Chapter 11 Biogeochemical Uranium Redox Transformations: Potential Oxidants of Uraninite

Developments in Earth & Environmental Sciences, 7 Mark O. Barnett and Douglas B. Kent (Editors) r 2008 Elsevier B.V. All rights reserved DOI 10.1016/S1571-9197(07)07011-5

Chapter 11

Biogeochemical Uranium Redox Transformations: Potential Oxidants of Uraninite Matthew Ginder-Vogel1,2 and Scott Fendorf1, 1

Department of Geological and Environmental Sciences, Stanford University, Stanford, CA 94305-2115, USA 2 Present Address: Department of Plant and Soil Sciences, University of Delaware, Newark, DE 19716, USA

ABSTRACT In aerobic environments, uranium is generally found in the hexavalent oxidation state, is quite soluble, and readily forms complexes with calcium and carbonate. 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 natural attenuation of uranium and is being explored as a potential remediation option for this hazardous element. 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. Given their environmental predominance, Fe(III) (hydr)oxides, which act as both U(VI) reductants and U(IV) oxidants, exert a prominent role in determining uranium chemical fate and associated mobility. Here, we examine several factors controlling the extent of uraninite oxidation by Fe(III) (hydr)oxides, including iron oxide and bicarbonate concentration, in addition to iron oxide type. Our analysis reveals that the extent of uraninite oxidation increases under conditions that increase the thermodynamic probability of the reaction; however, recrystallization of poorly crystalline Fe(III) (hydr)oxides to more crystalline forms may ultimately limit the uraninite oxidation reaction. Nevertheless, uraninite oxidation by Fe(III) (hydr)oxides may be a limiting factor on uraninite stability in the environment.

Corresponding author. Tel.: +1 650 723 5238; Fax: +1 650 725 2199;

E-mail: [email protected] (S. Fendorf).

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11.1. Introduction The release of radionuclides, toxic heavy metals, and organic co-contaminants into the environment during the nuclear age poses a unique long-term environmental problem (Riley et al., 1992; Ginder-Vogel et al., 2005). Among these contaminants, uranium is of particular concern because of its carcinogenicity, long half-life, widespread distribution, and mobility (Riley et al., 1992). Uranium contamination of ground and surface waters has been detected at numerous sites throughout the world, including agricultural evaporation ponds (Bradford et al., 1990), nuclear weapons manufacturing areas, and mine tailings sites (Riley et al., 1992). The mobility of uranium within surface and subsurface environments is determined, in part, by its geochemical speciation. In oxic environments, uranium is generally present in the hexavalent oxidation state as the uranyl [U(VI)O2+ 2 ] species and under most environmental conditions is quite soluble – the solubility of U(VI) is particularly enhanced by complexation with carbonate, a common groundwater ligand (Grenthe et al., 1992). However, U(VI) also forms several sparingly soluble complexes with phosphate (Langmuir, 1978; Sandino and Bruno, 1992) and readily forms inner-sphere complexes with many transition metal hydroxides – absorption being most extensive at neutral pH (Barnes and Cochran, 1993; Barnett et al., 2000; Moyes et al., 2000; Bostick et al., 2002; Davis et al., 2004; Curtis et al., 2006). Nevertheless, U(VI) tends to be relatively soluble and thus subject to migration within ground and surface water. Conversely, U(IV) is sparingly soluble, even in the presence of common groundwater ligands such as carbonate, and thus tends to be relatively immobile. Therefore, the oxidation state of uranium, which may be controlled by a wide variety of biogeochemical processes (Fig. 11.1), will play an important role in determining its mobility in surface and subsurface environments. In fact, in situ transformation of mobile U(VI) species into immobile U(IV) uraninite is being explored as a possible uranium remediation technique (Wu et al., 2006a,b, 2007).

11.2. Uranium Oxidation–Reduction Reactions A host of reduction pathways exist for U(VI) and include both abiotic and biotic reactions. Abiotic (chemical) reduction of U(VI) may proceed via several pathways, albeit typically under a limited set of conditions, within low temperature geochemical systems (Fig. 11.1). Sulfide minerals, for

Biogeochemical Uranium Redox Transformations

295

Fe(III) Fe(II) Fe(III) Fe(II)

U(VI) (Soluble)

H2, OCreduced DMRB CO2, H2O, OCoxidized

Mn(IV)

H2, OCreduced SRB CO2, H2O, OCoxidized

Mn(II) 2-

0

S , SO4 HS-

-

NO, N2O, NO2

-

HS N2

2-

SO4

≡Fe(II)

T. denitrificans and others

Fe(III)

-

NO3

O2

AH2DS

N2 H2O

U(IV) (Uraninite)

AQDS

Figure 11.1: Biological and Abiotic Processes that Affect the Redox State of Uranium. Ovals Represent Biologically Catalyzed Processes and Include Uranium Reduction by Dissimilatory Metal-Reducing Bacteria (DMRB) and Sulfate-Reducing Bacteria (SRB). Other Abbreviations Include Surface Bound Iron (
Fe), Reduced Organic Carbon (OCreduced), Oxidized Organic Carbon (OCoxidized), 9, 10-Anthraquinone-2, 6-Disulfonic Acid (AQDS), and Reduced AQDS (AH2DS). example, are commonly found in association with supergene uranium deposits, suggesting sulfide reduction of U(VI) (Langmuir and Chatman, 1980; Nash et al., 1981). Additionally, laboratory studies have demonstrated that partial U(VI) reduction by sulfide mineral surfaces occurs with the concomitant production of polysulfides (Wersin et al., 1994; Livens et al., 2004). Furthermore, complete reduction of dissolved U(VI) to microcrystalline uraninite by aqueous sulfide occurs at low pH and low bicarbonate concentrations, where the predominant species of U(VI) is UO2þ 2ðaqÞ – the most reactive uranyl species toward sulfide (Hua et al., 2006). Although Fe2þ ðaqÞ does not appear to be a facile reductant of U(VI), adsorbed Fe(II) and ferrousbearing minerals are kinetically viable reductants of U(VI). Uranium(VI) reduction by ferrous iron bound to the surface of microcrystalline hematite, goethite, smectite, and natural solids, has been observed at near-neutral pH (Liger et al., 1999; Jeon et al., 2005), as has reduction by hydroxysulfate

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green rust (O’Loughlin et al., 2003). At low pH (o5), U(VI) is reduced to U(IV) in the interlayer of ferrous iron-bearing micas (Ilton et al., 2004, 2006), and partial uranyl reduction by magnetite is also observed (El Aamrani et al., 2000; Missana et al., 2003; Scott et al., 2005). In contrast to abiotic U(VI) reduction, numerous common dissimilatory metal-reducing bacteria (DMRB) and sulfate-reducing bacteria (SRB) 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) (Gorby and Lovley, 1992; Fredrickson et al., 2000), a sparingly soluble phase. However, the formation of the tenary Ca2UO2(CO3)3(aq) species kinetically limits biological uranium reduction (Brooks et al., 2003; Stewart et al., 2007). Additionally, the presence of nitrate or Fe(III) (hydr)oxides as alternate electron acceptors and potential U(IV) oxidants impedes biological U(VI) reduction (Wielinga et al., 2000; Stewart et al., 2007). Although DMRB are capable of reducing solid Fe(III) (hydr)oxides, solid U(VI) is not available for biological reduction. The reduction of NaBoltwoodite (NaUO2SiO3OH  1.5H2O) by Shewanella oneidensis strain MR-1 requires the sequential coupling of U(VI) dissolution with microbial reduction (Liu et al., 2006). Sediment-bound U(VI) also appears to be unavailable for microbial reduction (Ortiz-Bernad et al., 2004). However, because U(VI) reduction occurs concurrently with Fe(III) and SO2 4 reduction, systems with active populations of DMRB and SRB should contain abundant Fe(II) and HS1, both of which can abiotically reduce U(VI) sorbed to synthetic and natural iron (hydr)oxides (Liger et al., 1999; Jeon et al., 2005). Bacterial uranium reduction may also occur via the reduction of electron shuttling compounds, such as quinones, which then reduce U(VI) to U(IV) (Fredrickson et al., 2000; Nevin and Lovley, 2000). In order to determine the viability of in situ biological uranium remediation and to discern the role of reductive processes in natural uranium cycling, it is critical to determine the stability of biologically precipitated uraninite, and in particular, it is important to identify environmentally relevant oxidative processes (Fig. 11.1). Potential UO2 oxidants include molecular oxygen, nitrate, nitrate reduction intermediates, Mn(IV) (hydr)oxides, and Fe(III) (hydr)oxides (Figs. 11.1 and 11.2). Additionally, UO2 oxidation may be catalyzed by biological activity. The oxidation of biologically precipitated uraninite by molecular oxygen remains poorly characterized; however, the oxidative dissolution of synthetic uraninite is relatively well characterized. In those studies, the uraninite grains examined were generally 100 mm in diameter (Torrero et al., 1997; Pierce et al., 2005), rather than 10 nm in diameter, as observed for biologically precipitated uraninite (Fredrickson et al., 2000; Singer et al., 2006). However, such

Biogeochemical Uranium Redox Transformations

EH (pH = 7)

Oxidized

Reduced

Oxidized

Reduced

297

pε (pH = 7)

[volt] 1.0 15 O2 NO3-

0.8

H2O N2 10

0.6

MnO2(s) NO3-

0.4

MnCO3 NO25

0.2 0 -0.2

UO22+

UO2

2-

UO2 UO2

UO2(CO3)2 Ca2UO2(CO3)3

0 Fe(OH)3

Fe

2+

-5

-0.4 [U(VI)] = 1 x 10-6 M

[Fe(II)] = 1 x 10-5 M

Figure 11.2: Representative Redox Couples for Dominant Constituents within Soils/Sediments and Their Comparison to U(IV)/U(VI). studies still provide insight into the oxidation mechanism of, and possible geochemical limits on, uraninite oxidation by molecular oxygen. Below pH 7, the oxidation rate of uraninite increases as pH decrease, and an increase in either the bicarbonate or O2(aq) concentration results in an increase in the oxidation rate (Torrero et al., 1997; Peper et al., 2004; Pierce et al., 2005). Above pH 7 and in the presence of low bicarbonate concentrations, the uraninite oxidation rate decreases due to the precipitation of U(VI)-phases on the uraninite surface (Torrero et al., 1997; Peper et al., 2004; Pierce et al., 2005). Nitrate, a common co-contaminant with uranium (Riley et al., 1992), not only impedes biological uranium reduction (Finneran et al., 2002; Senko et al., 2002; Istok et al., 2004), but also may oxidize U(IV). Uraninite oxidation by nitrate is a thermodynamically favored process under environmental conditions (Fig. 11.2); however, it is rate-limited (Senko et al., 2005a,b). The  biological transformation of NO 3 into NO2 , NO, and N2O increases the oxidation rate; however, the rate remains quite slow as compared to oxidation by Fe(III) and O2 (Senko et al., 2005a). Additionally, Thiobacillus denitrificans and Geobacter metallireducens are capable of catalyzing nitratedependent U(IV) oxidation, although it is currently not known if the bacteria obtain energy from this process (Finneran et al., 2002; Beller, 2005).

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Furthermore, the indirect oxidation of U(IV) by nitrite, through production of Fe(III) (hydr)oxides, also increase this reaction rate (Senko et al., 2005a,b). In addition to their role as intermediates in U(IV) oxidation by nitrite, Fe(III) (hydr)oxide minerals have been implicated in U(IV) oxidation under sulfate-reducing (Sani et al., 2004, 2005) and methanogenic conditions (Wan et al., 2005). The redox couples for U(IV)/U(VI) and Fe(III) (hydr)oxide/ Fe(II) occur at similar potentials under common groundwater conditions (Fig. 11.3); therefore, small changes in aqueous and solid-phase chemistry can result in UO2 oxidation by Fe(III) (hydr)oxide oscillating between thermodynamic viability and nonviability (Fig. 11.3) (Ginder-Vogel et al., 2006). However, acidic, ferric iron-containing solutions rapidly oxidize and dissolve U(IV) from uranium ore minerals (Harrison et al., 1966; Vuorinen et al., 1985). This reaction is further enhanced by the activity of acidophilic, iron-oxidizing bacteria, such as Thiobacillus ferrooxidans (Dispirito and Tuovinen, 1982a,b). Iron(III) (hydr)oxides will likely play an important role in controlling the long-term stability of biologically precipitated uranium. Not only are they ubiquitous in soils and sediments (Cornell and Schwertmann, 2003), they accelerate uraninite oxidation by nitrite (Senko et al., 2005b) and are generated in environments undergoing redox cycling. However, the geochemical conditions conducive to uraninite oxidation by Fe(III) (hydr)oxides remain poorly constrained. In particular, the transformation of ferrihydrite into more thermodynamically stable Fe(III) (hydr)oxide minerals may ultimately limit uraninite oxidation. Accordingly, here we examine the potential impact of Fe(III) (hydr)oxide mineralogy on biogenic uraninite oxidation.

11.3. Experimental 11.3.1. Materials All acids were trace-metal grade, and chemicals were ACS grade or better. Anaerobic solutions used in this study were prepared using distilled deionized (DDI) water that had been treated to remove dissolved O2, by boiling, while purging with N2 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 were equilibrated in the anaerobic chamber for 24 h prior to use.

Biogeochemical Uranium Redox Transformations

EH

Ox

(V)

[U(IV)] = 10-9 M [U(VI)] = 10-6 M

-0.10

UO2(CO3)2-2

Red

Ox

Red

299

pε

[Fe(II)] = 10-5 M

Fe(OH)3

UO2

Fe2+

-0.15

-2.5

+3

-0.20

UO2(CO3)2-2

U(OH)

-0.25

UO2(CO3)2-2

U(OH)4

UO2(CO3)2-2

U(CO3)4

γ-FeOOH

Fe2+

α-FeOOH

Fe2+

Fe2O3

Fe2+

-4

-0.30

-5.0

-0.35

-0.40

-0.45

-7.5 -6

UO2(CO3)2-2

U(CO3)5

UO2(CO3)2-2

U

-0.50

-0.55

-0.60

+4

-10.0

Figure 11.3: Representative Fe(III)/Fe(II) and U(VI)/U(IV) Redox Couples 6 M of Each U(VI) Species, at pH 7 with 3  103 M HCO 3 , 1  10 9 1  10 M of Each Dissolved U(IV) Species, and 1  105 M Fe(II).

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11.3.2. Biogenic Uraninite Cell suspensions were prepared by growing Shewanella putrefaciens strain CN32 aerobically on Tryptic Soy Broth at 301C to late log phase. Cells were harvested by centrifugation (4,000 g, 5 min), 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, 3 mM NH4Cl, 40 mM lactate, and 10 mL Wolfe’s vitamins) with 100 mL of bacterial suspension (108 cells mL1). Media was then stirred continuously in an anaerobic glove box with an atmosphere of 95% N2 and 5% H2 (Coy Laboratory Products). After 4 days (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, DDI water. X-ray diffraction (XRD) patterns were identical to those previously reported for biogenic UO2 (Fredrickson et al., 2000). The N2-BET surface area of the biogenic uraninite was 129 m2 g1.

11.3.3. Ferrihydrite Synthesis Two-line ferrihydrite was synthesized by rapidly titrating a ferric chloride solution with NaOH to a pH of 7.5 (Hansel et al., 2005). The Fe(III) (hydr)oxide flocs were washed by centrifugation twice with 1% HCl and three times with distilled H2O. The Fe(III) (hydr)oxide flocs were then resuspended and degassed by bubbling with N2 for 24 h. Oxide mineralogy and purity was confirmed with XRD. The N2-BET surface areas of the ferrihydrite, goethite, and hematite were 219, 80, and 60 m2 g1, respectively.

11.3.4. Uraninite 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 2 weeks, and ferrihydrite was used within 2 d of preparation. Changes in reactivity of the two solid phases were not observed over this time frame. A set of oxidation experiments was performed to examine the effect of ferrihydrite concentration on reaction extent and solid-phase Fe and U speciation during UO2 oxidation. These reactions were carried out in 1 L of

Biogeochemical Uranium Redox Transformations

301

media in 1 L polypropylene Nalgene bottles with a single uraninite concentration of 30.7 m2 L1 and ferrihydrite concentrations of 23.5, 46.0, and 80.2 m2 L1. All reactions were continuously mixed, 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. Ferrihydrite was added to the reaction media first, followed by uraninite, in less than 10 s for all reaction conditions studied. 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 L1. 11.3.5. Sampling and Analytical Procedures The extent of biogenic UO2 oxidation by ferrihydrite was quantified, primarily using ferrozine extractable Fe(II) concentrations (Stookey, 1970). The system-partitioning coefficient for Fe(II) was used to determine total Fe(II), which was in excellent agreement with the stoichiometric amounts of U(VI) produced. Use of dissolved U(VI) to quantify the extent of reaction was less reliable because of the potential for variable U(VI) sorption on Fe(III) (hydr)oxides under the varying reaction conditions examined. Ferrozineextractable Fe(II) was used, rather than soluble or acid-extractable Fe(II), because (i) of the propensity for Fe(II) uptake by Fe(III) (hydr)oxides (Morrison et al., 1995; Moyes et al., 2000; Duff et al., 2002; Walter et al., 2003), and (ii) UO2(biogenic) oxidation by Fe(III) is more energetically favorable (Ginder-Vogel et al., 2006) and rapid under acidic conditions. In fact, after 24 h, we observe near-complete oxidation of 30.7 m2 L1 biogenic uraninite by 32.1 m2 L1 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 mm 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. Total iron and uranium were determined at the conclusion of each oxidation experiment by acidic dissolution of the reaction slurry with concentrated HNO3 and 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 mm polycarbonate filter, which was then acidified and measured by ICP-OES. Soluble Fe(II) in the filtrate was measured by the ferrozine method, using 10 mM ferrozine, and soluble U(VI) was measured

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spectroflourometrically. 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. 11.3.6. X-Ray Diffraction and X-Ray Absorption Spectroscopy 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 (12732.137 eV) and a MAR 345 image plate. The resulting images were processed using FIT2D (Hammersley, 1997). The sampleto-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. Extended X-ray absorption fine 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 a wide-angle (Lytle) fluorescence chamber. Incident and transmitted X-ray intensities were measured with in-line ionization chambers. The energy range studied was 200 to +1,000 eV around the Fe K-edge (7,112 eV). All samples were internally referenced to a Fe-metal standard, placed between the second and third in-line ionization chambers. Two to four individual spectra were averaged for each sample. EXAFS spectra were processed using the SixPACK (Webb, 2005) interface to IFEFFIT (Newville, 2001). 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 (Newville, 2001). Linear combination fitting routines were used to reconstruct the

Biogeochemical Uranium Redox Transformations

303

experimental spectrum and to determine the relative percentages of iron mineral phases (Hansel et al., 2003; Hansel et al., 2005). Each spectrum was fit using 2-line ferrihydrite, lepidocrocite, and goethite, which were detected in the XRD patterns for the 23.5 m2 L1 ferrihydrite experiment (Fig. 11.4A).

11.3.7. Thermodynamic Reaction Modeling Except for the calcium–uranyl–carbonate species, Gibbs free energies of formation for all uranium species were obtained from Guillaumont et al. (2003). 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 (Giammar and Hering, 2001; Fredrickson et al., 2002; Suzuki et al., 2002). Amorphous uraninite (UO2(am)) is not thermodynamically well defined, but it provides a more realistic prediction of reactivity over shorter time-scales than crystalline UO2 (uraninite) (Wan et al., 2005). The Gibbs free energy of formation for Ca2UO2(CO3)3 and CaUO2(CO3)2 3 were calculated from stability constants provided in Dong and Brooks (2006). The value for hematite (Fe2O3) was obtained from Cornell and Schwertmann (2003), the values for lepidocrocite and goethite were obtained from Majzlan et al. (2003), and the value for 2-line ferrihydrite (Fe(OH)3) was obtained from Majzlan et al. (2004). All values were checked for internal consistency and are tabulated in Ginder-Vogel et al. (2006). The Gibbs free energy of reaction for specific conditions was calculated using standard convention at 298 K and noted in Table 11.1.

11.4. Results 11.4.1. Effect of Ferrihydrite Concentration Uraninite Oxidation Extent The ability of ferrihydrite to oxidize UO2 in the absence of calcium was investigated in a series of batch experiments in 3 mM HCO 3 at pH 7. After 48 h of reaction, the proportion of uraninite oxidized (calculated as the difference between the control and experimental dissolved U(VI) concentration) increases from 7.5% at 23.5 m2 L1 to 21.2% at 80 m2 L1 ferrihydrite (Fig. 11.5C).

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Intensity (Counts)

chi(k)*k3

A

0

10

20

30

40

50

60

chi*k3

chi*k3

2 theta (degrees)

Figure 11.4: Iron K-Edge EXAFS Spectra (Solid Lines) and Linear Combination Fits (Dotted Lines) Used to Construct Fig. 11.6A–C. XRD Patterns Used to Identify Fe Minerals for Linear Combination Fits (Left Side of A); g ¼ Goethite, l ¼ Lepidocrocite, and u ¼ Uraninite.

Table 11.1: Gibb’s Free Energy of Reactions at Standard State Conditions (DGr1) and at Experimental Conditions  Represented by pH 7, 1  106 M U(VI) Species, 5  107 M Fe2+, and 3  103 M HCO 3 (DGr ). Oxidation reaction

60.08 58.97 48.53 33.98 58.40 44.02 44.14 42.38 41.28 30.83 16.28 40.70 26.32 26.44 51.38 50.28 39.83 25.28 49.70 35.32 35.44 39.82 38.72 28.28 13.73 38.15 23.77 23.89

6.6 5.0 7.3 5.5 12.8 7.0 5.2 24.4 12.7 10.4 12.1 4.84 10.7 12.5 15.4 3.7 1.4 3.14 4.1 1.7 3.5 26.9 15.2 12.9 14.7 7.4 13.2 15.1

305

Thermodynamic data sources are described in the ‘‘Experimental’’ section. 1Standard state. Convention for noting the chemical gradients which are included in the thermodynamic value.

DGr (kJ mol1)

Biogeochemical Uranium Redox Transformations

Fe(OH)3+0.5UO2+3H+2Fe2++0.5UO2+ 2 +3H2O 2+ +3H2O Fe(OH)3+0.5UO2+2.5H++0.5HCO 3 20.5UO2CO3+Fe 2+ 2 +3H2O Fe(OH)3+0.5UO2+2H++HCO 3 20.5UO2(CO3)2 +Fe 2+ 4 +3H2O Fe(OH)3+0.5UO2+1.5H++1.5HCO 3 20.5UO2(CO3)3 +Fe  + 2+ Fe(OH)3+0.5UO2+1.5H +1.5HCO3 +Ca 20.5Ca2UO2(CO3)3+Fe2++3H2O 2+ 2+ 20.5CaUO2(CO3)2 +3H2O Fe(OH)3+0.5UO2+1.5H++1.5HCO 3 +0.5Ca 3 +Fe 2+  20.25(UO ) CO (OH) +Fe +2.25H Fe(OH)3+0.5UO2+2H++0.25HCO 3 2 2 3 3 2O 2+ +2H2O a-FeOOH+0.5UO2+3H+20.5UO2+ 2 +Fe 2+ +2H2O a-FeOOH+0.5UO2+2.5H++0.5HCO 3 20.5UO2CO3+Fe 2+ 2 20.5UO (CO ) +Fe +2H2O a-FeOOH+0.5UO2+2H++1HCO 3 2 3 2 2+ 4 +H2O a-FeOOH+0.5UO2+1.5H++1.5HCO 3 20.5UO2(CO3)3 +Fe 2+ 20.5Ca2UO2(CO3)3+Fe2++2H2O a-FeOOH+0.5UO2+1.5H++1.5HCO 3 +Ca 2+ 2+ 20.5CaUO2(CO3)2 +2H2O a-FeOOH+0.5UO2+1.5H++1.5HCO 3 +0.5Ca 3 +Fe 2+  +5/4H2O a-FeOOH+0.5UO2+2H++0.25HCO 3 20.25(UO2)2CO3(OH)3 +Fe 2+ +2H2O g-FeOOH+0.5UO2+3H+20.5UO2+ 2 +Fe +  g-FeOOH+0.5UO2+2.5H +0.5HCO3 20.5UO2CO3+Fe2++2H2O 2+ 2 +2H2O g-FeOOH+0.5UO2+2H++1HCO 3 20.5UO2(CO3)2 +Fe 2+ 4 20.5UO (CO ) +Fe +H2O g-FeOOH+0.5UO2+1.5H++1.5HCO 3 2 3 3 + 2+  g-FeOOH+0.5UO2+1.5H +1.5HCO3 +Ca 20.5Ca2UO2(CO3)3+Fe2++2H2O 2+ 2+ 20.5CaUO2(CO3)2 +2H2O g-FeOOH+0.5UO2+1.5H++1.5HCO 3 +0.5Ca 3 +Fe 2+  20.25(UO ) CO (OH) +Fe +5/4H g-FeOOH+0.5UO2+2H++0.25HCO 3 2 2 3 3 2O 0.5Fe2O3+0.5UO2+3H+2Fe2++0.5UO2+ 2 +1.5H2O 2+ +1.5H2O+0.5UO2CO3 0.5Fe2O3+0.5UO2+2.5H++0.5HCO 3 2Fe 2 2+ 2Fe +1.5H 0.5Fe2O3+0.5UO2+2H++1HCO 3 2O+0.5UO2(CO3)2 2+ +1.5H2O+0.5UO2(CO3)4 0.5Fe2O3+0.5UO2+1.5H++1.5HCO 3 2Fe 3 2+ 2Fe2++1.5H2O+0.5Ca2UO2(CO3)3 0.5Fe2O3+0.5UO2+1.5H++1.5HCO 3 +Ca 2+ 2Fe2++1.5H2O+0.5CaUO2(CO3)2 0.5Fe2O3+0.5UO2+1.5H++1.5HCO 3 +0.5Ca 3 2+  +0.75H2O 0.5Fe2O3+0.5UO2+2H++0.25HCO 3 +0.25(UO2)2CO3(OH)3 +Fe

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A

B

C

Figure 11.5: Percentage of Initial Uraninite Oxidized after 48 h as a Function of Fe(III) (Hydr)Oxide Type (A), Bicarbonate Concentration (B), and Ferrihydrite Concentration (C). Data Points for A and B are Calculated from Ginder-Vogel et al. (2006). Not Detected (ND).

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11.4.2. Fe(III) (Hydr)Oxide Mineral Evolution In the presence of small amounts of Fe(II) (o0.67 mmol Fe(II) g1 ferrihdyrite), ferrihydrite transforms into lepidocrocite and goethite (Hansel et al., 2005). During 48 h of oxidation, dissolved Fe(II) remains below detection limits; extractable Fe(II) at all three ferrihydrite suspension concentrations reaches a maximum of 0.2 mmol Fe(II) g1 ferrihydrite (Fig. 11.6D), and dissolved U(VI) also plateaus at 75, 112, and 211 mM for ferrihydrite concentrations of 23.5, 46.0, and 80.2 m2 L1 ferrihydrite (data not shown). At all ferrihydrite suspension concentrations, the solid phase remains predominantly ferrihydrite over the first 2 h of the experiment (Figs. 11.4 and 11.6). Crystalline Fe(III) (hydr)oxides are not detected in XRD diffraction patterns prior to 5 h of reaction time (Fig. 11.4). As the reaction progresses, the goethite concentration continues to increase with detectable amounts of lepidocrocite accumulating only after 6 h 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 (Fig. 11.6A). The transformation of ferrihydrite is notably more complete at increasing ferrihydrite concentration, with only 8% of the ferrihydrite remaining in the 46 m2 L1 reaction and no ferrihydrite remaining in the 80.2 m2 L1 reaction after 48 h (Fig. 11.6B and C). Uraninite oxidation was confirmed by comparing the solid phase-uranium oxidation state of the last sample from each reaction to the oxidation state of biogenic uraninite. Uranium(VI) content of the biogenic uraninite is 5%, and 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 L1 ferrihydrite (data not shown).

11.5. Discussion 11.5.1. Extent of Uraninite Oxidation The prevalence of materials capable of oxidizing biologically precipitated uraninite and geochemical characteristics will influence the long-term stability of biologically precipitated uraninite. Although microcrystalline uraninite has an estimated solubility of 108 to 1010 at pH 7 (Casas et al., 1998), it can be oxidized and remobilized by several common environmental constituents (Fig. 11.1). Given Fe(III) (hydr)oxide’s environmental

Figure 11.6: Transformation of Ferrihydrite by Fe(II) Generated during Uraninite Oxidation. Reactions were Conducted with 3.0 mM KHCO3, 30.7 m2 L1 UO2, and Either (A) 23.5, (B) 46.0, or (C) 80.2 m2 L1 Ferrihydrite in 3 mM HCO 3 at pH 7.2. Percentages (75%) were Determined from Linear Combination Fits of k3-Weighted Fe EXAFS Spectra (k ¼ 1–14) (Fig. 11.4). (D) Ferrozine Extractable Fe(II) for Each Reaction. 1  103 mmol Fe(II) m2 Ferrihydrite is Equivalent to 0.2 mM Fe(II) g1 Ferrihydrite.

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prevalence and tendency to form in redox-active environments (Cornell and Schwertmann, 2003), it is crucial to develop a thorough understanding of the conditions that enhance uraninite oxidation by Fe(III) (hydr)oxides. The energetic favorability of uraninite oxidation by Fe(III) (hydr)oxides depends on Fe(III) (hydr)oxide type and abundance, in addition to geochemical conditions. For instance, we (Ginder-Vogel et al., 2006) observed that in 48 h 1.8% of the initial biogenic uraninite is oxidized by 2-line ferrihydrite in 3 mM HCO 3 at pH 7 (Fig. 11.5B); in contrast, uraninite exposure to goethite and hematite results in less, or no, uraninite oxidation (Fig. 11.5A). A similar increase in the amount of uraninite oxidized also occurs with increasing HCO 3 concentration, with 1.8% oxidized at 3 mM  HCO 3 and nearly 10% oxidized at 100 mM HCO3 . However, increasing the ferrihydrite concentration has the most dramatic effect on the extent of uraninite oxidation, with 7.4% oxidized at 23.5 m2 L1 and >20% oxidized at 80.2 m2 L1 ferrihydrite. As opposed to changes in the Fe(III) (hydr)oxide type and HCO 3, changes in ferrihydrite concentration do not affect the thermodynamic favorability of the uraninite oxidation reaction. This suggests that the higher concentration of ferrihydrite is acting as a sink for Fe(II) and/or U(VI) produced during uraninite oxidation, which may otherwise inhibit uraninite oxidation. At 3 mM bicarbonate and 10 mM U(VI)(aq), the majority of uranium should be in solution and thus only slightly affected by changes in ferrihydrite concentration. However, uraninite oxidation appears to cease as extractable Fe(II) concentrations approach 1  103 mmol m2 ferrihdyrite, despite the accumulation of o1 mM dissolved Fe(II). Assuming that ferrihydrite has two surface sites per nanometer square (Davis and Kent, 1990), equivalent to 3.7  103 mmol surface site meter square, the 1.0  103 mmol m2 (0.2 mmol g1) extractable Fe(II) present after 10 h reaction time at all ferrihydrite concentrations (Fig. 11.6D) accounts for 25% of the total initial surface sites available and may prevent U(IV) from reaching reactive surface sites. Additionally, Fe(II) sorption to ferrihydrite may result in electron delocalization within the Fe (hydr)oxide structure, thereby lowering its redox potential (Williams and Scherer, 2004; Iordanova et al., 2005; Kerisit and Rosso, 2005; Larese-Casanova and Scherer, 2007). This suggests that 0.2 mmol g1 ferrihydrite is a thermodynamic threshold for the energetic favorability of uraninite oxidation. The alteration of ferrihydrite into more crystalline, lower surface area Fe(III) (hydr)oxide phases will further reduce the number of surface sites available for uraninite oxidation.

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11.5.2. Evolution of Fe(III) (Hydr)Oxide Mineralogy In addition to limiting uraninite oxidation, Fe(II) generated during reaction will catalyze the transformation of ferrihydrite into an assemblage of more thermodynamically stable iron (hydr)oxide minerals. This mineral assemblage, depending on initial chemistry and Fe(II) concentration, is often dominated by lepidocrocite, goethite, or magnetite (Fredrickson et al., 1998; Benner et al., 2002; Hansel et al., 2003), all of which are less favorable oxidants than ferrihydrite (Cornell and Schwertmann, 2003). Indeed, as uraninite oxidation proceeds, Fe(II) induces the transformation of ferrihydrite into lepidocrocite and goethite (Fig. 11.6). Despite the similarity in the amount of Fe(II) generated relative to initial surface area, ferrihydrite conversion into lepidocrocite and goethite is noticeably more complete at 46.0 m2 L1 ferrihydrite (Fig. 11.6B) and proceeds to completion at 80.2 m2 L1 ferrihydrite (Fig. 11.6C). Although it is not possible to determine the exact cause of the observed differences in Fe(III) (hydr)oxide mineralogy from this study, the higher U(VI) concentrations produced in conjunction with the higher ferrihydrite concentration may consume dissolved bicarbonate through complexation reactions. The decrease in dissolved bicarbonate concentration may alter the reaction products produced during ferrihydrite recrystallization (Schwertmann and Cornell, 2000; Hansel et al., 2005); additionally, the increase in U(VI) concentration may, itself, alter the products of ferrihydrite transformation (Duff et al., 2002). Based upon thermodynamic considerations and previous studies, lepidocrocite accumulation should be observed prior to goethite accumulation (Hansel et al., 2005); however, goethite accumulates prior to lepidocrocite at all three ferrihydrite concentrations (Fig. 11.6). Since the oxidation of biogenic uraninite by lepidocrocite remains thermodynamically favorable at the low Fe(II)(aq) and U(VI)(aq) concentrations observed during the first 5 h of reaction, it may be a transient species that is rapidly consumed by uraninite oxidation. Indeed, uraninite oxidation by ferrihydrite, lepidocrocite, and goethite is no longer thermodynamically favorable at the reaction conditions observed after 10 h of reaction (100 mM U(VI)(aq), pH 7.2, 3 mM HCO 3 , and 1 mM Fe(II)(aq)).

11.6. Implications for Biogeochemical Uranium Cycling During uranium reduction in biostimulated environments, Fe(II) concentrations are frequently 10–40 mM (Wu et al., 2007), which will likely limit

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uraninite oxidation by Fe(III) (hydr)oxides. However, intrusion of nitrate or molecular oxygen into a previously reduced (anaerobic) environment may consume Fe(II), resulting in the formation of poorly crystalline Fe(III) (hydr)oxides capable of oxidizing biologically precipitated uraninite. Consideration of this oxidation pathway is particularly important in the case of nitrate, a common co-contaminant with uranium, because abiotic and biotic oxidation of uraninite coupled to nitrate reduction occurs more slowly than oxidation by Fe(III) (hydr)oxides (Beller, 2005; Senko et al., 2005a). Additionally, the extent of biogenic uraninite oxidation by molecular oxygen increases below pH 7 in the presence of >20 mM Fe(II) (Zhong et al., 2005). Geochemical conditions play a critical role in determining the energetic favorability of the uraninite oxidation reaction with the favorability increasing with increasing carbonate and calcium concentrations and decreasing pH (Ginder-Vogel et al., 2006). However, the ratio of ferrihydrite to biogenic uraninite appears to be the most important factor affecting the extent of uraninite oxidation at neutral pH, with the amount of uraninite oxidized scaling linearly with ferrihydrite concentration (Fig. 11.5C); however, a 33-fold increase in the bicarbonate concentration only results in a 5-fold increase in the amount of uraninite oxidized (Fig. 11.5B). In batch systems, Fe(II) produced during the oxidation reaction inhibits uraninite oxidation at concentrations of 1.0  103 mmol m2 ferrihydrite; however, it is likely that transport of reaction products, including Fe(II) and U(VI), in groundwater systems would enhance the extent of uraninite oxidation. Additionally, Fe(II) catalyzes the transformation of ferrihydrite into thermodynamically more stable Fe(III) (hydr)oxide phases, which affects the viability and extent of uraninite oxidation (Figs. 11.3 and 11.5). The importance of considering Fe(III) (hydr)oxide crystallinity on the reoxidation of biogenic UO2 by Fe(III) (hydr)oxides is exemplified by several field and soil-column experiments (Fig. 11.7). Surprisingly, for all conditions reported, oxidation of biogenic UO2 by ferrihydrite is thermodynamically probable at Fe(II) concentrations of 20 mM. However, in the presence of 20 mM Fe(II), uraninite oxidation by lepidocrocite is thermodynamically favored for two conditions, and oxidation by goethite is only thermodynamically probable for one condition (Fig. 11.7). A host of U(IV) oxidants exist in surface and subsurface environments that include poorly crystalline Fe(III) (hydr)oxide phases. Therefore, while, iron(II) could serve to consume molecular oxygen in oxic waters, resulting ferric (hydr)oxides may continue to serve as oxidants of U(IV). However, long-term redox cycling promotes the conversion of poorly crystalline Fe(III) (hydr)oxide minerals into more crystalline forms, e.g., goethite or hematite (Thompson et al., 2006) and may be one method of limiting uraninite

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0

UO2 Oxidation Favored

-1

log (HCO3 ) (M)

-2

e-

Senko

II)

Istok M 0µ

( Fe

Sani Wan-2

2

Anderson

it eth

II)

Go

µM

-3

( Fe

Wu-3

Wu-2 20 et i c Wu-1 o r

oc

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-4

Wan-1

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Le

0

ite dr

-5

-2

µM

( Fe

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r Fe

2-

UO2(CO3)2 Reduction Favored

-6 5.5

6.0

6.5

7.0 pH

7.5

8.0

8.5

Figure 11.7: Thermodynamic Viability of UO2 (biogenic) Oxidation by Fe(III) (Hydr)Oxides for Conditions Reported for Various Field Sites and SoilColumn Experiments (Anderson et al., 2003; Istok et al., 2004; Sani et al., 2005; Senko et al., 2005a; Wan et al., 2005; Wu et al., 2006b). Conditions Representing Equilibrium with 0.126 mM UO2(CO3)2 2 (the Drinking Water Maximum) and 20 mM Fe(II) are Illustrated for Each Iron Oxide. Reoxidation is Viable at Bicarbonate and pH Conditions Plotting Above the Line and Nonviable for those Plotting Below. Reactions for Each Line are Listed in Table 11.1. Relevant Geochemical Conditions and References are Detailed in Ginder-Vogel et al. (2006). oxidation by Fe(III) (hydr)oxides. Nonetheless, oxidation by ferric (hydr)oxides is thermodynamically favorable and may limit uranium sequestration under mildly reducing conditions.

ACKNOWLEDGMENTS We would like to thank two anonymous reviewers and the editors for providing input that greatly improved the manuscript. We thank Kathleen Ginder-Vogel for her constructive input on the manuscript. We also thank Sam Webb, John Bargar, and Joe Rogers for their assistance in the collection of X-ray data at SSRL. This work was funded by the Office of Science

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Biological and Environmental Research ERSD Program, U.S. Department of Energy (grant number ER63609-1021814) and the Stanford NSF Environmental Molecular Sciences Institute (grant number NSF-CHE-0431425). 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.

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