation pathway at high pH and/or

very high carbonate concentrations (> 500 mM).

At carbonate concentrations exceeding 30 mM, UO2 oxidation in the presence

of ferrihydrite, goethite, and hematite has been noted (22,26,28). Carbonate

concentrations of this magnitude increase the thermodynamic favorability of UO2

(biogenic) oxidation by changing the speciation of the oxidation product to UO2(CO3)34-

or UO2(CO3)22- rather than UO2CO3 or (UO2)2CO3(OH)3

- (Table 2.1, Figure 2A.1)

and, of course, increase the concentration gradient favoring reaction products. The

role of carbonate concentration in increasing UO2 (biogenic) oxidation is further

exemplified by column experiments; U in the sediment was 87 (± 26 %) U(IV) at 8

mM carbonate and only 58 (± 22 %) U(IV) when carbonate concentrations exceeds 13

mM (28). Moreover, the impact of carbonate concentration on UO2 (biogenic) oxidation

by Fe(III) (hydr)oxides is illustrated by high U(VI) (aq) and ferrozine extractable Fe(II)

at 100 mM carbonate (Figure 2.4). At lower bicarbonate concentrations (3 and 10

mM), oxidation is again noted but must be viewed by production of Fe(II) (soluble and

ferrozine extractable) owing to U(VI) sorption on mineral surfaces, which is only

exceeded at the highest carbonate concentrations studied (Figure 2.4, Table 2.2).

Additionally, monitoring uranium as a sole indicator of oxidation is complicated by

47

the increased efficiency of U(VI) dissolution from the biogenic UO2 used here

(hexavalent uranium constitutes 4% of the total U within the solid) with increased

carbonate concentrations (46).

It is important to note that two common methods of extracting U(VI) from

solid phase samples, acidic dissolution and strong bicarbonate wash, may artificially

bias the measurement of Fe(II) and U(VI) in systems containing ferric (hydr)oxides

and UO2 (biogenic). Both of these extraction methods result in the oxidation of UO2 by

Fe(III) (hydr)oxides becoming energetically favorable (Figures 2.2 and 2.3).

48

2.5 IMPLICATIONS FOR URANIUM NATURAL ATTENUATION AND BIOREMEDIATION The importance of geochemical conditions on the reoxidation of biogenic UO2

by Fe(III) (hydr)oxides is exemplified by field and soil-column experiments (Figure

2.5). Surprisingly, for nine different conditions reported, oxidation of biogenic UO2

by Fe(III) (hydr)oxides is thermodynamically probable, depending on the aqueous

Fe(II) concentration. For example, for carbonate concentrations exceeding 30 mM (as

measured in refs. 22, 25 and 29 and portrayed in Figure 2.5), more than 200 µM

Fe(II) is required to maintain dissolved uranium concentrations below the

Environmental Protection Agency’s (EPA) Maximum Contaminant Level (MCL)

drinking water limit (0.126 µM) if ferrihydrite is present, even in the absence of

calcium. In biostimulated subsurface environments, transient Fe(II) concentrations

may exceed 200 µM (25); however, precipitation of Fe(II)-bearing mineral phases

such as FeCO3 (siderite) will ultimately control aqueous Fe(II) levels at 63 µM at

circumneutral pH and bicarbonate concentrations of 3 mM HCO3-. If carbonate

concentrations are less than 15 mM (as observed by refs. 19, 20, and 27 and projected

in Figure 2.5), between 10 and 100 µM Fe(II) are required to maintain soluble U(VI)

below 0.126 µM in the absence of Ca2+. Maintenance of low HCO3- concentrations in

the presence of Ca2+ is particularly important; even HCO3- concentrations as low as

1.2 mM require greater than 200 µM Fe(II) to maintain U(VI) concentrations below

the MCL for uranium (Figure 2.5). Therefore, if UO2 oxidation by Fe(III)

(hydr)oxides is controlling soluble U(VI) concentrations, maintenance of soluble

49

U(VI) concentrations below the EPA’s MCL of 0.126 µM (30 µg U / L) requires a

combination of low Ca2+, low HCO3-, and high Fe(II) concentrations.

The thermodynamic favorability of UO2 oxidation by ferric (hydr)oxides is

highly variable and slight changes in pH, aqueous Fe(II), HCO3-, or U(VI)

concentrations can oscillate the reaction between being thermodynamic viable and

non-viable. In fact, the oxidation reaction itself is self-suppressing due to the

generation of Fe(II) and U(VI). However, secondary sinks of Fe(II), such as mineral

surfaces or the generation of ferrous bearing precipitates (e.g., siderite), will alleviate

thermodynamic suppression of the oxidation reaction. Nevertheless, the impact of

UO2 oxidation by Fe(III) (hydr)oxides can be limited by low Ca2+ (< 1 mM), low

HCO3- (< 3mM), and reasonably high aqueous Fe(II) (>100 µM) concentration.

However, oxidation by ferric (hydr)oxides is thermodynamically viable and thus may

limit uranium sequestration under mildly reducing conditions while Fe(III) persists.

50

2.6 ACKNOWLEDGEMENTS

We would like to thank four anonymous reviewers and the Associate Editor,

Janet Hering, for their excellent suggestions that strengthened the manuscript. We also

thank Kathleen Beman and Thomas Borch for their constructive input on the

manuscript. This work was funded by the Environmental Remediation Sciences

Program, U.S. Department of Energy (grant numbers ER63609-1021814 and

DOEAC05-00OR22725). A portion of this work was conducted at Stanford

Synchrotron Radiation Laboratory, a national user facility operated by Stanford

University on behalf of the U.S. Department of Energy, Office of Basic Energy

Sciences. The SSRL Structural Molecular Biology Program is supported by the

Department of Energy, Office of Biological and Environmental Research, and by the

National Institutes of Health, National Center for Research Resources, Biomedical

Technology Program.

51

2.7 LITERATURE CITED (1) Riley, R. G.; Zachara, J. M.; Wobber, F. J. "Chemical contaminants on DOE lands and selection of contaminant mixtures for subsurface science research," U.S. Department of Energy, 1992. (2) Ginder-Vogel, M.; Borch, T.; Mayes, M. A.; Jardine, P. M.; Fendorf, S. Chromate reduction and retention processes within arid subsurface environments. Environ. Sci. Technol. 2005, 39, 7833-7839. (3) Morrison, S. J.; Metzler, D. R.; Carpenter, C. E. Uranium precipitation in a permeable reactive barrier by progressive irreversible dissolution of zerovalent iron. Environ. Sci. Technol. 2001, 35, 385-390. (4) Morrison, S. J.; Spangler, R. R.; Tripathi, V. S. Adsorption of uranium(VI) on amorphous ferric oxyhydroxide at high concentrations of dissolved carbon(IV) and sulfur(VI). Jour. Contam. Hydrol. 1995, 17, 333-346. (5) Grenthe, I.; Fuger, J.; Konings, R. J. M.; Lemire, R. J.; Muller, A. B.; Nguyen-Trung, C.; Wanner, H. Chemical thermodynamics of uranium; North-Holland Elsevier Science Publishers B.V.: Amsterdam, 1992; Vol. 1. (6) Langmuir, D. Uranium solution-mineral equilibria at low temperature with applications to sedimentary ore deposits. Geochim. Cosmochim. Acta 1978, 42, 547-569. (7) Sandino, A.; Bruno, J. The solubility of (UO2)3(PO4)2·4H2O(s) and the formation of U(VI) phospate complexes: Their influence in uranium speciation in natural waters. Geochim. Cosmochim. Acta 1992, 56, 4135-4145. (8) Moyes, L. N.; Parkman, R. H.; Charnock, J. M.; Vaughan, D. J.; Livens, F. R.; Hughes, C. R.; Braithwaite, A. Uranium Uptake from Aqueous Solution by Interaction with Goethite, Lepidocrocite, Muscovite, and Mackinawite: An X-ray Absorption Spectroscopy Study. Environ. Sci. Technol. 2000, 34, 1062-1068. (9) Bostick, B. B.; Fendorf, S.; Barnett, M. O.; Jardine, P. M.; Brooks, S. C. Uranyl Surface Complexes Formed on Subsurface Media from DOE Facilities. Soil Sci. Soc. Am. J. 2002, 66, 99-108. (10) Barnett, M. O.; Jardine, P. M.; Brooks, S. C.; Selim, H. M. Adsorption and transport of uranium(VI) in subsurface media. Soil Sci. Soc. Am. J. 2000, 64, 908-917. (11) Barnes, C. E.; Cochran, J. K. Uranium geochemistry in esturaine sediments: Controls on removal and release processes. Geochim. Cosmochim. Acta 1993, 57, 555-569. (12) Abdelouas, A.; Lutze, W.; Nutall, H. E. Oxidative dissolution of uraninite precipitated on Navajo sandstone. Journal of Contaminant Hydrology 1999, 36, 353-375. (13) Casas, I.; dePablo, J.; Gimenez, J.; Torrero, M. E.; Bruno, J.; Cera, E.; Finch, R. J.; Ewing, R. C. The role of pe ; pH ; and carbonate on the solubility of UO2 and uraninite under nominally reducing conditions. Geochim. Cosmochim. Acta 1998, 62, 2223-2231. (14) Wersin, P.; Hochella, M. F.; Persson, P.; Redden, G.; Leckie, J. O.; Harris, D. W. Interaction between aqueous uranium(VI) and sulfide minerals :

52

Spectroscopic evidence for sorption and reduction. Geochim. Cosmochim. Acta 1994, 58, 2829-2843. (15) Liger, E.; Charlet, L.; Cappellen, P. V. Surface catalysis of uranium(VI) reduction by iron(II). Geochim. Cosmochim. Acta 1999, 63, 2939-2955. (16) Fredrickson, J. K.; Zachara, J. M.; Kennedy, D. W.; Duff, M. C.; Gorby, Y. A.; Li, S. M. W.; Krupka, K. M. Reduction of U(VI) in goethite (alpha-FeOOH) suspensions by a dissimilatory metal-reducing bacterium. Geochim. Cosmochim. Acta 2000, 64, 3085-3098. (17) Lovley, D. R.; Phillips, E. J. P. Bioremediation of uranium contamination with enzymatic uranium reduction. Environ. Sci. Technol. 1992, 26, 2228-2234. (18) Gorby, Y. A.; Lovley, D. R. Enzymatic Uranium Precipitation. Environ. Sci. Technol. 1992, 26, 205-207. (19) Wu, W.-M.; Carley, J.; Gentry, T.; Ginder-Vogel, M.; Fienen, M.; Mehlhorn, T.; Yan, H.; Caroll, S.; Pace, M.; Nyman, J.; Luo, J.; Gentile, M.; Fields, M. W.; Hickey, R.; Watson, D. B.; Cirpka, O.; Zhou, J.; Fendorf, S.; Kitanidis, P.; Jardine, P. M.; Criddle, C. Pilot-scale in situ bioremediation of uranium in a highly contaminated aquifer. 2. Geochemical control of U(VI) bioavailability and evidence of U(VI) reduction. Environ. Sci. Technol. 2006, 40, 3986-3995. (20) Anderson, R. T.; Vrionis, H. A.; Ortiz-Bernard, I.; Resch, C. T.; Long, P. E.; Dayvault, R.; Karp, K.; Marutzky, S.; Metzler, D. R.; Peacock, A. D.; White, D. C.; Lowe, M.; Lovley, D. R. Stimulating the in situ activity of Geobacter species to remove uranium from the groundwater of a uranium-contaminated aquifer. Appl. Environ. Microb. 2003, 69, 5884-5891. (21) Fredrickson, J. K.; Zachara, J. M.; Kennedy, D. W.; Liu, C. G.; Duff, M. C.; Hunter, D. B.; Dohnalkova, A. Influence of Mn oxides on the reduction of uranium(VI) by the metal-reducing bacterium Shewanella putrefaciens. Geochim. Cosmochim. Acta 2002, 66, 3247-3262. (22) Senko, J. M.; Mohamed, Y.; Dewers, T.; Krumholz, L. R. Role for Fe(III) minerals in nitrate-dependent microbial U(IV) oxidation. Environ. Sci. Technol. 2005, 39, 2529-2536. (23) Finneran, K. T.; Housewright, M., E.; Lovley, D. R. Multiple influences of nitrate on uranium solubility during bioremediation of uranium-contaminated subsurface sediments. Environ. Microbiol. 2002, 4, 510-516. (24) Senko, J. M.; Istok, J. D.; Suflita, J. M.; Krumholz, L. R. In-situ evidence for uranium immobilization and remobilization. Environmental Science & Technology 2002, 36, 1491-1496. (25) Istok, J. D.; Senko, J. M.; Krumholz, L. R.; Watson, D.; Bogle, M.; Peacock, A. D.; Chang, Y.-J.; White, D. C. In situ bioreduction of technetium and uranium in a nitrate contaminated aquifer. Environ. Sci. Technol. 2004, 38, 468-475. (26) Sani, R. K.; Peyton, B. M.; Dohnalkova, A.; Amonette, J. E. Reoxidation of reduced uranium with iron(III) (hydr)oxides under sulfate-reducing conditions. Environ. Sci. Technol. 2005, 39, 2059-2066.

53

(27) Sani, R. K.; Peyton, B. M.; Amonette, J. E.; Geesey, G. G. Reduction of uranium(VI) under sulfate-reducing conditions in the presence of Fe(III)-(hydr)oxides. Geochim. Cosmochim. 2004, 68, 2639-2648. (28) Wan, J.; Tokunaga, T. K.; Brodie, E.; Wang, Z.; Zheng, Z.; Herman, D.; Hazen, T.; Firestone, M. K.; Sutton, S. R. Reoxidation of bioreduced uranium under reducing conditions. Environ. Sci. Technol. 2005, 39, 6162-6169. (29) Schwertmann, U.; Cornell, R. M. Iron oxides in the laboratory: Preparation and characterization; Wiley-VCH: Weinhein, 2000. (30) Hansel, C. M.; Benner, S. G.; Neiss, J.; Dohnalkova, A.; Kukkadapu, R. K.; Fendorf, S. Secondary mineralization pathways induced by dissimilatory iron reduction of ferrihydrite under advective flow. Geochim. Cosmochim. Acta 2003, 67, 2977-2992. (31) Casas, I.; Casabona, L.; Duro, L.; dePablo, J. The influence of hematite on the soprtion of uranium(VI) onto granite filling fractures. Chem. Geol. 1994, 113, 319-326. (32) Ticknor, K. V. Uranium sorption on geological materials. Radiochim. Acta 1994, 64, 229-236. (33) Stookey, L. L. A new spectrophotometric reagent for iron. Anal. Chem. 1970, 42, 779-781. (34) Bertsch, P. M.; Hunter, D. B. In Situ Chemical Speciation of Uranium in Soils and Sediments by Micro X-ray Absorption Spectroscopy. Environmental Science and Technology 1994, 28, 980-984. (35) Guillaumont, R.; Fanghanel, T.; Neck, V.; Fuger, J.; Palmer, D. A.; Grenthe, I.; Rand, M. H. "Update on the Chemical Thermodynamics of Uranium, Neptumium, Plutonium, Americium, and Technetium," Nuclear Energy Agency, 2003. (36) Suzuki, Y.; Kelly, S. D.; Kemner, K. M.; Banfield, J. F. Radionuclide contamination : Nanometer-size products of uranium bioreduction. Nature 2002, 419, 134-134. (37) Giammar, D. E.; Hering, J. G. Time scales for sorption-desorption and surface precipitation of uranyl on goethite. Environ. Sci. Technol. 2001, 35, 3332-3337. (38) Bernhard, G.; Geipel, G.; Reich, T.; V., B.; Amayri, S.; Nitsche, H. Uranyl(VI) carbonate complex formation: Validation of the Ca2UO2(CO3)3 species. Radiochim. Acta 2001, 89, 511-518. (39) Cornell, R. M.; Schwertmann, U. The Iron Oxides: Structure, Properties, Reactions, Occurrences and Uses; Wiley-VCH, 2003. (40) Majzlan, J.; Navrotsky, A.; Schwertmann, U. Thermodynamics of iron oxides: Part III. Enthalpies of formation and stability of ferrihydrite, schwertmannite, and ε-Fe2O3. Geochim. Cosmochim. Acta 2004, 68, 1049-1059. (41) Allison, J. D.; Brown, D. S.; Nova-Gradac, K. J. MINTEQA2/PRODEFA2, A geochemical assessment model for environmental systems: Version 3.0 user's manual; U.S. EPA: Athens, GA, 1990. (42) Senko, J. M.; Suflita, J. M.; Krumholz, L. R. Geochemical controls on microbial nitrate-dependent U(IV) oxidation. Geomicrobiology Journal 2005, 22, 371-378.

54

(43) Fredrickson, J.; Zachara, J.; Kennedy, D.; Dong, H.; Onstott, T.; Hinman, N.; Li, S. Biogenic iron mineralization accompanying the dissimilatory reduction of hydrous ferric oxide by a groundwater bacterium. Geochimica et Cosmochimica Acta 1998, 62, 3239-3257. (44) Hansel, C. M.; Benner, S. G.; Neiss, J.; Dohnalkova, A.; Kukkadapu, R. K.; Fendorf, S. Secondary mineralization pathways induced by dissimilatory iron reduction of ferrihydrite under advective flow. Geochimica et Cosmochimica Acta 2003, 67, 2977-2992. (45) Benner, S.; Hansel, C.; Wielinga, B.; Barber, T.; Fendorf, S. Reductive dissolution and biomineralization of iron hydroxide under dynamic flow conditions. Env. Sci. Technol. 2002, 36, 1705-1711. (46) Zhou, P.; Gu, B. Extraction of oxidized and reduced forms of uranium from contaminated soils: Effects of carbonate concentration and pH. Environ. Sci. Technol. 2005, 39, 4435-4440.

55

Table 2.1 Gibb’s Free Energy of Reactions at Standard State conditions (∆G°r), or at experimental conditions represented by pH 7, 1 x 10-6 M U(VI) species, 5 x 10-7 M Fe2+, and 3 x 10-3 M HCO3

- (∆G°*r). Sources of data are provided within Table 2A.1.

Oxidation Reaction ∆G°r (kJ

mol-1)

∆G°*r (kJ

mol-1)

Fe(OH)3 + 0.5 UO2 + 3 H+ ↔ Fe2+ + 0.5 UO22+ + 3H2O -60.08 6.6

Fe(OH)3 + 0.5 UO2 + 2.5 H+ + 0.5 HCO3- ↔ 0.5 UO2CO3 + Fe2+ + 3H2O -58.97 -5.0

Fe(OH)3 + 0.5 UO2 + 2 H+ + HCO3- ↔ 0.5 UO2(CO3)2

2- + Fe2+ + 3H2O -48.53 -7.3

Fe(OH)3 + 0.5 UO2 + 1.5 H+ + 1.5 HCO3- ↔ 0.5 UO2(CO3)3

4- + Fe2+ + 3 H2O -33.98 -5.5

Fe(OH)3 + 0.5 UO2 + 1.5 H+ + 1.5 HCO3- + Ca2+ ↔ 0.5 Ca2UO2(CO3)3 + Fe2+ + 3 H2O -58.40 -12.8

Fe(OH)3 + 0.5 UO2 + 1.5 H+ + 1.5 HCO3- + 0.5 Ca2+ ↔ 0.5 CaUO2(CO3)3

2- + Fe2+ + 3 H2O -44.02 -7.0

Fe(OH)3 + 0.5 UO2 + 2 H+ + 0.25 HCO3- ↔ 0.25 (UO2)2CO3(OH)3

- + Fe2+ + 2.25 H2O -44.14 -5.2

FeOOH + 0.5 UO2 + 3 H+ ↔ 0.5 UO22+ + Fe2+ + 2 H2O -48.58 18.2

FeOOH + 0.5 UO2 + 2.5 H+ + 0.5 HCO3- ↔ 0.5 UO2CO3 + Fe2+ + 2 H2O -47.48 6.5

FeOOH + 0.5 UO2 + 2 H+ + 1 HCO3- ↔ 0.5 UO2(CO3)2

2- + Fe2+ + 2H2O -37.03 4.2

FeOOH + 0.5 UO2 + 1.5 H+ + 1.5 HCO3- ↔ 0.5 UO2(CO3)3

4- + Fe2+ + H2O -22.48 5.9

FeOOH + 0.5 UO2 + 1.5 H+ + 1.5 HCO3- + Ca2+ ↔ 0.5 Ca2UO2(CO3)3 + Fe2+ + 2 H2O -46.90 -1.4

FeOOH + 0.5 UO2 + 1.5 H+ + 1.5 HCO3- + 0.5 Ca2+ ↔ 0.5 CaUO2(CO3)3

2- + Fe2+ + 2 H2O -32.52 4.5

FeOOH + 0.5 UO2 + 2 H+ +0.25 HCO3- ↔ 0.25 (UO2)2CO3(OH)3

- + Fe2+ + 5/4 H2O -32.64 6.3

0.5 Fe2O3 + 0.5 UO2 + 3 H+ ↔ Fe2+ + 0.5 UO2

2+ + 1.5 H2O -39.82 26.9

0.5 Fe2O3 + 0.5 UO2 + 2.5 H+ + 0.5 HCO3- ↔ Fe2+ + 1.5 H2O + 0.5 UO2CO3 -38.72 15.2

0.5 Fe2O3 + 0.5 UO2 + 2 H+ + 1 HCO3- ↔ Fe2+ + 1.5 H2O + 0.5 UO2(CO3)2

2- -28.28 12.9

0.5 Fe2O3 + 0.5 UO2 + 1.5 H+ + 1.5 HCO3- ↔ Fe2+ + 1.5 H2O + 0.5 UO2(CO3)3

4- -13.73 14.7

0.5 Fe2O3 + 0.5 UO2 + 1.5 H+ +1.5 HCO3- + Ca2+ ↔ Fe2+ + 1.5 H2O + 0.5 Ca2UO2(CO3)3 -38.15 7.4

0.5 Fe2O3 + 0.5 UO2 + 1.5 H+ +1.5 HCO3- + 0.5 Ca2+ ↔ Fe2+ + 1.5 H2O + 0.5 CaUO2(CO3)3

2- -23.77 13.2

0.5 Fe2O3 + 0.5 UO2 + 2 H+ + 0.25 HCO3- + 0.25 (UO2)2CO3(OH)3

- + Fe2+ + 0.75 H2O -23.89 15.1

56

Table 2.2 Geochemical conditions at the conclusion of each UO2 oxidation experiment. ∆G reaction values were calculated using extractable Fe(II) and soluble U(VI) concentrations, along with the reactions and thermodynamic values listed in Table 2.1.

Iron

Oxide [HCO3

-] (mM)

Fe(II)(aq) (µM)

Extractable Fe(II) (µM)

Soluble U(VI) (µΜ)

Soluble U(VI) (µM)

Control

Predominant U(VI) Species in

∆G reaction

(kJ/mole) Ferrihydri

te 3 0.47 23.3 0.70 0.60 9.3 UO2(CO3)2

2- 3.4 % UO2(CO3)3

4- 1.6 % UO2CO3(aq)

85.5 (UO2)2CO3(OH)3-1

0.5 2.2 2.8 3.4

Ferrihydrite

10 0.52 38.8 0.93 0.77 24.3% UO2(CO3)22-

55.0% UO2(CO3)34-

18.2% (UO2)2CO3(OH)3-1

-1.0 -0.1 3.6

Ferrihydrite

50 1.22 104.4 3.76 1.98 2.5% UO2(CO3)22-

97.2% UO2(CO3)34-

-3.9 -1.2

Ferrihydrite

100 1.66 120.6 4.88 2.57 99.2% UO2(CO3)34- -3.4

Goethite 3 ND1 2.8 0.50 0.60 11.6% UO2(CO)22-

4.4 % UO2(CO3)34-

2.0% UO2CO3 81.9% (UO2)2CO3(OH)

3-1

4.9 5.5 5.0 10.0

Hematite 3 ND ND 0.38 0.60 66.9% UO2(CO)22-

20.3 % UO2(CO3)34-

11.0% UO2CO3 1.6% (UO2)2CO3(OH)3

-1

NA2

1. Not detected 2. Not applicable

57

Figure Captions Figure 2.1 Representative Fe(III)/Fe(II) and U(VI)/U(IV) redox couples at pH 7;

concentrations of 3 x 10-3 M HCO3-, 1 x 10-6 M U(VI), 1 x 10-3 M Ca2+, and either 5 x

10-7 or 1 x 10-5 M Fe(II) are portrayed.

Figure 2.2 Free energy of reaction for UO2 (biogenic) oxidation by ferrihydrite (A),

goethite (B) and hematite (C) as a function of pH with 3 x 10-3 M HCO3-, 1 x 10-6 M

U(VI), 1 x 10-3 M Ca2+, and 1 x 10-5 M Fe(II).

Figure 2.3 Free energy of reaction for the oxidation of UO2 (biogenic) by ferrihydrite as

a function of HCO3- concentration at pH 7, 1 x 10-6 M U(VI), 1 x 10-3 M Ca2+, and

1 x 10-5 M Fe(II).

Figure 2.4 Effect of carbonate concentration on the oxidation of UO2 (biogenic) by

ferrihydrite.

Figure 2.5 Thermodynamic viability of UO2 (biogenic) oxidation by ferrihydrite for

conditions reported for various field sites and soil-column experiments in the absence

(A) and presence (B) of 1 mM Ca2+ (numbers indicate reference for data point).

Conditions representing equilibrium with 0.126 uM U(VI) (aq) (the drinking water

maximum) are illustrated; for specific Fe(II) concentrations, reoxidation is viable at

bicarbonate and pH conditions plotting above the line and non-viable for those

plotting below. Colors are used to denote the potential for probable Fe(III) (hydr)oxide

oxidation of UO2 reflected by typical Fe(II) concentrations for subsurface

environments. Relevant geochemical conditions and references are detailed in Table

2A.2.

58

Figure 2.1 Representative Fe(III)/Fe(II) and U(VI)/U(IV) redox couples at pH 7; concentrations of 3 x 10-3 M HCO3

-, 1 x 10-6 M U(VI), 1 x 10-3 M Ca2+, and either 5 x 10-7 or 1 x 10-5 M Fe(II) are portrayed.

59

pH

2 3 4 5 6 7 8 9 10 11

Gib

b's

Free

Ene

rgy

of R

eact

ion

(kJ/

mol

)

-80

-60

-40

-20

0

20

40

60

80

100UO2CO3

UO2(CO3)22-

UO2(CO3)34-

Ca2UO2(CO3)3

CaUO2(CO3)22-

UO22+

(UO2)2CO3(OH)3-

A - Ferrihydrite

UO2 Oxidation Favored

UO2 Oxidation not Favored

pH

2 3 4 5 6 7 8 9 10 11

Gib

bs F

ree

Ener

gy o

f Rea

ctio

n (k

J/m

ol)

-60

-40

-20

0

20

40

60

80

100UO2CO3

UO2(CO3)22-

UO2(CO3)34-

Ca2UO2(CO3)3

CaUO2(CO3)22-

UO22+

(UO2)2CO3(OH)3-

B - Goethite



60

pH

2 3 4 5 6 7 8 9 10 11

Gib

bs F

ree

Ener

gy o

f Rea

ctio

n (k

J/m

ol)

-60

-40

-20

0

20

40

60

80

100

UO2CO3

UO2(CO3)22-

UO2(CO3)34-

Ca2UO2(CO3)3

CaUO2(CO3)22-

UO22+

(UO2)2CO3(OH)3-

C - Hematite



Figure 2.2 Free energy of reaction for UO2 (biogenic) oxidation by ferrihydrite (A), goethite (B) and hematite (C) as a function of pH with 3 x 10-3 M HCO3

-, 1 x 10-6 M U(VI), 1 x 10-3 M Ca2+, and 1 x 10-5 M Fe(II).

61

HCO3- (mM)

0 20 40 60 80 100

Gib

bs F

ree

Ener

gy o

f Rea

ctio

n (k

J/m

ole)

-20

-10

0

10

20

UO2CO3

(UO2)2CO3(OH)3-

UO2(CO3)22-

UO2(CO3)34-

Ca2UO2(CO3)3

CaUO2(CO3)32-

UO22+



Figure 2.3 Free energy of reaction for the oxidation of UO2 (biogenic) by ferrihydrite as a function of HCO3

- concentration at pH 7, 1 x 10-6 M U(VI), 1 x 10-3 M Ca2+, and 1 x 10-5 M Fe(II).

62

HCO3- (mM)0 20 40 60 80 100 120

Sol

uble

U(V

I) (µ

M)

0

1

2

3

4

5

6

Experimental

Controlwithout Ferrihydrite

HCO3- (mM)

0 20 40 60 80 100 120

Ext

ract

able

Fe(

II) (µ

M)

0

20

40

60

80

100

120

140

160

Experimental

Controlwithout UO2

Figure 2.4 Effect of carbonate concentration on the oxidation of UO2 (biogenic) by ferrihydrite.

63

Figure 2.5 Thermodynamic viability of UO2 (biogenic) oxidation by ferrihydrite for conditions reported for various field sites and soil-column experiments in the absence (A) and presence (B) of 1 mM Ca2+ (numbers indicate reference for data point). Conditions representing equilibrium with 0.126 uM U(VI) (aq) (the drinking water maximum) are illustrated; for specific Fe(II) concentrations, reoxidation is viable at bicarbonate and pH conditions plotting above the line and non-viable for those plotting below. Colors are used to denote the potential for probable Fe(III) (hydr)oxide oxidation of UO2 reflected by typical Fe(II) concentrations for subsurface environments. Relevant geochemical conditions and references are detailed in Table 2A.2.

64

APPENDIX 2A


65

Table 2A.1 Gibb’s Free Energy of formation for species used in thermodynamic calculations.

Species ∆Gf° (kJ/mol) Source 2-line Fe(OH)3 -708.5 (1) α-FeOOH -482.9 (2)

Fe2O3 -746.2 (2) UO2 (am) -995.8 (3) UO2 (c) -1031.8 (3)

U(OH)4 (aq) -1421.3 (3) U(CO3)4

4- -2841.9 (3) UO2

2+ (aq) -952.5 (3) UO2CO3(aq) -1537.2 (3)

UO2(CO3)22- (aq) -2103.2 (3)

UO2(CO3)34- (aq) -2660.9 (3)

Ca2UO2(CO3)3 (aq) -3815.4 (4) CaUO2(CO3)3

2- (aq) -3233.8 (4) Ca2+ (aq) -552.8 (3)

HCO3- (aq) -586.8 (3)

Fe2+ (aq) -78.9 (3) H2O (l) -237.1 (3)

66

Table 2A.2 References and relevant geochemical conditions used in Figure 2.5.

Point in Figure [HCO3-]

(mM) pH [Fe(II)]

(µM) [U(VI)] (µM)

Source Notes

19-1 0.6 6.25 NA1 2.5 Wu et al. (5)

Day 399 – Prior to EtOH addition

19-2 1.2 6.6 NA1 1.5 Wu et al. (5)

Day 401 – Beginning of EtOH addition

19-3 0.9 6.8 NA1 2.9 Wu et al. (5)

Day 404 – Conclusion of EtOH addition

25 100.02 6.8 0-500 1-5 Istok et al. (6)

Test 40 Well FW034

20 3.83 7 100 0.2 - 1 Anderson et al. (7)

Well M-13

22 40.0 6.8 0 ~ 100 Senko et al. (8)

Soluble Fe(II) is negligible after NO2

- addition

29 30.0 7.0 NA1 60 Sani et al. (9)

Soluble U(VI) at 20 Figure 3

27-1 9.0 7.3 NA1 0.1 Wan et al. (10)

Day 107 of Figure 1

27-2 12.7 7.4 .042 1.4 Wan et al. (10)

Day 346 of Figure 1 – Steady State

1. Not Available 3. Istok et al. did not report the specific [HCO3

-] used in each study, but gave a range of 80-130 mM for [HCO3

-]; therefore, we have chosen 100 mM as the representative [HCO3

-]. 2. Anderson et al. did not report [HCO3

-], because this is an arid environment and the study used amended native groundwater we have assumed [HCO3

-] was controlled by equilibrium with calcite at pH 7.

67

pH

2 3 4 5 6 7 8 9 10 11

% U

(VI)

0

20

40

60

80

100

120

UO2CO3(OH)3-

UO2(CO3)22-

UO2(CO3)34-

UO2CO3

UO22+

UO2OH+

A

pH

2 3 4 5 6 7 8 9 10 11

% U

(VI)

0

20

40

60

80

100

120

Ca2UO2(CO3)2 UO2

2+

UO2(CO3)22-

UO2CO3

B

68

HCO3- (mM)

0 10 20 30 40 50 60 70 80 90 100

% U

(VI)

0

20

40

60

80

100

120

UO2(CO3)34-

(UO2)2CO3(OH)3-

UO2(CO3)22-

UO2CO3

C

HCO3- (mM)

0 20 40 60 80 100

% U

(VI)

0

20

40

60

80

100

120

UO2(CO3)22-

UO2(CO3)34-

Ca2UO2(CO3)3

CaUO2(CO3)22-

D

Figure 2A.1 Soluble uranium speciation as a function of pH with 1 x 10-6 M of each U(VI) species, and 1 x 10 -3 M HCO3

- in the absence (A) or presence (B) of 1 x 10-3 M Ca2+. The influence of carbonate concentration on uranyl speciation at pH 7 with 1 x 10-6 M of each U(VI) species in the absence (C) or presence (D) of 1 x 10-3 M Ca2+.

69

Ext

ract

able

Fe(

II) (µ

Μ)

0

10

20

30

40

23.3

Sol

uble

U(V

I) (µ

M)

0.0

0.2

0.4

0.6

0.8

Ferrihydrite Goethite Hematite Control

Ferrihydrite Goethite Hematite Control

3.6

ND ND

A

0.70

0.50

0.38

0.60

B

Figure 2A.2 Extractable Fe(II) (A) and soluble U(VI) (B) at the conclusion of the oxidation reactions.

70

Figure 2A.3 Representative Fe(III)/Fe(II) and soluble U(VI)/U(IV) redox couples at pH 7 with 3 x 10-3 M HCO3

-, 1 x 10-6 M each U(VI) species, 1 x 10-9 M each U(IV) species, 1 x 10-3 M Ca2+, and either 5 x 10-7 or 1 x 10-5 M Fe2+.

71

(1) Majzlan, J.; Navrotsky, A.; Schwertmann, U. Thermodynamics of iron oxides: Part III. Enthalpies of formation and stability of ferrihydrite, schwertmannite, and ε-Fe2O3. Geochim. Cosmochim. Acta 2004, 68, 1049-1059. (2) Cornell, R. M.; Schwertmann, U. The Iron Oxides: Structure, Properties, Reactions, Occurrences and Uses; Wiley-VCH, 2003. (3) Guillaumont, R.; Fanghanel, T.; Neck, V.; Fuger, J.; Palmer, D. A.; Grenthe, I.; Rand, M. H. "Update on the Chemical Thermodynamics of Uranium, Neptumium, Plutonium, Americium, and Technetium," Nuclear Energy Agency, 2003. (4) Bernhard, G.; Geipel, G.; Reich, T.; V., B.; Amayri, S.; Nitsche, H. Uranyl(VI) carbonate complex formation: Validation of the Ca2UO2(CO3)3 species. Radiochim. Acta 2001, 89, 511-518. (5) Wu, W.-M.; Carley, J.; Gentry, T.; Ginder-Vogel, M.; Fienen, M.; Mehlhorn, T.; Yan, H.; Caroll, S.; Pace, M.; Nyman, J.; Luo, J.; Gentile, M.; Fields, M. W.; Hickey, R.; Watson, D. B.; Cirpka, O.; Zhou, J.; Fendorf, S.; Kitanidis, P.; Jardine, P. M.; Criddle, C. Pilot-scale in situ bioremediation of uranium in a highly contaminated aquifer. 2. Geochemical control of U(VI) bioavailability and evidence of U(VI) reduction. Environ. Sci. Technol. 2006, 40, 3986-3995. (6) Istok, J. D.; Senko, J. M.; Krumholz, L. R.; Watson, D.; Bogle, M.; Peacock, A. D.; Chang, Y.-J.; White, D. C. In situ bioreduction of technetium and uranium in a nitrate contaminated aquifer. Environ. Sci. Technol. 2004, 38, 468-475. (7) Anderson, R. T.; Vrionis, H. A.; Ortiz-Bernard, I.; Resch, C. T.; Long, P. E.; Dayvault, R.; Karp, K.; Marutzky, S.; Metzler, D. R.; Peacock, A. D.; White, D. C.; Lowe, M.; Lovley, D. R. Stimulating the in situ activity of Geobacter species to remove uranium from the groundwater of a uranium-contaminated aquifer. Appl. Environ. Microb. 2003, 69, 5884-5891. (8) Senko, J. M.; Mohamed, Y.; Dewers, T.; Krumholz, L. R. Role for Fe(III) minerals in nitrate-dependent microbial U(IV) oxidation. Environ. Sci. Technol. 2005, 39, 2529-2536. (9) Sani, R. K.; Peyton, B. M.; Dohnalkova, A.; Amonette, J. E. Reoxidation of reduced uranium with iron(III) (hydr)oxides under sulfate-reducing conditions. Environ. Sci. Technol. 2005, 39, 2059-2066. (10) Wan, J.; Tokunaga, T. K.; Brodie, E.; Wang, Z.; Zheng, Z.; Herman, D.; Hazen, T.; Firestone, M. K.; Sutton, S. R. Reoxidation of bioreduced uranium under reducing conditions. Environ. Sci. Technol. 2005, 39, 6162-6169.

72

CHAPTER 3

Kinetic and Mineralogical Constraints on the Oxidation of Biogenic Uraninite by Ferrihydrite

Matthew Ginder-Vogel and Scott Fendorf

Department of Geological and Environmental Sciences Stanford University, Stanford, CA 94305

In review for: Kinetic and mineralogical constrains on the oxidation of biogenic uraninite by ferrihydrite. Ginder-Vogel, M.; Fendorf, S. Environ. Sci. Technol.

73

ABSTRACT

The oxidation state of uranium plays a major role in determining uranium

mobility in the environment. Uranium is generally found in the hexavalent form, is

quite soluble, and readily forms complexes with calcium and carbonate in aerobic

environments. However, under anaerobic conditions, common metal respiring bacteria

can enzymatically reduce U(VI) to U(IV), resulting in the formation of sparingly

soluble UO2 (uraninite). Uranium(VI) reduction, therefore, has a prominent role in

uranium natural attenuation and is being explored as a potential uranium remediation

technique. The stability of biologically precipitated uraninite is critical for determining

the long-term fate of uranium and is not well characterized within soils and sediments.

Here, we demonstrate that uraninite oxidation by ferrihydrite, a disordered Fe(III)

(hydr)oxide, proceeds through a soluble U(IV) intermediate and results in the

concomitant production of Fe(II) and dissolved U(VI). Uraninite oxidation rates are

accelerated under conditions that increase its solubility, which include high

bicarbonate concentration and pH values deviating from neutrality. Additionally,

Fe(II) produced during uranium oxidation catalyzes the transformation of ferrihydrite

into goethite and lepidocrocite, which, combined with elevated Fe(II)(aq) and U(VI)(aq),

may ultimately limit UO2 oxidation over longer time frames. Thus, our results

demonstrate that UO2 oxidation by Fe(III) (hydr)oxides is controlled by the rate of

uraninite dissolution and that this process may limit uranium sequestration under

mildly reducing conditions.

74

3.1 INTRODUCTION

Uranium contamination of ground and surface waters has been detected at

numerous sites throughout the world, including agricultural evaporation ponds (1),

U.S. Department of Energy nuclear weapons manufacturing areas, and mine tailings

sites (2). Uranium is generally found in the hexavalent oxidation state in oxic systems.

While U(VI) forms several sparingly soluble complexes with phosphate (3,4), and

readily complexes with many transition metal (hydr)oxides (5-8), it commonly

remains relatively soluble, and thus mobile, due to the formation of stable aqueous

complexes with carbonate and calcium (9). Conversely, U(IV) forms sparingly

soluble solids, even in the presence of common groundwater ligands, such as

carbonate, and thus tends to be relatively immobile. Therefore, the oxidation state of

uranium will play an important role in determining its environmental mobility.

Abiotic (chemical) reduction of U(VI) is essentially limited to aqueous sulfide

(at low pH and low bicarbonate concentration) (10), sulfide minerals (11), and surface

bound Fe(II) (12). However, numerous, common dissimilatory metal (DMRB) and

sulfate reducing bacteria (SRB), including Shewanella, Geobacter, and Desulfovibrio

species, couple the oxidation of organic matter and H2 to the reduction of U(VI),

resulting in U(IV) and the subsequent precipitation of uraninite (UO2) (13-15), a

sparingly soluble phase. Therefore, biological uranium reduction plays an important

role in the cycling and natural attenuation of uranium. Additionally, this process is

being explored as an in situ remediation technique (16-18).

In order to determine the viability of in situ biological uranium remediation

and discern the role of reductive processes in natural uranium cycling, it is critical to

75

determine the stability of biologically precipitated uraninite. Microcrystalline

uraninite has an estimated solubility of 10-8 to 10-12 at pH 7 (19-21); however, it can be

oxidized and remobilized by several common environmental constituents. Molecular

oxygen and Mn oxides rapidly and extensively oxidize UO2 (22). Nitrate, a common

co-contaminant with uranium (2), not only impedes biological uranium reduction (23-

25), but it also induces U(IV) oxidation through the production of reactive

intermediates (i.e. NO2-, NO, and N2O) in denitrification (26) and through direct

respiration (27).

Recent work has confirmed the role of Fe(III) (hydr)oxide minerals in the

oxidation of biogenic UO2 (reaction 1) (26,28-33), which reflects the overall

oxidation reaction by ferrihydrite.

UO2(biogenic) + 2Fe(OH)3 + 2HCO3- + 4H+ → UO2(CO3)2

2-(aq) + 2Fe2+

(aq) + 6H2O

(1)

The redox couples for U(IV)/U(VI) and Fe(III) (hydr)oxide/Fe(II) occur at similar

potentials under common groundwater conditions (Figure 3A.1), and, therefore, small

changes in aqueous chemistry can result in UO2 oxidation by Fe(III) (hydr)oxide

oscillating between thermodynamic viability and non-viability (33). Oxidation of

biogenic U(IV) and concomitant release of Fe(II) was observed in the presence of

Fe(III) (hydr)oxide minerals produced by nitrate dependent oxidation of Fe(II) (26).

Iron(III) (hydr)oxide minerals have also been implicated in the oxidation of U(IV)

under sulfate reducing conditions. The rate and mechanism of biogenic uraninite

oxidation by Fe(III) (hydr)oxide minerals remains unresolved. Accordingly, the

objective of this study is to evaluate the oxidation of biogenic uraninite by Fe(III)

76

(hydr)oxides. Here we examine the kinetics of UO2 oxidation by ferrihydrite, a

common Fe(III) (hydr)oxide found in soils and sediments (34), in order to determine

the rate and rate-controlling parameters. Additionally, we characterize the speciation

of the iron (hydr)oxide solid-phases formed during biogenic uraninite oxidation in

order to determine the influence of iron mineralogy on the extent of uraninite

oxidation.

77


3.2.1 Chemical Synthesis Procedures

All chemicals were ACS reagent grade or better and used without further

purification. Solutions used in this study were prepared using distilled deionized water

that had been treated to remove dissolved O2 by boiling while purging with nitrogen

gas that had been passed over hot Cu-metal filings. The water was cooled for 12 h

while being purged with N2 and immediately transferred to an anaerobic glovebox

(Coy Laboratory Products) with a 95% N2 and 5% H2 atmosphere. All glassware and

equipment was equilibrated in the anaerobic chamber for 24 h prior to use. Biogenic

uraninite was prepared using the method described previously (33). X-ray diffraction

patterns were identical to those previously reported for biogenic UO2 (13,33). The

U(VI) content of the biogenic uraninite was 5% as determined by extraction with 100

mM potassium bicarbonate. Two-line ferrihydrite was synthesized by rapidly titrating

a ferric chloride solution with NaOH to a pH of 7.5 (35). The ferrihydrite floc was

washed 4 times by centrifugation with deionized distilled water, then resuspended and

degassed by bubbling with N2 for 24 h. The N2 BET surface area of ferrihydrite was

210 m2 g-1 (36) and for uraninite was 129 m2 g-1 (37).

3.2.2 Oxidation Experiments

Ferrihydrite and uraninite were maintained as aqueous suspensions to avoid

diminishing their reactivity and were added to each oxidation reaction as a slurry.

Biogenic uraninite was used within two weeks and ferrihydrite was used within two

days of preparation. Changes in reactivity of the two solid phases were not observed

over this time frame. Two sets of oxidation experiments were carried out at room

78

temperature, one to determine the rate of UO2 oxidation by ferrihydrite, and the

second to examine solid-phase Fe and U speciation during UO2 oxidation. Reactions

to examine the rate of UO2 oxidation were carried out in 125 mL sealed polypropylene

Nalgene bottles, using 100 mL of the appropriate media; reactions to examine solid-

phase Fe and U speciation were carried out in 1 L of media in 1 L Nalgene bottles.

All reactions were continuously stirred using an overhead stirrer at 75 rpm to avoid

solid-phase abrasion, and, unless otherwise noted, were performed in 3 mM KHCO3 at

pH 7.2. Kinetic experiments run for 10 min were performed in triplicate, while

singular experiments were performed for solid-phase characterization. Except for

reactions to test the effect of pre-equilibration on reaction rate, ferrihydrite was added

to the reaction media first, followed by the uraninite, in less than 10 s. Initial and final

pH of all reactions varied by less than 0.1 unit. The uraninite and ferrihydrite

concentrations used in each experiment were determined by acidic dissolution of the

reaction slurry and are denoted as m2 L-1. In this system 1 mM-Fe as ferrihydrite is

equivalent to 55.7 mg (11.7 m2) ferrihydrite L-1, while 1 mM-U as UO2 is equivalent to

238 mg (30.7 m2) uraninite L-1.

To determine the effect of pH, bicarbonate, ferrihydrite, and biogenic uraninite

concentration on oxidation rate, each component was systematically varied while all

other reaction conditions were held constant (Table 3.1). Three experiments to

examine solid phase U and Fe evolution during uraninite oxidation were performed

with a single uraninite concentration of 30.7 m2 L-1 and ferrihydrite concentrations of

23.5, 46.0, and 80.2 m2 L-1. Solid-phase samples were collected by vacuum filtration

79

of 50 mL of the reaction slurry, which was then dried with 20 mL of ethanol in order

to halt uraninite oxidation.

3.2.3 Sampling and Analytical Procedures

The rate of biogenic UO2 oxidation by ferrihydrite was quantified primarily

using ferrozine extractable Fe(II) concentrations (38). The system partitioning

coefficient for Fe(II) was used to determine total Fe(II), which was in excellent

agreement with stoichiometric amounts of U(VI) produced (Figure 3A.2).

Measurement of dissolved U(VI) was considered less reliable because of the potential

for differing concentrations of U(VI) to sorb on Fe(III) (hydr)oxides under the varying

reaction conditions examined. Ferrozine-extractable Fe(II) was used rather than

soluble or acid-extractable because (i) of the propensity for Fe(II) uptake by Fe(III)

(hydr)oxides (5,39-41) and (ii) UO2(biogenic) oxidation by Fe(III) is more favorable (33)

and more rapid (vida infra) under acidic conditions. In fact, after 24 h, we observe

near complete oxidation of 30.7 m2 L-1 biogenic uraninite by 32.1 m2 L-1 ferrihydrite

in 0.5 M HCl. Extractable Fe(II) was determined by adding 1.5 mL of reaction slurry

to 1.5 mL of ferrozine reagent and reacting for 20 s; the sample was then passed

through a 0.2 µm polycarbonate filter and Fe(II) quantified by the absorbance at 562

nm. Prior to reaction, neither uraninite nor ferrihydrite slurries contained detectable

quantities of Fe(II) as measured by this technique. The observed rate coefficient for

each set of reaction conditions was calculated by linear regression of 10 extractable-

Fe(II) data points (R2 of all regression lines is > 0.90).

Total iron and uranium were determined at the conclusion of each oxidation

experiment by acidic dissolution of the reaction slurry with concentrated HNO3 and

80

HCl and quantified with inductively coupled plasma – optical emission spectrometry

(ICP-OES). Total dissolved iron and uranium were determined by passing 5 mL of

reaction slurry through a 0.2 µm polycarbonate filter, which was then acidified and

also measured by ICP-OES. Soluble Fe(II) in the filtrate was measured by the

ferrozine method, while soluble U(VI) was measured spectroflourometically.

Uranium samples were diluted 1:30 in 10% phosphoric acid, and the fluorescence of

the uranyl-phosphate complex was measured at 515.4 nm. All measurements were

referenced to the fluorescence of the background matrix.

3.2.4 X-ray Diffraction and Absorption Spectroscopies

X-ray diffraction (XRD) was used to identify crystalline iron phases after

uraninite oxidation. XRD patterns were collected on beamline 11-3 of the Stanford

Synchrotron Radiation Laboratory (SSRL) in transmission geometry, using

monochromatic radiation (12,732.137 eV) and a MAR 345 image plate. The resulting

images were processed using FIT2D (42). The sample-to-detector distance and

geometric corrections were calculated from the pattern of LaB6. After these

corrections were applied, the 2D images were integrated radially to yield 1D powder

diffraction patterns which could then be analyzed using standard techniques. Peak

identification and background correction, including removal of the scattering from the

lexan window, were performed in JADE 6.5 (Materials Data, Inc., Livermore, CA).

Samples were mounted in the anaerobic chamber between Lexan windows sealed with

double-sided tape to limit sample oxidation during analysis.

X-ray absorption near-edge structure (XANES) spectroscopy was used to

determine the relative uranium oxidation state, while extended X-ray absorption fine

81

structure (EXAFS) spectroscopy was used to quantify the iron solid-phase distribution.

Samples were powdered using a mortar and pestle, diluted with boron nitride,

mounted on a Teflon plate, and sealed with Kapton polyamide film in an anaerobic

glovebox to prevent sample oxidation while minimizing X-ray absorption.

Fluorescence data were collected at SSRL beamline 11-2, using either a 30-element

Ge semiconductor detector (U) or a Lytle fluorescence chamber (Fe). Incident and

transmitted X-ray intensities were measured with in-line ionization chambers. The

energy range studied was -200 to +800 eV around the U LIIIα-edge of U (17,166 eV)

and -200 to +1,000 eV around the Fe K-edge (7,112 eV). All samples were internally

referenced to either a U(VI) nitrate or Fe-metal standard, placed between the second

and third in-line ionization chambers. Two to four individual spectra were averaged

for each sample.

XANES and EXAFS spectra were processed using the SixPACK (43) interface

to IFEFFIT (44). XANES data were background-subtracted and normalized to a unit-

edge step. The relative amount of reduced uranium (± 10%) in each sample was

determined by comparison of the half-height edge position of each sample to a

standard curve obtained from samples with varying known mole ratios of U(VI)/U(VI)

(45). After background subtraction and normalization, EXAFS data were extracted

and k3-weighted. A set of reference standards for Fe was utilized to perform linear

combination k3-weighted EXAFS spectral fitting, using SixPACK’s least-squares

fitting module, which is a graphical interface to IFEFFIT’s minimization function

(44). Linear combination fitting routines were used to reconstruct the experimental

spectrum and to determine the relative percentages of iron mineral phases. Each

82

spectra was fit using 2-line ferrihydrite, lepidocrocite, and goethite, which were

detected in the X-ray diffraction patterns for the 23.5 m2 L-1 ferrihydrite experiment

(Figure 3A.3).

83

3.3 RESULTS

3.3.1 Reaction Kinetics

The oxidation of biogenic uraninite by ferrihydrite (reaction 1) was determined

as a function of ferrihydrite and uraninite suspension densities, pH, and bicarbonate

concentrations. The concentration of extractable Fe(II) and dissolved U(VI) increased

during the reaction period under all conditions (Supporting Information). The rate of

uraninite consumption is calculated as the rate of ferrozine extractable Fe(II)

production (supporting information), and conforms to a pseudo first-order rate

expression for all conditions tested in this study (Figure 3.1). The pseudo first-order

rate constants increase with increasing uraninite suspension density (Table 3.1, Figure

3.2A). At ferrihydrite concentrations above 7.1 m2 L-1, the oxidation rate of uraninite

is independent of ferrihydrite concentration; however, below 7.1 m2 L-1 ferrihydrite,

uraninite oxidation rates decrease with decreasing ferrihydrite concentration (Table

3.1, Figure 3.2B).

In addition to the concentration of the reactants, the rate of UO2 oxidation is

also impacted by changes in pH and bicarbonate concentration. Reaction rate varies

non-linearly with pH and reaches a minimum at pH 7.2. Below pH 7.2 the reaction

rate rapidly increases, and above pH 7.2 the reaction rate increases more modestly

(Table 3.1, Figure 3.2C). Increasing bicarbonate concentration also results in a slight

increase in biogenic UO2 oxidation rate (Table 3.1, Figure 3.2D).

Equilibration of the ferrihydrite with the reaction media for 10 minutes prior to

reaction does not result in a change of the reaction rate (Table 3.1, Figure 3.2B).

However, uraninite equilibration with the reaction media for 10 minutes results in 2

84

µM of dissolved uranium of which 7 nM is U(IV), consistent with the predicted

solubility of microcrystalline uraninite and the U(VI) concentration of the biogenic

uraninite used in the study. Upon the addition of ferrihydrite, not only does the

extractable Fe(II) at t = 1 min increase (Figure 3A.4), the rate of oxidation also

increases from 8.3 x 10-12 mol s-1 without equilibration to 1.0 x 10-11 mol s-1 after

equilibration, this change is larger than the uncertainty in the slope of the linear

regression (Table 3.1, Figure 3.1A).

The rate dependence of the oxidation reaction on uraninite and bicarbonate

concentration can be determined by considering the conditional rate expression for the

pseudo-first-order rate coefficient.

kobs (s-1) = k’species[reactant]x (2)

Regressing the logarithm of the observed (pseudo-first-order) rate coefficient against

that of a varied reactant concentration (Figures 3.4A and 3.4D) provides the rate-

dependence on that variable. The rate dependence on uraninite concentration (m2 L-1)

is described (x =) 0.4, with a with a rate coefficient (k’UO2) of 3.2 x 10-12; the

dependence on bicarbonate concentration is (x=) 0.2, with a rate coefficient (k’HCO3)

of 2.8 x 10-10. The relationship of pH to reaction rate across the entire range studied

here (Figure 3.2C) is described by the polynomial relationship,

kpH (s-1) = 5x10-13 + 0.11[H+]1.55 + 1.5x10-12[H+]-0.1 (3)

The overall dependence of uraninite oxidation rate at ferrihydrite

concentrations > 7.1 m2 L-1 is described by:

roxidation (mol s-1 L-1) = k * ([HCO3-]0.2) * ([UO2]0.4)) (4)

where k = 0.9*kpH and UO2 is measured in m2 L-1.

85

3.3.2 Iron (hydr)oxide Evolution

In the presence of small amounts of Fe(II) (<0.67 mmoles Fe(II) g-1

ferrihdyrite) ferrihydrite transforms into lepidocrocite and goethite (35). In order to

assure that the ferrihydrite was the dominant Fe-phase during the 10 min kinetic

studies, the evolution of ferrihydrite was investigated in a series of three longer

timescale experiments (~48 h). Additionally, since uraninite oxidation by goethite and

lepidocrocite is less thermodynamically favorable (Figure 3A.1), Fe(III) mineralogy

may limit uraninite oxidation over longer time periods. During the 48 h of oxidation,

dissolved Fe(II) remains below detection limits; extractable Fe(II) at all three

ferrihydrite concentrations reaches a maximum of ~ 0.2 mmoles Fe(II) g-1 ferrihydrite

(Figure 3.3D), and dissolved U(VI) also plateaus at 75, 112 and 211 µM for

ferrihydrite concentrations of 23.5, 46.0, and 80.2 m2 L-1 ferrihdyrite. At all

ferrihydrite concentrations, the solid phase remains predominantly ferrihydrite over

the first two hours of the experiment (Figures 3.3A-C). As the reaction progresses,

goethite continues forming, with detectable amounts of lepidocrocite accumulating

only after more than six hours of reaction. At the lowest ferrihydrite concentration, the

solid-phase speciation stabilizes after 24 h of reaction at a distribution of 36%

ferrihydrite, 19% lepidocrocite, and 44% goethite (Figure 3.3A). The transformation

of ferrihydrite is notably more complete at increasing ferrihydrite concentration, with

only 8% of the ferrihydrite remaining in the 46 m2 L-1 reaction and no ferrihydrite

remaining in the 80.2 m2 L-1 reaction after 48 h (Figure 3.3B and C).

Uraninite oxidation was confirmed by comparing the solid phase U oxidation

state of the last sample from each reaction to the oxidation state of biogenic uraninite.

86

The U(VI) content of the biogenic uraninite is approximately 5%, while the U(VI)

content at the conclusion of each oxidation reaction is 10, 18, and 26%, respectively,

for 23.5, 46.0, and 80.2 m2 L-1 ferrihydrite (Figure 3A.5).

87

3.4 DISCUSSION

The prevalence of materials capable of oxidizing biologically precipitated

uraninite and groundwater geochemical characteristics will influence the long-term

stability of UO2(biogenic). In the case of UO2(biogenic) oxidation by Fe(III) (hydr)oxides,

the thermodynamic viability of the reaction is dependent not only on Fe(III)

(hydr)oxide mineralogy, and the aqueous concentration of Ca2+, Fe(II), U(VI), and

HCO3- (33), but also on the U(IV) species involved in the oxidation reaction (Figure

3A.1). However, in order to determine the environmental significance of biogenic

uraninite oxidation by Fe(III) (hydr)oxides, it is necessary to determine the rate and

identify the mechanism of oxidation. It is particularly important to determine if a

soluble intermediate is involved in uraninite oxidation, as this would relieve the

necessity of solid-solid contact for electron transfer.

Iron(III) (hydr)oxides and biologically precipitated uraninite are generally

considered sparingly soluble under common environmental conditions; at pH 7, the

total dissolved Fe(III) concentration in equilibrium with ferrihydrite is predicted to be

approximately 1 x 10-10 M (34). While the solubility of poorly crystalline uraninite is

between 10-8 and 10-12 M, these measurements are frequently impacted by the

presence of U(VI) (19-21). In the absence of carbonate, the dissolution rate of

uraninite under anoxic conditions can be described by equation 5 below pH 7 and

equation 6 above pH 7 (46).

rdiss(UO2)(mol s-1 m-2) = 1.4(±0.3) x 10-8 x [H+]0.53±0.08 (5)

rdiss(UO2)(mol s-1 m-2) = 1.9(±0.8) x 10-12 (6)

88

The calculated uraninite oxidation rate at pH 7 (equation 4) is 7.3 x 10-12 M s-1

at pH 7 while the predicted dissolution rate of biogenic uraninite ({UO2} = 14 m2 L-1)

is 4.0 x 10-11 M s-1 at pH 7 (equation 5). Comparison of uraninite dissolution rates

(equations 5 and 6) and uraninite oxidation rates (equation 4) at the uraninite

concentrations examined here illustrates a good correlation between the two

predictions (Figure 3.4A).

The rate of dissolution will increase as the degree of undersaturation rises and

should therefore increase with U(IV)-binding ligand concentration. Indeed, Frazier et

al. (47) observe ligand enhanced dissolution rates of UO2 by desferrioxamine-B

(DFO-B). The role of bicarbonate in enhancing uraninite oxidation rates is

exemplified by the observed increase in oxidation rate at increasing bicarbonate

concentrations (Figure 3.2D), which is not predicted by equations 5 and 6.

The pH dependence of uraninite oxidation rate further supports the possibility

that the generation of a soluble U(IV) intermediate is the rate limiting step in uraninite

oxidation by ferrihydrite. Uraninite dissolution rates are poorly correlated with the

uraninite oxidation rates at varying pH (Figure 3.4B), owing to the inability of

equations 5 and 6 to account for the effect of U(IV)-carbonate complexes. While

uraninite dissolution rates have not been studied as a function of carbonate, Frazier et

al. (47) observed increasing dissolution rates above pH 7 in the presence of DFO-B.

Additionally, comparison of measured oxidation rates to calculated concentrations of

U(IV) species (which vary most with pH and bicarbonate) in equilibrium with

uraninite (UOH+ and U(CO3)56-) (Figure 3.5) reveal a strong correlation (Figure 3.4C

and 3.4D). Equation 4 overcomes the limitations of equations 5 and 6 by accounting

89

for a greater number of parameters (e.g. bicarbonate) which impact UO2 dissolution

and subsequent oxidation.

The dissolution rate of ferrihydrite also varies with pH, reaching a minimum at

pH 8 and increasing above and below this pH. If the dissolution of ferrihydrite is the

rate controlling step in uraninite oxidation, then the rate of uraninite oxidation should

reach a minimum at pH 8 and vary with the ferrihydrite suspension density. However,

the minimum oxidation rate is observed at pH 7.2, and at ferrihydrite concentrations

above 7.1 m2 L-1 the oxidation rate is independent ferrihydrite concentrations (Table

3.1, Figure 3.1B). At lower ferrihydrite concentration, a decrease in the oxidation rate

is observed, likely due to limited availability of ferrihydrite surface sites for uraninite

oxidation. A ferrihydrite concentration of 7.1 m2 L-1 is equivalent to a ferrihydrite

surface site concentration of ~30 µM, assuming 2 sites nm-2. Although the

concentration of dissolved U(IV) is likely in the nanomolar range, micromolar

concentrations of sorbed Fe(II) will likely play a role in limiting U(IV) surface access

at lower ferrihydrite concentrations. Additionally, pre-equilibration of ferrihydrite

with the reaction media does not result in an increase in the uraninite oxidation rate

(Table 3.1, Figure 3.1B), while pre-equilibration of the reaction media with uraninite,

prior to the addition of ferrihydrite, increases the initial reaction (Table 3.1, Figures

3.1A and 3A.4).

Although reaction 1 describes the overall oxidation reaction, given the data

presented, it is likely that the actual mechanism of uraninite oxidation by ferrihydrite

proceeds through several reaction steps (equations 8 to 11). Initially, uraninite

dissolution occurs (equation 8), followed by rapid equilibration of dissolved U(IV)

90

species (equation 9). Dissolved U(IV) then sorbs to the ferrihydrite surface (reaction

10), followed by rapid electron transfer and U(VI) release to solution (equation 11).

91

>U-OH2+ + (n-1)H2O → U(OH)n

(4-n) + n(H+) (8)

U(OH)n

(4-n) + mHCO3- ↔ U(CO3)m

(4-2m) + (m-n)H+ (9)

>Fe(III)2 + U(OH)n(4-n) + n(H+) → >Fe(III)≡U(IV) + n(H2O) (10)

≡Fe(II) + U(VI)O22+ + nHCO3

- → ≡Fe(II) + UO2(CO3)n(2-n) + nH+ (11)

This series of reactions accounts for the fractional dependence of the oxidation rate on

both bicarbonate and proton concentrations as they both are involved in several of the

oxidation reaction steps. Given the evidence presented here, it appears likely that the

rate controlling step in uraninite oxidation is dissolution (equation 8), since conditions

that increase uraninite solubility also increase the oxidation rate. Equation 6

incorporates these conditions (pH and bicarbonate) to provide a more accurate

depiction of the reactivity of biogenic uraninite in the presence of ferrihydrite.

3.4.1 Evolution of Iron (hydr)oxides

The transformation of ferrihydrite into less favorable oxidants may ultimately

limit biogenic uraninite oxidation over longer time frames. As uraninite oxidation

proceeds, Fe(II) induces the transformation of ferrihydrite into lepidocrocite and

goethite, both of which are less favorable oxidants than ferrihydrite (Figure 3A.1).

Although, over the time frame of the oxidation rate studies (10 min), the solid phase

remains > 95% ferrihydrite (Figures 3.3A-C). At all three ferrihydrite concentrations

studied, extractable ferrous iron increases rapidly during the first 1 to 2 h of the

reaction and can be predicted using equation 4 (Figure 3.3D). The extractable Fe(II)

concentration then levels off after approximately 10 h at ~ 0.2 mmoles Fe(II) g-1

92

ferrihydrite (Figure 3.3D), which is not accounted for in equation 4. Despite the

apparent cessation of uraninite oxidation after ~ 10 h of reaction time, uraninite

remains after 48 h of reaction (Figure 3A.5). This may indicate that 0.2 mmoles Fe(II)

g-1 ferrihydrite is a thermodynamic threshold for the energetic favorability of uraninite

oxidation, or that the build-up of a reaction results in the termination of the reaction

(33). Indeed, uraninite oxidation by ferrihydrite, lepidocrocite, and goethite is no

longer thermodynamically favorable at the reaction conditions observed after 10 h of

reaction (~100 µM U(VI)(aq), pH 7.2, 3 mM HCO3-, and 1 µM Fe(II)(aq)). The

oxidation of dissolved U(IV) species by lepidocrocite is thermodynamically favorable

at the low Fe(II)(aq) and U(VI)(aq) concentrations during the first 5 hours of reaction,

and since its presence is not initially detected, it may be a transient species being

rapidly consumed by uraninite oxidation.

93

3.5 IMPLICATIONS FOR BIOGEOCHEMICAL URANIUM CYCLING

The oxidation of biologically precipitated uraninite by Fe(III) (hydr)oxides

may limit the stability of biologically precipitated uraninite under mildly reducing

conditions. Prevailing geochemical conditions can drastically affect the rate of

uraninite oxidation by ferrihydrite and geochemical conditions that result in an

increase in the degree of undersaturation for uraninite (acidic or alkaline pH and

increased bicarbonate concentration) increase the rate of uraninite oxidation by

ferrihydrite. These results are consistent with uraninite dissolution being the rate

limiting step in the oxidation reaction. The generation of Fe(II) during uraninite

oxidation causes ferrihydrite to transform into more thermodynamically stable Fe(III)

(hydr)oxides, and when combined with Fe(II) and U(VI) generation, limits the extent

of uraninite oxidation in static systems where reaction product transport is limited. In

sum, uraninite oxidation by ferrihydrite likely proceeds through a soluble intermediate

and is accelerated under conditions that increase the solubility of uraninite. Ultimately

, the rate of U(VI) formation is well described by a fractional dependence on uraninite

suspension density and bicarbonate concentration, with a non-linear dependence on

pH – all of which impact the rate of UO2 dissolution.

94

3.6 ACKNOWLEDGEMENTS

We thank Kathleen Beman and Brandy Stewart for their constructive input on

the manuscript. This work was funded by the Office of Science Biological and

Environmental Research ERSD Program, U.S. Department of Energy (grant number

ER63609-1021814). A portion of this work was conducted at Stanford Synchrotron

Radiation Laboratory, a national user facility operated by Stanford University on

behalf of the U.S. Department of Energy, Office of Basic Energy Sciences. The SSRL

Structural Molecular Biology Program is supported by the Department of Energy,

Office of Biological and Environmental Research, and by the National Institutes of

Health, National Center for Research Resources, Biomedical Technology Program.

95


(1) Bradford, G. R.; Bakhtar, D.; Westcot, D. Uranium, vanadium, and molybdenum in saline waters of California. Journ. Environ. Qual. 1990, 19, 105-108. (2) Riley, R. G.; Zachara, J. M.; Wobber, F. J. "Chemical contaminants on DOE lands and selection of contaminant mixtures for subsurface science research," U.S. Department of Energy, 1992. (3) Langmuir, D. Uranium solution-mineral equilibria at low temperature with applications to sedimentary ore deposits. Geochim. Cosmochim. Acta 1978, 42, 547-569. (4) Sandino, A.; Bruno, J. The solubility of (UO2)3(PO4)2·4H2O(s) and the formation of U(VI) phospate complexes: Their influence in uranium speciation in natural waters. Geochim. Cosmochim. Acta 1992, 56, 4135-4145. (5) Moyes, L. N.; Parkman, R. H.; Charnock, J. M.; Vaughan, D. J.; Livens, F. R.; Hughes, C. R.; Braithwaite, A. Uranium Uptake from Aqueous Solution by Interaction with Goethite, Lepidocrocite, Muscovite, and Mackinawite: An X-ray Absorption Spectroscopy Study. Environ. Sci. Technol. 2000, 34, 1062-1068. (6) Bostick, B. B.; Fendorf, S.; Barnett, M. O.; Jardine, P. M.; Brooks, S. C. Uranyl Surface Complexes Formed on Subsurface Media from DOE Facilities. Soil Sci. Soc. Am. J. 2002, 66, 99-108. (7) Barnett, M. O.; Jardine, P. M.; Brooks, S. C.; Selim, H. M. Adsorption and transport of uranium(VI) in subsurface media. Soil Sci. Soc. Am. J. 2000, 64, 908-917. (8) Barnes, C. E.; Cochran, J. K. Uranium geochemistry in esturaine sediments: Controls on removal and release processes. Geochim. Cosmochim. Acta 1993, 57, 555-569. (9) Grenthe, I.; Fuger, J.; Konings, R. J. M.; Lemire, R. J.; Muller, A. B.; Nguyen-Trung, C.; Wanner, H. Chemical thermodynamics of uranium; North-Holland Elsevier Science Publishers B.V.: Amsterdam, 1992; Vol. 1. (10) Hua, B.; Xu, H.; Terry, J.; Deng, B. Kinetics of uranium(VI) reduction by hydrogen sulfide in anoxic aqueous systems. Environ. Sci. Technol. 2006, 40, 4666-4671. (11) Wersin, P.; Hochella, M. F.; Persson, P.; Redden, G.; Leckie, J. O.; Harris, D. W. Interaction between aqueous uranium(VI) and sulfide minerals : Spectroscopic evidence for sorption and reduction. Geochim. Cosmochim. Acta 1994, 58, 2829-2843. (12) Liger, E.; Charlet, L.; Cappellen, P. V. Surface catalysis of uranium(VI) reduction by iron(II). Geochim. Cosmochim. Acta 1999, 63, 2939-2955. (13) Fredrickson, J. K.; Zachara, J. M.; Kennedy, D. W.; Duff, M. C.; Gorby, Y. A.; Li, S. M. W.; Krupka, K. M. Reduction of U(VI) in goethite (alpha-FeOOH) suspensions by a dissimilatory metal-reducing bacterium. Geochim. Cosmochim. Acta 2000, 64, 3085-3098. (14) Lovley, D. R.; Phillips, E. J. P. Bioremediation of uranium contamination with enzymatic uranium reduction. Environ. Sci. Technol. 1992, 26, 2228-2234.

96

(15) Gorby, Y. A.; Lovley, D. R. Enzymatic Uranium Precipitation. Environ. Sci. Technol. 1992, 26, 205-207. (16) Wu, W.-M.; Carley, J.; Fienen, M.; Mehlhorn, T.; Lowe, K.; Nyman, J.; Luo, J.; Gentile, M.; Rajan, R.; Wagner, D.; Hickey, R.; Gu, B.; Watson, D. B.; Cirpka, O.; Kitanidis, P.; Jardine, P. M.; Criddle, C. Pilot-scale in situ bioremediation of uranium in a highly contaminated aquifer. 1. Conditioning of a treatment zone. Environmental Science & Technology 2006, 40, 3978-3985. (17) Wu, W.-M.; Carley, J.; Gentry, T.; Ginder-Vogel, M.; Fienen, M.; Mehlhorn, T.; Yan, H.; Caroll, S.; Pace, M.; Nyman, J.; Luo, J.; Gentile, M.; Fields, M. W.; Hickey, R.; Watson, D. B.; Cirpka, O.; Zhou, J.; Fendorf, S.; Kitanidis, P.; Jardine, P. M.; Criddle, C. Pilot-scale in situ bioremediation of uranium in a highly contaminated aquifer. 2. Geochemical control of U(VI) bioavailability and evidence of U(VI) reduction. Environ. Sci. Technol. 2006, 40, 3986-3995. (18) Anderson, R. T.; Vrionis, H. A.; Ortiz-Bernard, I.; Resch, C. T.; Long, P. E.; Dayvault, R.; Karp, K.; Marutzky, S.; Metzler, D. R.; Peacock, A. D.; White, D. C.; Lowe, M.; Lovley, D. R. Stimulating the in situ activity of Geobacter species to remove uranium from the groundwater of a uranium-contaminated aquifer. Appl. Environ. Microb. 2003, 69, 5884-5891. (19) Abdelouas, A.; Lutze, W.; Nutall, H. E. Oxidative dissolution of uraninite precipitated on Navajo sandstone. Journal of Contaminant Hydrology 1999, 36, 353-375. (20) Casas, I.; dePablo, J.; Gimenez, J.; Torrero, M. E.; Bruno, J.; Cera, E.; Finch, R. J.; Ewing, R. C. The role of pe ; pH ; and carbonate on the solubility of UO2 and uraninite under nominally reducing conditions. Geochim. Cosmochim. Acta 1998, 62, 2223-2231. (21) Rai, D.; Felmy, A. R.; Ryan, J. L. Uranium(IV) hydrolysis constants and solubility product of UO2 xH2O(am). Inorg. Chem. 1990, 29, 260-264. (22) Fredrickson, J. K.; Zachara, J. M.; Kennedy, D. W.; Liu, C. G.; Duff, M. C.; Hunter, D. B.; Dohnalkova, A. Influence of Mn oxides on the reduction of uranium(VI) by the metal-reducing bacterium Shewanella putrefaciens. Geochim. Cosmochim. Acta 2002, 66, 3247-3262. (23) Finneran, K. T.; Housewright, M., E.; Lovley, D. R. Multiple influences of nitrate on uranium solubility during bioremediation of uranium-contaminated subsurface sediments. Environ. Microbiol. 2002, 4, 510-516. (24) Senko, J. M.; Istok, J. D.; Suflita, J. M.; Krumholz, L. R. In-situ evidence for uranium immobilization and remobilization. Environmental Science & Technology 2002, 36, 1491-1496. (25) Istok, J. D.; Senko, J. M.; Krumholz, L. R.; Watson, D.; Bogle, M.; Peacock, A. D.; Chang, Y.-J.; White, D. C. In situ bioreduction of technetium and uranium in a nitrate contaminated aquifer. Environ. Sci. Technol. 2004, 38, 468-475. (26) Senko, J. M.; Mohamed, Y.; Dewers, T.; Krumholz, L. R. Role for Fe(III) minerals in nitrate-dependent microbial U(IV) oxidation. Environ. Sci. Technol. 2005, 39, 2529-2536.

97

(27) Beller, H. R. Anaerobic, nitrate-dependent oxidation of U(IV) oxide minerals by the chemolithoautotrophic bacterium Thiobacillus dentrificans. Appl. Environ. Microb. 2005, 71, 2170-2174. (28) Nevin, K. P., Derek R. Lovley Potential for Nonenzymatic Reduction of Fe(III) via Electron Shuttling in Subsurface Sediments. Environ. Sci. Technol. 2000, 34, 2472-2478. (29) Wan, J.; Tokunaga, T. K.; Brodie, E.; Wang, Z.; Zheng, Z.; Herman, D.; Hazen, T.; Firestone, M. K.; Sutton, S. R. Reoxidation of bioreduced uranium under reducing conditions. Environ. Sci. Technol. 2005, 39, 6162-6169. (30) Sani, R. K.; Peyton, B. M.; Amonette, J. E.; Geesey, G. G. Reduction of uranium(VI) under sulfate-reducing conditions in the presence of Fe(III)-(hydr)oxides. Geochim. Cosmochim. 2004, 68, 2639-2648. (31) Sani, R. K.; Peyton, B. M.; Dohnalkova, A.; Amonette, J. E. Reoxidation of reduced uranium with iron(III) (hydr)oxides under sulfate-reducing conditions. Environ. Sci. Technol. 2005, 39, 2059-2066. (32) Senko, J. M.; Suflita, J. M.; Krumholz, L. R. Geochemical controls on microbial nitrate-dependent U(IV) oxidation. Geomicrobiology Journal 2005, 22, 371-378. (33) Ginder-Vogel, M.; Criddle, C.; Fendorf, S. Thermodynamic constraints on the oxidation of biogenic UO2 by Fe(III) (hydr)oxides. Environ. Sci. Technol. 2006, 40, 3544-3550. (34) Cornell, R. M.; Schwertmann, U. The Iron Oxides: Structure, Properties, Reactions, Occurrences and Uses; Wiley-VCH, 2003. (35) Hansel, C. M.; Benner, S. G.; Fendorf, S. Competing Fe(II)-induced mineralization pathways of ferrihydrite. Environ. Sci. Technol. 2005, 39, 7147-7153. (36) Hansel, C. M.; Benner, S. G.; Neiss, J.; Dohnalkova, A.; Kukkadapu, R. K.; Fendorf, S. Secondary mineralization pathways induced by dissimilatory iron reduction of ferrihydrite under advective flow. Geochim. Cosmochim. Acta 2003, 67, 2977-2992. (37) Singer, D.; Farges, F.; Brown, G. E. Biogenic UO2: Characterization and surface reactivity. Physica Scripta 2006, In Press. (38) Stookey, L. L. A new spectrophotometric reagent for iron. Anal. Chem. 1970, 42, 779-781. (39) Morrison, S. J.; Spangler, R. R.; Tripathi, V. S. Adsorption of uranium(VI) on amorphous ferric oxyhydroxide at high concentrations of dissolved carbon(IV) and sulfur(VI). Jour. Contam. Hydrol. 1995, 17, 333-346. (40) Walter, M.; Arnold, T.; Reich, T.; Bernhard, G. Sorption of uranium(VI) onto ferric oxides in sulfate-rich acid waters. Environ. Sci. Technol. 2003, xx. (41) Duff, M. C.; Coughlin, J. U.; Hunter, D. B. Uranium co-precipitation with iron oxide minerals. Geochim. Cosmochim. Acta 2002, 66, 3533-3547. (42) Hammersley, A. P. "FIT2D: An Introduction and Overview," European Synchrotron Radiation Facility, 1997. (43) Webb, S. M. Sixpack: A graphical user interface for XAS analysis using IFEFFIT. Physica Scripta 2005, T115, 1011-1014.

98

(44) Newville, M. IFEFFIT: interactive XAFS analysis and FEFF fitting. J. Synchrotron Radiation 2001, 8, 322-324. (45) Bertsch, P. M.; Hunter, D. B. In Situ Chemical Speciation of Uranium in Soils and Sediments by Micro X-ray Absorption Spectroscopy. Environmental Science and Technology 1994, 28, 980-984. (46) Bruno, J.; Casas, I.; Puigdomenech, I. The kinetics of dissolution of UO2 under reducing conditions and the influence of an oxidized surface layer (UO2+x): Application of a continuous flow-through reactor. Geochim. Cosmochim. Acta 1991, 55, 647-658. (47) Frazier, S. W.; Kretzschmar, R.; Kraemer, S. M. Bacterial siderophores promote dissolution of UO2 under reducing conditions. Environ. Sci. Technol. 2005, 39, 5709-5715. (48) Allison, J. D.; Brown, D. S.; Nova-Gradac, K. J. MINTEQA2/PRODEFA2, A geochemical assessment model for environmental systems: Version 3.0 user's manual; U.S. EPA: Athens, GA, 1990. (49) Guillaumont, R.; Fanghanel, T.; Neck, V.; Fuger, J.; Palmer, D. A.; Grenthe, I.; Rand, M. H. "Update on the Chemical Thermodynamics of Uranium, Neptumium, Plutonium, Americium, and Technetium," Nuclear Energy Agency, 2003.

99

Table 3.1 Reaction conditions and observed first order rate coefficients.

pH UO2 (m2 L-1)

Ferrihydrite (m2 L-1)

HCO3-

(mM) kobs (s-1)

pH Variation 6.3 13.8 29.7 3 2.62 x 10-11

6.7 13.8 29.7 3 1.17 x 10-11 7.2 13.8 29.7 3 8.3 x 10-12 8 13.8 29.7 3 1.05 x 10-11 9 13.8 29.7 3 1.20 x 10-11

Uraninite Variation 7.2 7.0 29.7 3 6.2 x 10-12 7.2 13.8 29.7 3 8.3 x 10-12 7.2 20.6 29.7 3 1.1 x 10-11 7.2 27.6 29.7 3 1.2 x 10-11 7.2 13.8 29.7 3 1.0 x 10-11*

Bicarbonate Variation 7.2 13.8 29.7 3 8.3 x 10-12 7.2 13.8 29.7 6 9.5 x 10-12 7.2 13.8 29.7 9 1.1 x 10-11

Ferrihydrite Variation 7.2 13.8 2.0 3 2.1 x 10-12 7.2 13.8 4.0 3 5.3 x 10-12 7.2 13.8 7.1 3 7.9 x 10-12 7.2 13.8 13.2 3 8.0 x 10-12 7.2 13.8 22.6 3 8.9 x 10-12 7.2 13.8 29.7 3 8.3 x 10-12 7.2 13.8 29.7 3 8.2 x 10-12#

*Uraninite equilibrated with reaction for 10 minutes prior to addition of ferrihydrite #Ferrihydrite equilibrated for 10 minutes prior to the addition of uraninite

100

Figure Captions

Figure 3.1 First order rate plots as a function of uraninite (A), ferrihydrite (B), pH (C),

and HCO3- (D) concentration. Error bars show the standard deviation of the three

measurements at each time point, the R2 of all linear regression measurements is >

0.95.

Figure 3.2 Log kobs vs log uraninite concentration (A), ferrihydrite concentration (B),

pH (C) and bicarbonate concentration (D). The error for each point is calculated from

the standard error in the slope of the linear regression (Figure 3.1). Linear regression

lines and equations are shown for A and D, while the fit using equation 3 is shown in

C.

Figure 3.3 Transformation of ferrihydrite by Fe(II) generated during uraninite

oxidation. Reactions represented all contain 3.0 mM KHCO3, 1.0 mM UO2 and either

(A) 2.2, (B) 4.5, or (C) 7.3 mM ferrihydrite in 3 mM HCO3- at pH 7.2. Percentages (±

5%) were determined from linear combination fits of k3-weighted Fe EXAFS spectra

(k=1 - 14) (supporting information). (D) Extractable Fe(II) for each reaction.

Figure 3.4 Comparison of predicted uraninite dissolution rates (equations 5 and 6) to

uraninite oxidation rates (equation 4) as a function of uraninite concentration (A) and

pH (B). Comparison of calculated UOH3+ and U(CO3)56- concentrations oxidation

rates as a function of pH (C) and bicarbonate concentration (D). Linear regression

lines and R2 values are shown on each graph.

Figure 3.5 Predicted U(IV) speciation in 3 mM KHCO3 in equilibrium with UO2(am),

calculated with Visual Minteq (48) and thermodynamic data from Guillaumont et al.

(49).

101

Time (s)

0 100 200 300 400 500 600 700

ln(F

e(II)

) (M

)

-15

-14

-13

-12

-11

-10

-97.0 m2 L-1 UO2

13.8 m2 L-1 UO2

20.6 m2 L-1 UO2

27.6 m2 L-1 UO2

13.8 m2 L-1 UO2

10 min equilibration

A

Time (s)

0 100 200 300 400 500 600 700

ln(F

e(II)

) (M

)

-15

-14

-13

-12

-11

-102.0 m2 L-1 FHY4.0 m2 L-1 FHY7.1 m2 L-1 FHY13.2 m2 L-1 FHY22.6 m2 L-1 FHY29.7 m2 L-1 FHY29.7 m2 L-1 FHY10 min equilibration

B

102

Time (s)

0 100 200 300 400 500 600 700

ln(F

e(II)

) (M

)

-13.5

-13.0

-12.5

-12.0

-11.5

-11.0

-10.5pH 6.3 pH 6.7 pH 7.2 pH 8 pH 9

C

Time (s)

0 100 200 300 400 500 600 700

ln(F

e(II)

) (M

)

-13.4

-13.2

-13.0

-12.8

-12.6

-12.4

-12.2

-12.0

-11.8

-11.6

3 mM HCO3-

6 mM HCO3-

9 mM HCO3-

D

Figure 3.1 First order rate plots as a function of uraninite (A), ferrihydrite (B), pH (C),

and HCO3- (D) concentration. Error bars show the standard deviation of the

three measurements at each time point, the R2 of all linear regression measurements is > 0.95.

103

log (UO2) (m2 L-1)

0.8 0.9 1.0 1.1 1.2 1.3 1.4 1.5

log

(kob

s) (s

-1)

-11.25

-11.20

-11.15

-11.10

-11.05

-11.00

-10.95

-10.90A

R2 = 0.95y = 0.39x - 11.5

log (Ferrihdyrite) m2 L-1

0.2 0.4 0.6 0.8 1.0 1.2 1.4 1.6

log

k obs (

s-1)

-11.6

-11.4

-11.2

-11.0

-10.8

-10.6

B

104

pH

6.0 6.5 7.0 7.5 8.0 8.5 9.0 9.5

log

k obs (

s-1)

-11.2

-11.1

-11.0

-10.9

-10.8

-10.7

-10.6

-10.5

R2 = 0.96

C

log (HCO3-) (M)

-2.6 -2.5 -2.4 -2.3 -2.2 -2.1 -2.0

log

(kob

s) (s

-1)

-11.15

-11.10

-11.05

-11.00

-10.95

D

R2 = 0.99

y = 0.25x - 10.45

Figure 3.2 Log kobs vs log uraninite concentration (A), ferrihydrite concentration (B),

pH (C) and bicarbonate concentration (D). The error for each point is calculated from the standard error in the slope of the linear regression (Figure 3.1). Linear regression lines and equations are shown for A and D, while the fit using equation 3 is shown in C.

105

Time (h)

0 10 20 30 40 50

Mol

e P

erce

nt Ir

on

0

20

40

60

80

100

Ferrihydrite

Goethite

Lepidocrocite

A

Time (h)

0 10 20 30 40 50

Mol

e P

erce

nt Ir

on

0

20

40

60

80

100

Ferrihydrite

Goethite

Lepidocrocite

B

106

Time (h)

0 10 20 30 40 50

Mol

e P

erce

nt Ir

on

0

20

40

60

80

100

Ferrihydrite

Goethite

Lepidocrocite

C

Time (h)

0 10 20 30 40 50

Ext

ract

able

Fe(

II) (m

mol

g-1

ferri

hydr

ite)

0.00

0.05

0.10

0.15

0.20

0.25

23.5 m2 L-1 Ferrihydrite46.0 m2 L-1 Ferrihydrite80.2 m2 L-1 Ferrihdyirte

D

Figure 3.3 Transformation of ferrihydrite by Fe(II) generated during uraninite oxidation. Reactions represented all contain 3.0 mM KHCO3, 30.7 m2 L-1 UO2 and either (A) 23.5, (B) 46.0, or (C) 80.2 m2 L-1 mM ferrihydrite in 3 mM HCO3

- at pH 7.2. Percentages (± 5%) were determined from linear combination fits of k3-weighted Fe EXAFS spectra (k=1 - 14) (Figure 3A.3) (D) Extractable Fe(II) for each reaction.

107

log rateoxidation (M s-1)(Uraninite Oxidation Rate)

-10.90 -10.85 -10.80 -10.75 -10.70 -10.65

log

rate

diss

olut

ion (

M s

-1)

(Ura

nini

te D

isso

lutio

n R

ate)

-11.20

-11.15

-11.10

-11.05

-11.00

-10.95

A

R2 = 0.95

7.0 m2 L-1 UO2

13.8 m2 L-1 UO2

20.6 m2 L-1 UO2

27.6 m2 L-1 UO2


-11.2 -11.1 -11.0 -10.9 -10.8 -10.7 -10.6 -10.5

log

rate

diss

olut

ion (

Ms-1

)(U

rani

nite

Dis

solu

tion

Rat

e)

-10.6

-10.5

-10.4

-10.3

-10.2

-10.1

-10.0

-9.9

B

R2 = 0.69

pH 6.3

pH 6.7

pH 9.0

pH 8.0

pH 7.2

108


-11.2 -11.1 -11.0 -10.9 -10.8 -10.7 -10.6 -10.5

log

(UO

H3+

+ U

(CO

3)56-

) (M

)(P

redi

cted

Equ

ilibriu

m C

once

ntra

tion)

-23.5

-23.0

-22.5

-22.0

-21.5

-21.0

-20.5

-20.0

C

R2 = 0.88

pH 6.3

pH 6.7pH 9.0

pH 8.0

pH 7.2


-11.10 -11.08 -11.06 -11.04 -11.02 -11.00 -10.98 -10.96 -10.94

log

(UO

H3+

+ U

(CO

3)56-

) (M

)(P

redi

cted

Equ

ilibriu

m C

once

ntra

tion)

-23.5

-23.0

-22.5

-22.0

-21.5

-21.0

D

R2 = 0.99

9 mM HCO3-

6 mM HCO3-

3 mM HCO3-

Figure 3.4 Comparison of predicted uraninite dissolution rates (equations 5 and 6) to uraninite oxidation rates (equation 4) as a function of uraninite concentration (A) and pH (B). Comparison of calculated UOH3+ and U(CO3)5

6- concentrations oxidation rates as a function of pH (C) and bicarbonate concentration (D). Linear regression lines and R2 values are shown on each graph.

109

pH

2 4 6 8 10

log

conc

entra

tion

(M)

-80

-70

-60

-50

-40

-30

-20

-10

0

U(OH)4(aq)

UOH3+

U(CO3)56-

U(CO3)44-

U4+

Figure 3.5 Predicted U(IV) speciation in 3 mM KHCO3 in equilibrium with UO2(am), calculated with Visual Minteq (48) and thermodynamic data from Guillaumont et al. (49).

110

APPENDIX 3A


111

Figure 3A.1 Representative Fe(III)/Fe(II) and U(VI)/U(IV) redox couples with 3 x 10-

3 M HCO3-, 1 x 10-6 M of each U(VI) species, 1 x 10-9 M of each dissolved U(IV)

species, and 1 x 10-5 M Fe(II).

112

Time (s)

0 100 200 300 400 500 600 700

Ope

n S

ymbo

ls -

Ext

ract

able

Fe(

II) (µ

M)

Clo

sed

Sym

bols

- D

isso

lved

U(V

I) (µ

M)

0

2

4

6

8

10

12

14

16

7.0 m2 L-1 UO2

13.8 m2 L-1 UO2

20.6 m2 L-1UO2

27.6 m2 L-1 UO2

Figure 3A.2 Comparison of extractable Fe(II) and dissolved U(VI) concentrations as a

function of uraninite concentration.

113

Figure 3A.3 Fe K-edge EXAFS spectra (solid lines) and linear combination

fits (dotted lines) used to construct Figure 3A-C. XRD patterns used to identify Fe minerals for linear combination fits (left side of A), g=goethite, l=lepidocrocite, and u=uraninite.

k (Å-1)

2 4 6 8 10 12 14 16

chi(k

)*k3

0.5 hours

1.0 hours

1.5 hours

5.0 hours

8.5 hours

12.5 hours

46.5 hours

25.5 hours

A

Å

k (Å-1)

2 4 6 8 10 12 14 16

chi *

k3

C0.5 hours

1.5 hours

3.0 hours

6.0 hours

12.5 hours

25.0 hours

49.0 hours

36.0 hours

k (Å-1)

2 4 6 8 10 12 14 16

chi*k

3

B

0.5 hours

1.0 hours

2.0 hours

3.0 hours

6 hours

16 hours

45.0 hours

24.0 hours

2 theta (degrees)0 10 20 30 40 50 60

Inte

nsity

(Cou

nts)

Ferrihydrite

0.5 hours

1.0 hours

1.5 hours

5.0 hours

8.5 hours

12.5 hours

46.5 hours

25.5 hoursL

LL/G

L L/G

L/UL LL LL GLG

GG GGGGG

G/UUU

U

UG

G

114

Time (s)

0 100 200 300 400 500 600 700

Ext

ract

able

Fe(

II) (µ

M)

0

2

4

6

8

10

12

14

27.6 m2 L-1 UO2

20.6 m2 L-1 UO2

13.8 m2 L-1 UO2

7.0 m2 L-1 UO2

13.8 m2 L-1 UO2 10 m equilibration

A

Time (s)

0 100 200 300 400 500 600 700

Ext

ract

able

Fe(

II) (µ

M)

0

1

2

3

4

5

62.0 m2 L-1 Ferrihydrite4.0 m2 L-1 Ferrihydrite7.1 m2 L-1 Ferrihydrite13.2 m2 L-1 Ferrihydrite22.6 m2 L-1 Ferrihydrite29.7 m2 L-1 Ferrihydrite

B

115

Time (s)

0 100 200 300 400 500 600 700

Ext

ract

able

Fe(

II) (µ

M)

0

2

4

6

8

10

12

14

16

18

pH 6.3 pH 6.7 pH 7.2 pH 8.0 pH 9.0

C

Time (s)

0 100 200 300 400 500 600 700

Ext

ract

able

Fe(

II) (µ

M)

0

1

2

3

4

5

6

7

8

3 mM HCO3-

9 mM HCO3-

6 mM HCO3-

D

Figure 3A.4 Extractable Fe(II) produced during uraninite oxidation as a function of uraninite concentration (A), ferrihydrite concentration (B), pH (C), or bicarbonate concentration (D). Error bars show the standard deviation of the three measurements at each time point.

116

Energy (eV)

17160 17180 17200 17220

Nor

mal

ized

Inte

nsity

0.0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

1.6

1.8

Biogenic UraniniteUranyl Nitrate - U(VI)23.5 m2 L-1 Ferrihydrite - T = 48 h46.0 m2 L-1 Ferrihydrite - T = 46.5 h80.2 m2 L-1 Ferrihdyrite - T = 49 h

Figure 3A.5 Uranium LIII XANES spectra showing uranium oxidation state at the conclusion of each experiment.

117

CHAPTER 4

Chromate Reduction and Retention Processes within Arid Subsurface Environments

Matthew Ginder – Vogel1, Thomas Borch1, Melanie Mayes2, Phillip Jardine2, and Scott Fendorf1

1. Department of Geological and Environmental Science, Stanford University, Stanford, CA

2. Oak Ridge National Laboratory, Oak Ridge, TN Published as: Ginder-Vogel, M.; Borch, T.; Mayes, M. A.; Jardine, P. M.; Fendorf, S. Chromate reduction and retention processes within arid subsurface environments. Environ. Sci. Technol. 2005, 39, 7833-7839.

118

ABSTRACT

A number of chromate reductants, inclusive of Fe(II) (aq), magnetite, green

rust, ilmenite, and bacteria, have been extensively studied in model systems, but few

studies have investigated reduction processes within natural sediments, particularly

those from aerated systems. Accordingly, we examine chromate reduction by arid

sediments from the Hanford, WA site. Iron(III) (hydr)oxide coatings limit mineral

reactivity in this arid environment; high pH (pH = 14), representative of conditions

near high-level nuclear waste tanks, results in Fe(II) solubilization and concurrent

Cr(VI) reduction. The X-ray spectroscopic analysis of solids from columns that were

either acid pre-treated or included 10 M NaOH reveals that reduced chromium,

Cr(III), occurs in association with antigorite and lizardite, in addition to magnetite and

Fe(II)-bearing clay minerals. Additionally, in a column containing 10 M NaOH in the

feed solution, Cr(III) and Cr(VI) are found associated with portlandite, suggesting a

secondary mechanism for chromium retention at high pH. The data and analysis

presented herein provide novel information on reactions of chromate with ferrous-

bearing primary minerals within natural sediments.

119

4.1 INTRODUCTION

Chromium has been introduced into the environment primarily through its

widespread use in industrial applications, including tanning, metallurgy, and plating.

Once introduced into the environment, chromium persists as either Cr(VI) or Cr(III).

Hexavalent Cr exists primarily in groundwater systems as the oxyanion CrO42-

(chromate), or its protonated counterparts, which exhibit high water solubility over

much of the environmental pH range, are strong oxidants, and known mutagens,

teratogens, and carcinogens (1). In contrast to chromate, trivalent chromium is

relatively non-toxic and forms strong complexes with soil minerals and sparingly

soluble hydroxide precipitates at circumneutral pH (2).

Hexavalent chromium can be reduced to trivalent chromium by aqueous and

sorbed Fe(II) (3,4), organic matter (5), Fe(II)-bearing minerals (6-8), sulfide

compounds (9,10), and through microbial processes (11). These reductants all

ultimately lead to Cr(III); however, the resultant products and rates of reduction differ.

Enzymatic reduction of chromate may result in the formation of soluble organic

Cr(III) complexes (12-14), while chromate reduction by soluble Fe(II) results in the

formation of Cr1-xFex(OH)3 precipitates having limited solubility (2). Reaction of

magnetite with high concentrations of dissolved chromate result in the formation of

maghemite (γ-Fe2O3) and an over-layer of Cr(OH)3·nH2O (6,15).

Subsurface chromate contamination is a major environmental threat at the U.S.

Department of Energy’s Hanford Site (southeastern Washington State, USA), the

location of plutonium production starting during World War II. The climate of the

Hanford Site is arid, with an average rainfall of 15.9 cm y-1, resulting in a 10 m to 60

120

m deep vadose zone (16). In the 100 Area of the Hanford Site, chromate used as a

corrosion inhibitor in nuclear reactor cooling water was discharged to unlined surface

cribs, resulting in chromate contamination reaching the Columbia River (16), which is

not only a primary source of drinking water, but also a spawning ground for salmon.

Additionally, thousands of liters of hot (100 Co), caustic (pH>14), chromate-

containing high-level nuclear waste (HLW), from the Reduction-Oxidation (REDOX)

plutonium recovery process, has leaked into the vadose zone, as a result of multiple

failures of single-shell tanks in the S-SX tank farm (17). Should chromate from these

tanks reach the water table, rapid migration to the Columbia River would result.

Analysis of chromium’s redox state in a core obtained from a HLW plume beneath

tank SX 108 revealed that 29 – 75% of the total Cr had been reduced to immobile

Cr(III) (18).

In aerobic, arid environments, with limited organic carbon (such as the

Hanford Site), chromate reductants are essentially restricted to Fe(II)-bearing mineral

phases, primary in origin. The uppermost principal geologic unit of the Hanford Site

consists of material deposited in cataclysmic ice age floods during the past 12,000 to

700,000 y, termed the Hanford formation. Potential sources of solid-phase Fe(II) for

chromate reduction in the Hanford sediments include ferrous iron-bearing silicates

such as lizardite [MgxFe3-xSi2O5(OH)4] (this work), antigorite [(Mg,Fe)3Si2O5(OH)4]

(this work), biotite (19), magnetite, and ilmenite (19). Chromate reduction by Fe(II)-

bearing minerals is controlled by their surface reactivity. Mineral surface coatings

such as carbonates (20), silicates, and Fe(III) (hydr)oxides (6,15) inhibit electron

121

transfer from underlying Fe(II) to aqueous Cr(VI), all of which commonly form in

oxic, arid, vadose-zone environments.

In the Hanford subsurface, chromium contamination has occurred in two

distinct geochemical environments, termed the “near-field” and “far-field”

environments in this communication. The near-field environment is adjacent to

leaking tanks of HLW and is characterized by high pH, salt concentration, and

temperature, resulting in extensive mineral dissolution and surface modification. The

far-field environment is representative of the background geochemical conditions

(circumneutral pH, ambient temperature, etc.). These conditions would be relevant at

locations far from a leaking tank, or in arid environments with limited co-

contamination.

Chromate reduction by Fe(II) (aq), magnetite, ilmenite, and biotite has been

studied extensively in model systems at acidic and neutral pH (6,9,21,22), and in the

case of magnetite (23) and Fe(II) (aq) (24) at high pH as well. What remains elusive

are the minerals serving as chromate reductants and the factors limiting chromate

reduction within arid soils and sediments. In the present work, we examine the

reduction of chromate by sediment obtained from the Hanford formation beneath the

Interim Disposal Facility (IDF) at the Hanford Site. The objectives of this study are to

determine what geochemical conditions are required for chromate reduction by

Hanford formation sediments and which minerals are a source of Fe(II) for chromate

reduction under flow conditions. In addition to providing information about chromate

reduction specific to the Hanford Site, these experiments provide generally applicable

information about abiotic chromate reduction pathways in natural media.

122


4.2.1 Sediment Description

Sediment samples were obtained from borehole C3177, drilled during 2001 in

the 200 East Area, at the northeast corner of the Interim Disposal Facility (IDF) of the

Hanford Site. A detailed description of the core collection and geochemistry is

available in Horton et al. (25) and Walker et al. (26). Briefly, the < 2 mm size fraction

of five composite samples from depths of 14, 34, 46, 61, and 66 m below ground

surface in the Hanford flood sediments were used in this study. These sediment

samples are referred to as IDF-1 through IDF-5, in order of increasing depth. Horton

et al. (2003) determined that the pH of these sediments was 7.3, 7.4, 7.5, 7.5, and 7.6

for IDF-1 to -5, respectively, and that the total carbon content of each sediment sample

was less than 0.26 wt %. All samples are mineralogically similar, the sand-size

fraction dominated by quartz, plagioclase feldspar, potassium feldspar and mica (25),

and the clay-size fraction dominated by smectite, chlorite, illite, and kaolinite. Hanford

formation sediments contain trace amounts of magnetite, ilmenite, and biotite (19).

The total iron content of the samples was 4.4, 4.9, 4.3, 6.1, and 7.7 wt % for IDF 1 to

5, respectively, as determined by XRF (25).

4.2.2 Column Design and Reaction Conditions

All chemicals used were reagent grade or better. Various experiments were

conducted to determine factors responsible or limiting chromate reduction in

sediments from the Hanford site, which are summarized in Table 4.1. The porosity of

all packed columns was approximately 53%, with an average pore volume of 7.9 ± 0.1

cm3. Flow velocities upward through all columns were maintained at ca. 1.7 pore

123

volumes per day, and the effluent from all columns was sampled periodically. All

columns except the 10 M NaOH were run at least until the concentration of Cr(VI) in

the effluent was equal to the influent concentration (complete Cr(VI) breakthrough).

In order to simulate reaction conditions in arid environments, oxygen was not

excluded from any of the column systems, and the influent solutions were adjusted to

pH 8 to approximate the basic pH of these sediments (25). At the conclusion of each

experiment, breakthrough curves were obtained from each column using 0.2 mM

CaBr2 (pH 8); the sediments were then removed and stored at 10 ºC.

Chromate reduction under near-field conditions was investigated using two

columns packed with 15 g of dry IDF 5 sediment. The columns were saturated with

0.2 mM CaCl2 prior to the introduction of the feed solution, containing 0.2 mM

K2CrO4 in 10 M NaOH, into the experimental column. The control column’s influent

solution consisted of 0.2 mM K2CrO4 in 0.2 mM CaCl2 at pH 8.

The second set of experiments was conducted to investigate the influence of

various pre-treatments on chromate reduction. In this experiment, one column was

packed with 15 g of IDF 4 sediment that had been sonicated in Milli-Q water for 8 h.

Four other columns were packed with 15 g of dry IDF 4 sediment. After packing, one

of those four columns was left untreated as a control, while each of the remaining

three columns was pre-treated by pumping 10 pore volumes of one of three solutions

(10 mM oxalate, 0.5 M HCl, or 0.5 M NaOH) through the column. All five columns

were then equilibrated with 10 pore volumes of 0.2 mM CaCl2 at pH 8. Flow of a

solution of 0.2 mM K2CrO4 and 0.2 mM CaCl2, adjusted to pH 8, was then initiated.

124

A final series of experiments was conducted using three columns packed with

15 g of dry IDF 5 sediment. Three of these columns were treated with 10 pore

volumes of 0.5 M HCl, while an untreated column was used as a control. The four

columns were equilibrated with 10 pore volumes of 0.2 mM CaCl2 at pH 8. Flow of a

chromate-containing feed solution was then initiated. The feed solution for the control

column contained 0.2 mM K2CrO4, while the feed solutions of the acid-treated

columns contained either 0.2, 0.1, 0.02 mM K2CrO4. All feed solutions also contained

0.2 mM CaCl2 and were adjusted to pH 8 with 5 mM NaOH. Using a regression

model, the 95% confidence interval of the uptake curves was established at ± 0.0002

mmoles Cr g-1 sediment.

4.2.3 Solution and Solid Phase Measurements

Effluent pH was determined with an Orion ROSS electrode. Total soluble Cr,

Si, Al, Ca, and Fe were measured in samples passed through a 0.2 µm polymer filter,

using a Thermo Jarrell Ash IRIS inductively coupled plasma optical emission

spectrometer. Iron(II) and chromium(VI) were analyzed colorimetrically in a subset of

effluent samples by the ferrozine (27) and 1,5-diphenyl carbazide (28) methods,

respectively. Solid-phase chromium concentrations were determined by extraction of

0.5 g of sediment with 2 mL concentrated trace metal grade nitric acid for 48 h.

Samples were diluted and total Cr concentration was measured using ICP

spectrophotometry. Bromide concentrations were determined using ion

chromatography (Dionex, DX-500).

4.2.4 Solid-phase Micro-Analysis

125

Synchrotron micro-X-ray fluorescence (µ−XRF) mapping and micro-X-ray

absorption spectroscopy (µ−XAS) measurements were performed on beamline 10.3.2

at the Advanced Light Source (ALS, Berkeley, CA), GSE-CARS beamline 13-ID-C at

the Advanced Photon Source (APS, Argonne, IL), and beamline X26A at the National

Synchrotron Light Source (NSLS, Brookhaven, NY). Energy selection at the ALS and

NSLS was accomplished with a water-cooled Si(111) monochromator and at the APS

with a liquid N2-cooled Si (111) monochromator. The incident X-ray beam was

focused to a size of either 5 x 5 µm (ALS) or 5 x 10 µm (APS and NSLS), using two

Si mirrors in a Kirkpatrick-Baez geometry. Sediments from the batch reactions were

mounted on Kapton tape and attached to a x-y translation stage; the incident beam

intensity (Io) was measured with an in-line ionization chamber, and fluorescence yield

was measured using a multi-element Ge solid-state detector and normalized by I0. X-

ray fluorescence spectra were recorded on selected regions of the samples, on the basis

of elemental associations obtained from µ−XRF maps.

XAS data were processed using the SixPACK interface to (29) IFEFFIT (30).

XANES data were background-subtracted and normalized to a unit-edge step. After

background subtraction and normalization, the EXAFS data were extracted and k3-

weighted. Phase and amplitude functions for shell-by-shell fitting were generated

using FEFF 7 (31). A set of reference standards for Fe was utilized to perform linear

combination k3-weighted EXAFS spectral fitting, using SixPACK’s least squares

fitting module, which is a graphical interface for IFEFFIT’s minimization function

(30). Linear combination fitting routines were used to reconstruct the experimental

spectrum, in order to determine the relative percentages of mineral phases in the

126

sample. Shell-by-shell and linear combination fits were optimized by minimizing

reduced χ2 (32). The ratio of the Cr pre-edge peak to post-edge amplitude was used to

determine the concentration of Cr(VI) (10).

Μicro−XRD patterns were collected at ALS beamline 10.3.2 and NSLS

beamline X26A on select areas in transmission geometry using monochromatic

radiation (14, 16 or 17 keV) and a Bruker X-ray CCD camera. The resulting images

were processed using FIT2D (33). The sample-to-detector distance and geometric

corrections were calculated from the pattern of α-Al2O3. After these corrections were

applied, the 2D images were integrated radially to yield 1D powder diffraction

patterns that could then be analyzed using standard techniques. Peak identification and

background correction, including removal of the scattering from the Kapton tape, were

performed in JADE 6.5 (Materials Data, Inc., Livermore, CA).

127

4.3 RESULTS

4.3.1 Impact of Alkaline Conditions on Cr(VI) Retention Under Flow Conditions

To simulate mineral reactivity under near-field geochemical conditions

beneath a leaking tank, a feed solution of 10 M NaOH, containing 0.2 mM K2CrO4,

was utilized. Chromate removal is observed in the first pore-volume of effluent. After

40 pore volumes the effluent chromate concentration was 0.16 mM (Figure 4.1), and

the effluent chromate concentration remained below the influent concentration (0.2

mM) until the termination of the experiment. The removal of Cr(VI) does not appear

correlated with changes in aqueous Fe or Al concentration (Figure 4.1).

The presence of the characteristic Cr(VI) pre-edge feature (10,21) in XANES

spectra reveals Cr(VI) in all identified Cr hot spots ( 4.2, NaOH 1-3)–which is

confirmed by the presence of Cr-O shells at 1.6 Å in the EXAFS spectra from the

same spots (Figure 4.2, Table 4.2). The points of high Cr concentration analyzed in

XRF maps of the sediment from the 10 M NaOH column contain 13 to 26% Cr(VI)

(Figure 4.2, NaOH 1-3). Analysis of areas of high Cr concentration with chromium

EXAFS spectroscopy shows a Cr-Cr(Fe) structure similar to α/β-MeOOH (Me = Fe or

Cr) (Figure 4.2, Table 4.2), with scattering contributions from Al(Si) at ~3.1 Å,

consistent with aluminum substitution, and Cr(Fe) at ~3.5 Å. Comparatively, goethite

(α-FeOOH) has Fe-Fe shells at 3.01, 3.29 and 3.43 Å (34) and akaganeite (β-FeOOH)

at 3.03, 3.34, 3.50, and 3.55 Å (35). Lizardite and portlandite (CaOH) were

commonly associated with high chromium concentrations (4.3 A-D), and magnetite

was detected in one sample (Figure 4.3C). A chromium-rich grain of portlandite was

128

isolated (Figure 4.3B, 4.4A), in which 20% of the chromium resides in the hexavalent

state (data not shown).

Dissolution of ferrous iron-bearing minerals (antigorite, biotite, lizardite, and

magnetite) induced by the extreme alkalinity leads to the dispersal of Fe(II) and

chromate reduction/retention throughout the mineral matrix. The reduction of Cr(VI)

throughout the sediment is exemplified by a detailed examination of the chromium

distribution in a hand-picked mica grain (Figure 4.4A). XANES spectra obtained from

several points in the mica grain showing Cr(III), and line scans across the particle

indicate that Cr is distributed throughout the mineral surface (Figure 4.4A). Despite

the preponderance of Cr(III), Cr(VI) was noted within the sample, in part through

sequestration within portlandite (Figures 4.3B and 4.4A). As a result of the multiple

retention mechanisms, 1.1 x 10 -3 mmol Cr g-1 sediment is retained in the column

reacted with a feed solution containing 10 M NaOH (Table 4.1).

4.3.2 Influence of Sediment Pre-Treatment on Chromate Retention Under Flow

Conditions

The reduction of chromate by non-treated IDF 5 sediment (control) was

investigated under far-field conditions (those representative of the ‘native’

geochemistry) using a feed solution of 0.2 mM K2CrO4 in 0.2 mM CaCl2 adjusted to

pH 8. Relative to the Br- tracer, retardation of chromate breakthrough is not observed

(Figure 4.5). Several pre-treatments were then used to target different types of surface

coatings that may inhibit electron transfer from underlying Fe(II) to Cr(VI).

Sonication was used to simulate physical abrasion, 0.5 M NaOH to promote

dissolution of silicate phases, 10 mM oxalate to provide a complexant to dissolve

129

Fe(III) and other metal (hydr)oxide phases, and 0.5 M HCl to induce acidic dissolution

of metal oxide and carbonate phases. Sonication, 10 mM oxalate, and 0.5 M NaOH

pre-treatments did not result in removal of Fe(II) from the sediment; however, 0.16

mmoles Fe(II) g-1 sediment are removed with the acid treatment. Relative to the

control (no pre-treatment) and Br- tracer, retardation of Cr(VI) breakthrough occurred

only after treatment with 0.5 M HCl (Figure 4.5). Retained chromium after this

treatment is exclusively in the trivalent state (Figure 4.2, ATC 1-3). As a consequence

of Cr reduction in the 0.5 M HCl-treated column, nearly two orders of magnitude

more chromium are deposited than after the other pre-treatments (Table 4.1) The

effluent pH of the 0.5 M HCl treated column was initially 6.5 and gradually increased

to 7.5 over the duration of the experiment. The effluent pH of the other four columns

was 7.7 ± 0.2 for the duration of the experiment.

Three influent chromate concentrations (0.2, 0.1, and 0.02 mM) were utilized

to determine the effect of chromate concentration on uptake behavior of 0.5 M HCl

treated sediments. Chromate breakthrough occurs after the elution of 20 pore volumes

of 0.2 mM Cr(VI), 40 pore volumes of 0.1 mM Cr(VI), or 100 pore volumes of 0.02

mM Cr(VI) (Figure 4A.1). A comparison of the breakthrough trends on a mass-

introduced to mass-retained basis reveals two distinct regions of Cr uptake by the

sediment. An initial region occurs after introducing between 0 to 0.007 mmoles

Cr(VI). A second region develops after 0.007 mmoles Cr(VI) have been introduced

(Figure 4A.1). The 0.2 mM Cr(VI) feed solution rapidly breaks through, while the

lower feed solution concentrations of 0.1 and 0.02 mM Cr(VI) result in continued

Cr(VI) attenuation (Figure 4A.1), which is confirmed by nitric acid extraction of the

130

solid at the conclusion of the experiment (Table 4.1). The effluent pH of the three acid

pre-treated columns was initially 6.4 and gradually increased to 7.6 by the end of the

experiment.

At all three influent concentrations, chromate reduction occurs concurrently

with the presence of aqueous Fe(II) (Figure 4A.2). Once the aqueous Fe(II)

concentration decreases below 0.005 mM, chromate reduction ceases. Aqueous Fe(II)

concentrations are below 0.002 mM prior to the introduction of the chromate-

containing feed solution (CaCl2 wash; data not shown).

4.3.3 Reactive Solid Phases After Acid Treatment

High concentrations of reduced chromium associated with mafic minerals,

including lizardite, antigorite (Figure 4.3E), and nimite, a ferrous iron-bearing phyllite

(Figure 4.3F), were noted with µ-XRD. Chromium(III) was observed along the edges

of mica grains (Figure 4.4B). Two iron µ-EXAFS spectra, obtained from areas on

XRF maps of high Cr concentrations, are well fit by linear combinations of

ferrihydrite-nontronite and ferrihydrite-biotite (Figure 4A.3).

EXAFS spectroscopy was used to probe the local structure of the mixed Cr/Fe

(hydr)oxides precipitated at points of high Cr concentration in solid samples from the

acid-treated column. Similar to the alkaline system, shell-by-shell fitting reveals that

the short-range order of mixed Cr/Fe precipitates in three spots of high-Cr

concentration in the sediment from the acid-treated column (ATC) is indicative of an

α/β-MeOOH-like structure (Figure 4.2 and Table 4.2, ATC 1-3).

131

4.4 DISCUSSION

During weathering in arid, oxic subsurface environments, Fe(II)-bearing

mineral reactivity, with respect to chromate, becomes compromised by a variety of

weathering rinds and surface coatings. Zachara et al. (36), for example, observed

ferric (hydr)oxide rinds on biotite grains obtained from the Hanford formation

sediments. In fact, using XPS analysis we observe a ferric iron layer overlying

magnetite separated from IDF 5 sediments (Figure 4A.4). The presence of such an

oxidized layer diminishes the rate of chromate reduction, due to limited transfer of

electrons [or Fe(II)] from the underlying unoxidized mineral to the aqueous interface,

as illustrated for synthetic magnetite (6). Owing to the passivating layer on Fe(II)-

bearing minerals, physical abrasion from fast shaking speeds is required for chromate

reduction within batch reactions (Figure 4A.5) and acid or strong base treatment is

needed for reduction within column experiments performed here.

Under near-field conditions (10 M NaOH), the extreme pH induces base

catalyzed dissolution of both surface coatings (e.g., Fe-hydroxides) and Fe(II)-bearing

silicate minerals themselves. Under these conditions, the column having a 10 M

NaOH feed solution retains more chromium than either the acid pre-treated columns

or the control columns (Table 4.1). Additionally, chromate retention is still occurring

after the elution of 60 pore volumes (Figure 4.1) due to surface activation and

extensive mineral dissolution as noted by extensive Cr retention and reduction

throughout the mineral matrix. In this hyperalkaline system, XRF maps, combined

with localized XRD analysis, reveal the association of high local chromium

concentrations with several Fe(II)-bearing minerals, including antigorite, lizardite, and

132

magnetite (Figure 4.3). EXAFS analysis of localized Cr(III) phases reveals that they

have the short-range order consistent with an α/β MeOOH structure (Figure 4.2 and

Table 4.2), suggesting that homogeneous Cr(VI) reduction by Fe(II) (aq) is occurring,

in addition to the heterogenous Cr(VI) reduction observed on mica grains (Figure 4.4).

The presence of portlandite is also detected in every instance of Cr deposition.

Analysis of an isolated portlandite grain indicates Cr in association with the grain,

20% of which is in the hexavalent oxidation state (Figures 4.3 and 4.4). XANES

analysis of three additional areas of high chromium concentration reveal Cr(VI)

concentrations ranging from 13% to 26%. The presence of solid-phase Cr(VI)

suggests that in addition to chromate reduction and subsequent precipitation, other

retention mechanisms are partially responsible for the large amount of chromium

retained under simulated near-field (i.e., hyperalkaline) conditions. Reduction of

Cr(VI) to Cr(III) is likely inducing rapid nucleation and precipitation of portlandite,

thereby trapping Cr(VI) in the solid phase. The retention of Cr(VI) at the Hanford Site

has also been observed in sediment from a contaminated core obtained from beneath

tank SX-108, which contained Cr(VI) in addition to discrete Cr(III) mineral phases

(18).

The inhibition of chromate reduction by oxidized surface coatings is

emphasized by the results of the pre-treatment experiments, where chromate reduction

is only observed after the acidic dissolution of Fe(III) (hydr)oxide surface coatings

(Figure 4.5). Similar to the alkaline column, localized high-chromium concentrations

were associated with Fe(II)-bearing minerals including biotite, antigorite, lizardite,

nimite (a ferrous-bearing phyllosilicate) (Figure 4.3 and Figure 4A.3); however, less

133

chromium was retained by the sediment (Table 4.1). The relationship of dissolved Fe

and Cr in column-effluent (Figure 4A.2) suggests that initial Cr(VI) reduction is being

mediated by soluble Fe(II). Chromium(VI) concentrations are anticorrelated with

Fe(II) concentrations, and Cr(VI) is not detected in the effluent until Fe(II) decreases

below 0.005 mM. Acidic pre-treatment results in the activation of a limited number of

Fe(II) surface sites (Figure 4.4B), and as chromate reduction proceeds, it becomes

self-inhibited by the oxidation of the underlying mineral phase and precipitation of

mixed Cr(III)/Fe(III) (hydr)oxide phases (Figure 4.2, ATC 1-3, and Figure 4A.3). A

feed solution containing 0.2 mM Cr(VI) results in more rapid surface passivation than

the lower influent chromate concentrations of 0.1 and 0.02 mM (Figure 4A.1);

however, due to competition of chromate for Fe(II) with molecular oxygen in aerated

systems, less chromate is reduced in the 0.02 mM column than the 0.1 mM column—

owing, in part, to a 5-fold longer reaction period to achieve comparable Cr quantities

(Table 4.1 and Figure 4A.1).

134

4.5 ENVIRONMENTAL IMPLICATIONS

Attenuation of chromate in arid environments will be governed by the

reactivity of Fe(II)-bearing mineral phases, which is controlled by mineral solubility

and surface reactivity. Chromate reduction by native sediments of the Hanford

formation under far-field conditions (pH < 8) is inhibited by passivating surface

coatings and low mineral solubility. Acidic dissolution of the Fe(III) (hydr)oxide

layer restores the redox activity of the Fe(II) mineral surface and results in chromate

reduction; however, formation of newly oxidized surface layers further limits Cr(VI)

reduction. Conversely, near-field conditions (pH > 14) cause extensive base-induced

mineral dissolution and Fe(II) release, resulting in chromate retention throughout the

mineral matrix. In addition to magnetite, Fe(II)-bearing silicates such as biotite,

lizardite, and antigorite also play an important role in chromate reduction. Serpentine

subgroup minerals (i.e., lizardite and antigorite) are quite common and widely

distributed in mafic rocks (their quantification, unfortunately, is compromised by

kaolin minerals (37)). Up to 60% of the total iron in lizardite occurs in the reduced

state; likewise, 90% of the total iron in antigorite occurs in the reduced state (38,39).

This accounts for 3% and 40% of the octahedrally-coordinated cations in lizardite and

antigorite, respectively. Therefore, these minerals appear to represent an important

reservoir of Fe(II) for chromate reduction. Under conditions that induce silicate

mineral dissolution, such as the near-field geochemical environment at the Hanford

Site, chromate reduction will thus be enhanced by release of Fe(II). Thus, to

thoroughly evaluate chromate reduction in arid environments, viable reductants under

the relevant geochemical conditions must be considered. Given the paucity of organic

135

matter within such systems, primary Fe(II)-bearing mineral phases will likely be the

dominant reductant. The surface reactivity and solubility of such minerals thus need

to be appreciated in deciphering the rate and extent of chromate reduction.

136

4.6 ACKNOWLEDGMENTS

We are grateful to Royce Sparks and Molly Pace for laboratory assistance with

the batch experiments. We also thank R. Jeffrey Serne of Pacific Northwest National

Laboratory for providing the IDF sediments and Kathleen Beman for editorial input.

This research was supported by the U.S. Department of Energy’s EMSP Program

(grant number DE-FG07-02ER63516). Portions of this work were performed at

GeoSoilEnviroCARS (Sector 13), Advanced Photon Source (APS), Argonne National

Laboratory. GeoSoilEnviroCARS is supported by the National Science Foundation-

Earth Sciences (EAR-0217473), Department of Energy-Geosciences (DE-FG02-

94ER14466) and the State of Illinois. Use of the APS was supported by the U.S.

Department of Energy, Basic Energy Sciences, Office of Energy Research, under

Contract No. W-31-109-Eng-38. The Advanced Light Source is supported by the

Director, Office of Science, Office of Basic Energy Sciences, Materials Sciences

Division, of the U.S. Department of Energy under Contract No. DE-AC03-76SF00098

at Lawrence Berkeley National Laboratory. Research was also carried out at the

National Synchrotron Light Source, Brookhaven National Laboratory, which is

supported by the U.S. Department of Energy, Division of Materials Sciences and

Division of Chemical Sciences, under Contract No. DE-AC02-98CH10886.

137


(1) Fendorf, S. Surface reactions of chromium in soils and waters. Geoderma 1995, 67, 55-71. (2) Sass, B. M.; Rai, D. Solubility of amorphous chromium(III)-iron(III) hydroxide solid solutions. Inorg. Chem. 1987, 26, 2228-2232. (3) Fendorf, S.; Li, G. Kinetics of chromate reduction by ferrous iron. Environ. Sci. Technol. 1996, 30, 1614-1617. (4) Buerge, I. J.; Hug, S. J. Influence of mineral surfaces on Cr(VI) reduction by iron(II). Environ. Sci. Technol. 1999, 33, 4285-4291. (5) Deng, B.; Stone, A. T. Surface-catalyzed Cr(VI) reduction: Reactivity comparisons of different organic reductants and different oxide surfaces. Environ. Sci. Technol. 1996, 30, 2484-2494. (6) Peterson, M. L.; White, A. F.; Brown, G. E.; Parks, G. A. Surface passivation of magnetite by reaction with aqueous Cr(VI): XAFS and TEM results. Environ. Sci. Technol. 1997, 31, 1573-1576. (7) White, A. F.; Peterson, M. L. Reduction of aqueous transition metals on the surfaces of Fe(II)-containing oxides. Geochim. Cosmochim. 1996, 60, 3799-3814. (8) Bond, D. L.; Fendorf, S. Kinetics and structural constraints of chromate reduction by green rusts. Environ. Sci. Technol. 2003, 27, 2750-2757. (9) Eary, L. E.; Rai, D. Kinetics of chromate reduction by ferrous ions derived from hematite and biotite at 25 oC. Am. J. Sci. 1989, 289, 180-213. (10) Patterson, R. R.; Fendorf, S.; Fendorf, M. Reduction of hexavalent chromium by amorphous iron sulfides. Environ. Sci. Technol. 1997, 31, 2039-2044. (11) Fendorf, S.; Wielinga, B. W.; Hansel, C. M. Chromium transformations in natural environments: The role of biological and abiological processes in chromium(VI) reduction. Int. Geol. Rev. 2000, 42, 691-701. (12) James, B. J.; Bartlett, R. J. Behavior of chromium in soils: V. Fate of organically complexed Cr(III) added to soil. J. Environ. Qual. 1983, 12, 169-172. (13) Puzon, G. J.; Peterson, J. N.; Roberts, A. G.; Kramer, D. M.; Xun, L. A bacterial flavin reductase system reduces chromate to a soluble chromium(III)-NAD+ complex. Biochemical and Biophysical Research Communications 2002, 294, 76-81. (14) Brauer, S. L.; Hneihen, A. S.; Wetterhahn, K. E. Chromium(VI) forms thiolate complexes with gamma-glutamylcysteine, N-acetcylcysteine, cysteine, and the methyl ester of n-acetylcysteine. Inorg. Chem. 1996, 35, 373. (15) Kendelewicz, T.; Liu, P.; Doyle, C. S.; Brown, G. E. Spectroscopic study of the reaction of aqueous Cr(VI) with Fe3O4(111) surfaces. Surface Science 2000, 469, 144-163. (16) Poston, T. M.; Hanf, R. W.; Dirkes, R. L.; Morasch, L. F. "Hanford site environmental report for calendar year 2000," Pacific Northwest National Laboratory, 2001. (17) Jones, T. E.; Watrous, R. A.; G.T., M. "Inventory estimates for single-shell tank leaks in S and SX tank farms," CH2M Hill Hanford Group, Inc., 2000.

138

(18) Zachara, J. M.; Ainsworth, C. C.; Brown, G. E.; Catalano, J. G.; McKinley, J., P.; Qafoku, O.; Smith, S. C.; Szecsody, J. E.; Traina, S. J.; Warner, J. A. Chromium speciation and mobility in a high level nuclear waste vadose zone plume. Geochim. Cosmochim. 2004, 68, 13-30. (19) Zachara, J. M.; Ainsworth, C. C.; Brown, G. E.; Catalano, J. G.; McKinley, J., P.; Qafoku, O.; Smith, S. C.; Szecsody, J. E.; Traina, S. J.; Warner, J. A. Chromium Speciation and Mobility in a High Level nuclear Waste Vadoes Zone Plume. Geochim. Cosmochim. 2004, 68, 13-30. (20) Doyle, C. S.; Kendelewicz, T.; Brown, G. E. Inhibition of the reduction of Cr(VI) at the magnetite-water interface by calcium carbonate coatings. Applied Surface Science 2004, 230, 260-271. (21) Peterson, M. L.; Brown, G. E.; Parks, G. A.; Stein, C. L. Differential redox and sorption of Cr(III/VI) on natural silicates and oxide minerals: EXAFS and XANES results. Geochim. Cosmochim. Acta. 1997, 61, 3399-3412. (22) Ilton, E. S.; Veblen, D. R. Chromium sorption by phlogopite and biotite in acidic solutions at 25 C: Insights from X-ray photoelectron and electron microscopy. Geochim. Cosmochim. Acta. 1994, 58, 2777-2788. (23) He, Y. T.; Traina, S. J. Cr(VI) reduction and immobilization by magnetite under alkaline pH conditions: The role of passivation. Environ. Sci. Technol. 2005, 39, 4499-4504. (24) He, Y. T.; Chen, C. C.; Traina, S. J. Inhibited Cr(VI) reduction by aqueous Fe(II) under hyperalkaline conditions. Environ. Sci. Technol. 2004, 38, 5535-5539. (25) Horton, D. G.; Schaef, H. T.; Serne, R. J.; Brown, C. F.; Valenta, M. M.; Vickerman, T. S.; Kutnyakov, I. V.; Baum, S. R.; Geiszler, K. N.; Parker, K. E. "Geochemistry of Samples from Borehole C3177 (299-E24-21)," Pacific Northwest National Laboratory, 2003. (26) Walker, L. D. "Borehole summary report for the 2001 ILAW site characterization well," Bechtel Hanford, Inc., 2001. (27) Stookey, L. L. A new spectrophotometric reagent for iron. Anal. Chem. 1970, 42, 779-781. (28) Bartlett, R.; James, B. J. Behavior of chromium in soils: III. Oxidation. J. Environ. Qual. 1979, 8, 31-35. (29) Webb, S. M. Sixpack: A graphical user interface for XAS analysis using IFEFFIT. Physica Scripta 2005, T115, 1011-1014. (30) Newville, M. IFEFFIT: interactive XAFS analysis and FEFF fitting. J. Synchrotron Radiation 2001, 8, 322-324. (31) Ankudinov, A. L.; Rehr, J. J. Relativistic calculations of spin-dependent X-ray-absorption spectra. Phys. Rev. B: Condens. Matter 1997, 15, R1712-R1715. (32) Newville, M.; Ravel, B.; Haskel, D.; Rehr, J. J.; Stern, E. A.; Yacoby, Y. Analysis of multiple-scattering XAFS data using theoretical standards. Physica B 1995, 208 & 209, 154-155. (33) Hammersley, A. P. "FIT2D: An Introduction and Overview," European Synchrotron Radiation Facility, 1997.

139

(34) Hazemann, J. L.; Berar, J. F.; Manceau, A. Rietveld studies of the aluminum-iron substitution in synthetic goethite. Physics and Chemistry of Minerals 1991, 19, 25-38. (35) Post, J. E.; Buchwald, V., F. Crystal structure refinement of akaganeite. American Mineralogist 1991, 76, 272-277. (36) Zachara, J. M.; Smith, S. C.; McKinley, J., P.; Serne, J. N.; Gassman, P. L. Sorption of Cs+ to micaceous subsurface sediments from the Hanford site. Geochim. Cosmochim. Acta. 2002, 66, 193-211. (37) White, G. N.; Dixon, J. B. In Soil Mineraology with Environmental Applications; Dixon, J. B., Schulze, D. G., Eds.; Soil Science Society of America: Madison, WI, 2002; Vol. 7, pp 389-414. (38) Votyakov, S.; Chashchukhin, I.; Bykov, V.; Mironov, V. Behavior of Fe ions in minerals of ultrabasites during serpentinization. Geochem. Int. 1993, 30, 75-85. (39) Fuchs, Y.; Linares, J.; Mellini, M. Mossbauer and infrared spectrometry of lizardite-1T from Monte Fico, Elba. Phys. Chem. Minerals 1998, 26, 111-115.

140

Table 4.1 Experimental conditions used in columns to investigate chromate reduction by Hanford formation sediment. After pre-treatment, and prior to the introduction of the chromate influent solution, all columns were equilibrated with 10 pore volumes of 0.2 mM CaCl2 at pH 8.

Column ID Pre-treatment Influent Solution Influent pH

Cr Retained (mmol Cr g-1

sediment)

Control None 0.2 mM CaCl2, 0.2 mM K2CrO4

8 1.0 x 10-5

10 M NaOH None 10 M NaOH, 0.2 mM K2CrO4

>14 1.1 x 10-3

Sonicated 8 h sonication 0.2 mM CaCl2, 0.2 mM K2CrO4

8 1.0 x 10-5

10 mM Oxalate 10 PV* 10 mM oxalate

0.2 mM CaCl2, 0.2 mM K2CrO4

8 1.1 x 10-5

Acid-treated 10 PV 0.5 M HCl


8 8.3 x 10-4

Base-treated 10 PV 0.5 M NaOH


8 9.1 x 10-6

0.2 mM Cr(VI) 10 PV 0.5 M HCl


8 7.0 x 10-4

0.1 mM Cr(VI) 10 PV 0.5 M HCl


8 1.0 x 10-3

0.02 mM Cr(VI)

10 PV 0.5 M HCl


8 8.7 x 10-4

* PV = Pore Volumes

141

Table 4.2 Structural parameters derived from least squares fits to raw k3-weighted Cr-EXAFS spectra at three areas of high Cr concentration in sediments from acid pre-treated (ATC 1 - 3) and 10 M NaOH (NaOH 1 - 3) column experiments. Coordination number (CN), interatomic distance (r), and Debye-Waller factor (σ2) were obtained by fitting data with theoretical phase and amplitude functions. Fits were performed over a k range of 3.5 – 10 Å-1. The reduced χ2 values for each fit were between 0.6 and 0.9. Distances reported in the table differ from those in the radial distribution function (RDF) because the latter were not corrected for phase shift. Estimated errors at 95% confidence interval from the lease squares fit are given in parentheses.

ATC 1 ATC 2 ATC 3 NaOH 1 NaOH 2 NaOH 3 Shell r(Å) CN σ2 r(Å) CN σ2 r(Å) CN σ2 r(Å) CN σ2 r(Å) CN σ2 r(Å) CN σ2 Cr-O NA NA NA NA NA NA NA NA NA 1.64(1) 2.9(2) 0.003(2) 1.66(9) 3.3(4) 0.015(4) 1.62(4) 2.9(3) 0.012(4) Cr-O 1.98(1) 5.4(6) 0.002(2) 1.99(1) 5.8(4) 0.002(3) 1.99(1) 5.7(3) 0.003(2) 2.00(1) 4.8(2) 0.003(1) 1.99(1) 5.3(4) 0.003(5) 1.99(1) 5.0(3) 0.008(3) Cr-

Al(Si) NA NA NA NA NA NA NA NA NA 3.08(1) 1.0(2) 0.002(3) 3.14(5) 0.8(3) 0.002(3) NA NA NA

Cr-Cr(Fe) 3.18(8) 0.4(5) 0.002(1) 3.11(4) 1.9(9) 0.010(3) 3.19(1) 2.4(6) 0.004(3) 3.48(2) 2.1(3) 0.013(2) 3.45(5) 2.8(4) 0.010(4) 3.48(5) 2.8(4) 0.007(4)

Cr-Cr(Fe) NA NA NA NA NA NA NA NA NA 4.68(4) 1.1(2) 0.003(1) NA NA NA 3.64(8) 1.7(2) 0.007(3)

Cr-Cr(Fe) NA NA NA NA NA NA NA NA NA 5.20(4) 0.5(2) 0.003(1) NA NA NA NA NA NA

142

0

10

20

Con

cent

ratio

n (m

M)

0.150

0.175

0.200

0.225

Pore Volumes0 20 40 60

0.00

0.75

1.50

2.25

Aluminum

Chromium

Iron

Figure 4.1 Effluent metal concentration for control (●) and 10 M NaOH columns (▲).

143

k (Å-1)3 4 5 6 7 8 9 10 11

χ (k)

*k3

R (Å)0 1 2 3 4 5 6

Four

ier T

rans

form

Mag

nitu

deEnergy (eV)

5985 6000 6015 6030

ATC 2

ATC 3

NaOH 1

NaOH 2

NaOH 3Nor

mal

ized

Flu

ores

cenc

e In

tens

ity

ATC 1

Figure 4.2 Chromium K-edge XANES (left), EXAFS (center), and Fourier-transformed EXAFS (right) spectra from three areas of high Cr concentration in sediments from acid pre-treated (ATC 1 - 3) and 10 M NaOH (NaOH 1 - 3) column experiments. EXAFS data (center; solid lines) were fit (center, dotted lines) over k = 3.5 – 10 Å-1 (fit data provided in Table 4.2). Fourier transforms of data (right, solid lines) and fits (right, dotted lines) were calculated over the same k range.

144

Figure 4.3 XRD patterns obtained from areas of high chromium concentration in sediments from 10 M NaOH (a-d) and acid pre-treated (e-f) columns. Minerals identified in the patterns include portlandite (P), lizardite (L), magnetite (M), quartz (Q), nimite (N), antigorite (An), and albite (Al). Unidentifiable single scattering peaks are indicated with an “*”.

145

Figure 4.4 Microanalysis of mica grains from 10 M NaOH (A) and acid pre-treated (B) columns. In XRF maps, Cr distribution is shown in red and Fe in green.

146

Pore Volumes0 5 10 15 20 25 30

Cr (

C/C

0)

0.0

0.2

0.4

0.6

0.8

1.0

1.2

No TreatmentSonication10 mM Oxalate0.5 M HCl0.5 M NaOHBr- Tracer

Figure 4.5 Effect of sediment pre-treatment on Cr(VI) breakthrough. The sediment was pre-treated with 10 pore volumes of the indicated treatment, followed by equilibration with 10 pore-volumes 0.2 mM CaCl2 at pH 8, prior to the introduction of the 0.2 mM K2CrO4 solution.

147

APPENDIX 4A


148

Figure 4A.1. Mass of Cr(III) retained by acid treated column solids based on the mass of Cr(VI) introduced into the column. Breakthrough curves for the control and three acid treated columns are noted within the inset.

149

0.00

0.05

0.10

0.15

0.20

0.25

0.000

0.005

0.010

0.015

0.020

Cr(V

I) (m

M)

0.00

0.03

0.06

0.09

0.12

Fe(II

) (m

M)

0.00

0.01

0.02

0.03

0.1 mM K2CrO4

0.2 mM K2CrO4

Cr(VI)

Fe(II)

Cr(VI)

Fe(II)

Pore Volume

0 10 20 30 40 500.000

0.005

0.010

0.015

0.020

0.000

0.015

0.030

0.045Cr(VI)

Fe(II)

0.02 mM K2CrO4

Figure 4A.2. Effluent concentrations for 0.5 M HCl treated columns at various influent chromium concentrations.

150

k (Å-1)2 4 6 8 10 12 14

χ (k)

*k3

Batch - IDF 4

Batch - IDF 5

ATC 1, 0.2 mM K2CrO4

ATC 2, 0.2 mM K2CrO4

Goethite Standard

Ferrihydrite Standard

Magnetite Standard

Biotite Standard

Nontronite Standard

Sample % Ferrihydrite % Goethite % Magnetite % Nontronite % Biotite xv

2 IDF 4 34 66 NA NA NA 0.7 IDF 5 45 NA 55 NA NA 1.2 ATC 1 23 NA NA 77 NA 0.1 ATC 2 53 NA NA NA 47 0.4

Figure 4A.3. Experimental Fe-EXAFS (solid) and linear combination fits (dotted) from batch and acid-treated column (ATC) experiments. Spectra were collected at points of high Cr concentration identified in XRF maps. Fe mineralogical composition of points of high Cr concentration obtained linear combination fitting are tabulated. Ferrihydrite and goethite are indicative of mixed Fe(III)/Cr(III) (hydr)oxide phases.

151

Figure 4A.4. X-ray photoelectron spectra of magnetic separates. Prior to sputtering, the Fe(2p3/2) peak at 711.8 eV is indicative of Fe(III) and appears similar to ferrihydrite (1). After three minutes sputtering, peaks are apparent at 710.6 and 706.8 eV, which are indicative of Fe(II) and Fe(III) within magnetite (1). (1) Harvey, D. T.; Linton, R. W. Chemical characterization of hydrous ferric oxides by X-ray photoelectron spectroscopy Analytical Chemistry 1981, 53, 1684-1688.

Binding Energy (eV)

700705710715720

711.8

710.6

706.8

(a)

(b)

152

Time (hours)

0 200 400 600 800 1000

C/C

0

0.0

0.2

0.4

0.6

0.8

1.0

1.2

IDF 1 (140 rpm)IDF 2 (140 rpm)IDF 3 (140 rpm)IDF 4 (140 rpm)IDF 5 (140 rpm)IDF 5 (80 rpm)

Figure 4A.5. Batch Experiments. Four grams of each sediment were added to 8 mL of oxygen-free 0.1 M NaCl in a 20 mL serum vial and adjusted to pH 8 with dilute NaOH. Reactions were initiated, with the addition of K2CrO4, to a final concentration of 0.2 mM. The serum vials were placed on a rotary shaker at 140 rotations per minute (rpm), except for one experiment (with IDF-5) in which the shaker speed was 80 rpm. Samples were taken at various time points during the reaction sequence. Sets of the five sediments were sacrificed at 7 and 15 d for analysis by X-ray absorption spectroscopy (XAS). Results: At shaking speeds of 140 rpm, aqueous chromate is depleted; however, when shaking is decreased to 80 rpm, little chromium is removed from solution, even after more than 1400 h. At the higher shaking speed, sample IDF 3 removes 2 x 10-4 mmol Cr g-1 sediment, while the other four sediments remove 4 x 10-4 mmol Cr g-1 sediment. At the lower shaking speed, less than 1 x 10-6 mmol Cr g-1 sediment is removed. Since chromate tends to weakly adsorb to oxidized sediments (2), we suspect reduction to Cr(III) at the higher shaking speeds. Indeed, Cr(VI) is not detected in any of the solid phases from the batch experiments reacted at 140 rpm.

153

CHAPTER 5

Conclusions

Matthew Ginder-Vogel

Department of Geological and Environmental Sciences, Stanford University, Stanford, CA 94305

154

5.1 CONCLUSIONS

Despite centuries of study, fundamental biogeochemical processes controlling

the dynamics of contaminants within complex environmental media remain poorly

defined. In addressing the fate of environmental contaminants, one must consider the

complete system, rather than solely considering a simplified fraction of the soil.

However, in natural environments, elemental cycling may be mediated by a variety of

processes, such as adsorption/desorption and oxidation or reduction (abiotic or biotic);

therefore, the focus of my research has been the identification and characterization of

biotic and abiotic pathways of metal transformation in soils and sediments.

Within this thesis, the abiotic (chemical) and biotic (bacterial) controls on

uranium and chromium mobility within natural systems are examined in order further

our understanding of the biogeochemical processes controlling contaminant mobility

in complex environmental systems. The hexavalent oxidation state of chromium and

uranium is the most susceptible to environmental transport. However, in their reduced

states, Cr(III) and U(IV), their aqueous solubility, and hence their mobility, is limited.

It is therefore vital to identify environmental processes affecting the oxidation state of

these elements in order to further our understanding of their environmental behavior.

The environmental chemistry of chromium is relatively well-understood; however, as

yet, relatively little is known about biogeochemical uranium cycling in real-world

environments (Figure 5.1), particularly biogenic uraninite’s long-term stability.

The reduction of soluble U(VI) to sparingly soluble U(IV) (uraninite) by

bacterial respiration in anaerobic environments will dramatically impact its mobility;

however, the geochemical factors affecting the rate of uranium bacterial uranium

155

reduction and the stability of biologically precipitated uraninite remain poorly

characterized, particularly within complex environmental systems (Figure 5.1). For

example, Fe(III) (hydr)oxides are ubiquitous in nature and can impact the mobility of

uranium through sorption (1-3), reduction (4), and oxidation (Chapters 2 and 3)

reactions. Additionally, Fe(III) (hydr)oxides inhibit bacterial uranium reduction by

acting as competitive terminal electron acceptors (5,6). Aqueous geochemical

conditions also greatly impact the environmental mobility of uranium. The presence

of small amounts of dissolved Ca2+ severely inhibits bacterial uranium reduction (5,7),

although the presence of Fe(III) (hydr)oxides slightly relieves this inhibition through

calcium sorption. Changes in carbonate concentration and pH also have a large impact

on uranium mobility through sorption processes (3,8,9) and the oxidative dissolution

of uraninite by Fe(III) (hydr)oxides) (Chapters 2 and 3) and molecular oxygen (10,11)

(Figure 5.1). The oxidation state of uranium can also be affected by other common soil

constituents, including reduction by sulfide (12,13) and micas (14) or oxidation

Mn(III, IV) (hydr)oxides (15).

Many of the above studies have been performed solely in laboratory

environments and frequently in static systems far removed from environmental

conditions. In order to further our understanding of the importance of bacterial

uranium reduction in controlling the mobility of uranium in natural environments,

these studies need to be expanded to more closely approximate real-world conditions.

Therefore, we need to take what we have learned about the complexity of

biogeochemical uranium cycling through field studies, such as at the Oak Ridge

156

(16,17) and the Old Rifle (18) field sites, and identify factors that may enhance the

stability of biologically precipitated uraninite in natural environments.

In particular, we need to understand the role of physical and chemical

heterogeneity in uranium biogeochemical cycling. Although uraninite oxidation by O2,

Mn(IV), Fe(III), or NO2- is thermodynamically favorable (Figure 5.2), specific

geochemical conditions will determine the rate and extent of these reactions (Chapters

2 and 3). For example, since uraninite dissolution is the rate controlling step in

uraninite oxidation by Fe(III) (hydr)oxides (Chapter 3), conditions that limit the

solubility of uraninite, such as increasing the uraninite’s crystallinity and maintaining

a neutral pH, low bicarbonate environment, will necessarily reduce the oxidation rate.

Alternatively, the spatial isolation of bacterially reduced uranium from potential

oxidants (currently being investigated at the Oak Ridge Field Site) may be one method

to limit long-term uranium mobility. In addition to uranium sequestration as reduced

phases, low-solubility uranyl-phases (e.g. U(VI)-phosphates) or incorporation into

other low-solubility phases (e.g. Fe-(hydr)oxides (19)) may also be a pathway to limit

uranium mobility. However, as is the case with uraninite, slight changes in

geochemical conditions may result in the dissolution of these phases.

Our research at the Oak Ridge, TN FRC has demonstrated that even in

severely contaminated environments, stimulation of bacterial uranium reduction

results in the precipitation of uraninite and in dissolved uranium concentrations below

the Environmental Protection Agency’s Maximum Contaminant Level (MCL) of

0.126 µM (Chapter 4). Despite the presence of ~30% U(VI) in sediment retrieved after

bioremediation at the FRC, only a fraction of a percent is easily exchangeable

157

(Chapter 1). Ultimately, the microbial remediation may be successful at limiting

uranium mobility in reducing environments, but the threat for oxidation and

remobilization will remain.

158

5.2 LITERATURE CITED (1) Bargar, J. R.; Reitmeyer, R.; Davis, J. A. Spectroscopic confirmation of uranium(VI) - carbonato absorption complexes on hematite. Environ. Sci. Technol. 1999, 33, 2481-2484. (2) Ames, L. L.; McGarah, J. E.; Walker, B. E. Sorption of trace constituents from aqueous solution onto secondary minerals. 1. Uranium. Clays Clay Minerals 1983, 31, 145-164. (3) Casas, I.; Casabona, L.; Duro, L.; dePablo, J. The influence of hematite on the soprtion of uranium(VI) onto granite filling fractures. Chem. Geol. 1994, 113, 319-326. (4) Liger, E.; Charlet, L.; Cappellen, P. V. Surface catalysis of uranium(VI) reduction by iron(II). Geochim. Cosmochim. Acta 1999, 63, 2939-2955. (5) Stewart, B. D.; Neiss, J.; Fendorf, S. Quantifying constraints impoed by calcium and iron on bacterial reduction of uranium(VI). Journ. Env. Qual. 2006, In Press. (6) Wielinga, B.; Bostick, B.; Hansel, C. M.; Rosenzweig, R. F.; Fendorf, S. Inhibition of bacterially promoted uranium reduction: Ferric (hydr)oxides as competitive electron acceptors. Environmental Science & Technology 2000, 34, 2190-2195. (7) Brooks, S. C.; Fredrickson, J. K.; Carroll, S. L.; Kennedy, D. W.; Zachara, J. M.; Plymale, A. E.; Fendorf, S. Inhibition of bacterial U(VI) reduction by calcium. Environ. Sci. Technol. 2003, 37, 1850-1858. (8) Dodge, C. J.; Francis, A. J.; Gillow, J. B.; Halada, G. P.; Eng, C.; Clayton, C. R. Association of uranium with iron oxides typically formed on corroding steel surfaces. Environ. Sci. Technol. 2002, 36, 3504-3511. (9) Finneran, K. T.; Housewright, M., E.; Lovley, D. R. Multiple influences of nitrate on uranium solubility during bioremediation of uranium-contaminated subsurface sediments. Environ. Microbiol. 2002, 4, 510-516. (10) Pierce, E. M.; Icenhower, J. P.; Serne, J. N.; Catalano, J. G. Experimental determination of UO2(cr) dissolution kinetics: Effects of solution saturation state and pH. Jour. Nucl. Mat. 2005, 345, 206-218. (11) Torrero, M. E.; Baraj, E.; de Pablo, J.; Gimenez, J.; Casas, I. Kinetics of corrosion and dissolution of uranium dioxide as a function of pH. International Journal of Chemical Kinetics 1997, 29, 261-267. (12) Wersin, P.; Hochella, M. F.; Persson, P.; Redden, G.; Leckie, J. O.; Harris, D. W. Interaction between aqueous uranium(VI) and sulfide minerals : Spectroscopic evidence for sorption and reduction. Geochim. Cosmochim. Acta 1994, 58, 2829-2843. (13) Hua, B.; Xu, H.; Terry, J.; Deng, B. Kinetics of uranium(VI) reduction by hydrogen sulfide in anoxic aqueous systems. Environ. Sci. Technol. 2006, 40, 4666-4671. (14) Ilton, E.; Heald, S. M.; Smith, S. C.; Elbert, D. C.; Liu, C. Reduction of uranyl in the interlayer region of low iron micas under anoxic and aerobic conditions. Environ. Sci. Technol. 2006, 40, 5003-5009.

159

(15) Fredrickson, J. K.; Zachara, J. M.; Kennedy, D. W.; Liu, C. G.; Duff, M. C.; Hunter, D. B.; Dohnalkova, A. Influence of Mn oxides on the reduction of uranium(VI) by the metal-reducing bacterium Shewanella putrefaciens. Geochim. Cosmochim. Acta 2002, 66, 3247-3262. (16) Wu, W.-M.; Carley, J.; Fienen, M.; Mehlhorn, T.; Lowe, K.; Nyman, J.; Luo, J.; Gentile, M.; Rajan, R.; Wagner, D.; Hickey, R.; Gu, B.; Watson, D. B.; Cirpka, O.; Kitanidis, P.; Jardine, P. M.; Criddle, C. Pilot-scale in situ bioremediation of uranium in a highly contaminated aquifer. 1. Conditioning of a treatment zone. Environmental Science & Technology 2006, 40, 3978-3985. (17) Wu, W.-M.; Carley, J.; Gentry, T.; Ginder-Vogel, M.; Fienen, M.; Mehlhorn, T.; Yan, H.; Caroll, S.; Pace, M.; Nyman, J.; Luo, J.; Gentile, M.; Fields, M. W.; Hickey, R.; Watson, D. B.; Cirpka, O.; Zhou, J.; Fendorf, S.; Kitanidis, P.; Jardine, P. M.; Criddle, C. Pilot-scale in situ bioremediation of uranium in a highly contaminated aquifer. 2. Geochemical control of U(VI) bioavailability and evidence of U(VI) reduction. Environ. Sci. Technol. 2006, 40, 3986-3995. (18) Anderson, R. T.; Vrionis, H. A.; Ortiz-Bernard, I.; Resch, C. T.; Long, P. E.; Dayvault, R.; Karp, K.; Marutzky, S.; Metzler, D. R.; Peacock, A. D.; White, D. C.; Lowe, M.; Lovley, D. R. Stimulating the in situ activity of Geobacter species to remove uranium from the groundwater of a uranium-contaminated aquifer. Appl. Environ. Microb. 2003, 69, 5884-5891. (19) Duff, M. C.; Coughlin, J. U.; Hunter, D. B. Uranium co-precipitation with iron oxide minerals. Geochim. Cosmochim. Acta 2002, 66, 3533-3547.

160

Figure 5.1. Conceptual model of uranium biogeochemical cycling.

161

Figure 5.2. Representative redox couples of concern when predicting environmental

uraninite stability.

EH (pH = 7) [volt]

1.0

0.8

0.6

0.4

0.2

0

-0.2

-0.4

Oxidized Reduced Oxidized Reduced

Ca2UO2(CO3)3

UO2(CO3)22- UO2

UO2

UO22+ UO2

O2 H2O

NO3- N2

MnO2(s) MnCO3

NO3- NO2

-

Fe(OH)3 Fe2+

pε (pH = 7)

0

-5

5

10

15

[U(VI)] = 1 x 10-6 M [Fe(II)] = 1 x 10-5 M

defining abiotic and biotic pathways of metal...

Documents