Uptake of Uranium(VI) by Pyrite under Boom Clay Conditions

May 3, 2010 - Abiotic Reductive Immobilization of U(VI) by Biogenic Mackinawite. Harish Veeramani , Andreas C. Scheinost , Niven Monsegue , Nikolla P...
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Environ. Sci. Technol. 2010, 44, 4210–4216

Uptake of Uranium(VI) by Pyrite under Boom Clay Conditions: Influence of Dissolved Organic Carbon C. BRUGGEMAN* AND N. MAES Waste & Disposal Expert Group, Unit R&D Disposal, Belgian Nuclear Research Centre (SCK•CEN), Boeretang 200, Mol, Belgium

Received March 23, 2010. Accepted April 1, 2010.

The uptake of uranium(VI) by natural pyrite, FeS2, was studied under conditions relevant for geological disposal of radioactive waste (anoxic atmosphere, ∼0.014 mol · L-1 NaHCO3 electrolyte) with special emphasis on the role of dissolved organic matter. Solution analysis of batch experiments with different initial concentrations of uranium(VI) (10-8-10-4 mol · L-1) was combined with X-ray absorption spectroscopy on the solid phase to elucidate the speciation of uranium in these systems and to gain insight into the major reaction mechanisms between uranium and pyrite. The results showed that, under the conditions of the experiments, uranium(VI) was at least partly reduced to a UO2(s)-like precipitate, although the predominant valence state of uranium in solution was likely uranium(VI). All observations indicate that the uranium solid-liquid distribution is governed by both reduction and adsorption processes. No significant amounts of uranium colloids (either intrinsic UO2 colloids or complexes with natural organic matter) were found in any of the samples. The presence of dissolved organic matter did, however, increase the final uranium solution concentration and decrease the fraction of uranium(IV) found in the solid phase.

Introduction Geological disposal and storage of high-level radioactive waste and spent fuel can result, under certain circumstances, in the potential release of considerable quantities of uranium in the subsurface. In Belgium, the Boom Clay formation is studied as a reference host formation for phenomenological studies of the installation of such a disposal site. The safe disposal of nuclear waste depends, among other factors, on the capacity of the host formation to limit the migration of the radionuclides once released. To assess the extent of subsurface contamination by uranium and the resulting radioactive dose that could potentially reach the biosphere, the speciation of this element needs to be predicted correctly. In general, UVI species are thought to be more soluble and, hence, more mobile than UIV species. Thus, the transformation of uranium from an oxidized to a more reduced form is seen as a potential mechanism for its immobilization/ stabilization in situ. The major redox-controlling minerals in Boom Clay are considered to be pyrite, FeS2 (1-5 wt %), and, to a lesser extent, siderite, FeCO3 (0-1 wt %) (1). The current-state conceptual model for describing redox processes in Boom Clay suggests that the measured solution * Corresponding author e-mail: [email protected]; phone: +32 14 33 32 33. 4210

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redox potential (Eh) is probably controlled by the equilibrium of pyrite and siderite (1) FeS2 + 9H2O + CO2 S 18H+ + 2SO42- + FeCO3 + 14e(1) Under Boom Clay conditions (pCO2 of 10-2.4 atm and sulfate concentration of about 10-6 mol · L-1), the following Eh-pH relationship can be derived Eh(V) ) -0.076pH + 0.34

(2)

Pyrite has already been identified as a medium for the uptake and possible reduction of UVI compounds at pH 3-6 in anoxic conditions by several authors, using both solution analysis and spectroscopic methods (2-4). The reaction products were identified as UVI compounds (their fraction depending on the starting concentration); a mixed UO2+x(s) precipitate [poorly crystalline U3O8(s) is suggested]; and FeS2 oxidation products such as polysulfides, dissolved sulfate, and precipitated iron(III) (oxy)(hydr)oxides. Moreover, all studies mention the coexistence of multiple pathways, either parallel or sequential, in the reaction of UVI with FeS2. Mostly, the uptake behavior was interpreted as resulting from an initial surface sorption of UVI, followed by a subsequent partial reduction to UIV and precipitation of UO2+x(s) (5, 6). However, none of the aforementioned studies examined the pathways under conditions representative for natural carbonate-containing clay sediments, that is, at a moderately alkaline pH, in the presence of dissolved inorganic carbon and dissolved organic carbon. Also, the UVI concentrations that were used were mostly quite high, on the order of 10-4-10-3 mol · L-1 (3, 4). In a very recent study, the adsorption edge of UVI (4 × 10-8 mol · L-1) on pyrite was found to reach a maximum at pH 5.5 and seemed to slightly decrease at pH > 7 (6).The uptake of UVI by pyrite under Boom Clay conditions [pH 8.2, 0.014 mol · L-1 NaHCO3, 0.4% CO2(g), with and without the presence of organic carbon] was already part of a previous study (7, 8). A partial reduction of UVI to UIV [resulting in the probable formation of U3O8(s)] was observed using micro-X-ray absorption near-edge structure (µXANES) spectroscopy. The system also showed slow kinetics, with stable solution U concentrations ([U]aq) being reached after 15 days (10-5 mol · L-1 UVI) and 46 days (10-3 mol · L-1 UVI). A significant fraction (25-30%) of the U in solution was found to consist of colloidal species (evidenced from ultrafiltration measurements), especially when dissolved organic carbon was present. Abiotic UVI reduction by surface-catalyzed processes in the presence of FeII-bearing phases has been the focus of many recent studies (9-13). Most of these studies observed at least partial reduction of UVI, although the proposed reaction products varied from UO2(s) to UO2+x(s) to UV. Dissolved inorganic carbon is normally assumed to inhibit the reduction of UVI to UIV because of its ability to form strong negatively charged complexes with UVI (14). These complexes have been shown to interfere with FeII-catalyzed surface reduction processes, although surface uptake and reduction of UVI is not completely inhibited (10). Because dissolved organic carbon reduces UVI sorption under moderately alkaline conditions (15), a similar effect is expected. Apart from abiotic redox processes, microbial reduction of UVI to biogenic UO2(s) by dissimilatory metal-reducing bacteria or sulfate-reducing bacteria (16-18), although frequently encountered in soils, is probably of minor interest under disposal conditions because of the restricted pore size of the 10.1021/es100919p

 2010 American Chemical Society

Published on Web 05/03/2010

Boom Clay sedimentary formation, which effectively inhibits microbial activity (19). In the present study, UVI was added in a range of concentrations (from 10-8 to 10-4 mol · L-1) to FeS2 suspensions under anoxic conditions in the presence of dissolved inorganic carbon (0.014 mol · L-1 NaHCO3), both in the presence and in the absence of dissolved organic carbon (∼70 mg · L-1). These batch systems were equilibrated for 120 days, which, according to the previous study (7), should be more than sufficient to reach equilibrium. Thereafter, solution analysis of dissolved and colloidal U was performed, and both XANES and extended X-ray absorption fine structure (EXAFS) spectroscopy were used to analyze the U solid-phase speciation. To interpret XANES and EXAFS data, several U standards and similar batch systems equilibrated for only 10 days were analyzed as well. Our results show that UVI is at least partially reduced to a UO2(s)-like structure. Moreover, the UVI/UIV ratio appears to be largely influenced by dissolved organic carbon (DOC). These observations highlight the need to further elucidate the underlying mechanisms of heterogeneous redox processes in near-natural systems.

Experimental Section Batch Isotherm Experiments with Pyrite, FeS2: Solution Analysis. Two series of pyrite-containing batches were prepared, one in synthetic Boom Clay water (SBCW, ∼0.014 mol · L-1 NaHCO3, absence of dissolved organic carbon; for complete description and preparation procedure, see ref 1) and the other in real Boom Clay water (RBCW, piezometer water tapped from the HADES Underground Research Laboratory at Mol, Belgium, contains ∼70 mg · L-1 dissolved organic carbon; for a complete description, see ref 1), both at initial pH 8.5 ((0.2). The batch experiments were conducted under anoxic conditions in a glovebox with a controlled Ar/5% H2/0.4% CO2 atmosphere (O2 < 2 ppm). The FeS2 that was used originated from a cubic single crystal, which was ground and purified according to the procedure described elsewhere (20). The BET surface area of the resulting pyrite powder determined by N2 adsorption equaled 0.88 m2 · g-1. In a subsequent step, the FeS2 was preconditioned with the SBCW and RBCW electrolytes. UVI was added in a range of initial concentrations ([U]init) (from 10-8 to 10-4 mol · L-1), starting from a 10-1 mol · L-1 UO2(NO3)2 · 6H2O solution that was diluted in SBCW or RBCW, depending on the isotherm series. The final FeS2 solid/solution ratio in the U-containing samples was 5.1 ( 0.2 g · L-1. The systems were shaken gently and allowed to equilibrate for 120 days before phase separation by centrifugation, followed by 0.45-µm filtration. To identify any colloidal species, uranium solution concentrations ([U]aq) were measured by inductively coupled plasma mass spectrometry (ICP-MS) after 0.45-µm filtration [poly(vinylidene fluoride) (PVDF) membranes, VWR International) and 300- and 30-kDa centrifugal ultrafiltration [polyethersulfone (PES) membranes, Pall Corporation]. The dissolved organic carbon (DOC, Hach IL550 TOC-TN) concentration in all solutions was obtained spectrophotometrically as the UV absorbance at 280 nm (UV280, PerkinElmer Lambda40) using the following experimentally determined relationship: DOC (in mg · L-1) ) 33.028 × UV280 + 2.747. Fe concentrations after 0.45-µm filtration ([Fe]aq) were analyzed using ICP-AES. Dissolved SO42- ([SO42-]aq) and S2O32- ([S2O32-]aq) concentrations were obtained by anionexchange chromatography. Batch Experiments: X-ray Absorption Spectroscopy (XAS). A total of six batch samples containing FeS2 and SBCW or RBCW electrolyte were prepared for X-ray absorption spectroscopy (XAS) analysis according to the procedure described before, with an initial U concentration ([U]init) of 2 × 10-4 mol · L-1. Two samples were equilibrated for 120

days. These samples correspond to the isotherm series described above, but with a higher pyrite solid/solution ratio (20.0 ( 1.0 g · L-1) in order to have sufficient solid material and uranium for XAS analysis. The remaining four samples were equilibrated for only 10 days (FeS2 solid/solution ratios of 20.0 ( 1.0 and 4.1 ( 0.1 g · L-1). These samples were prepared to create a series of samples having similar matrixes and geochemical conditions, but with different fractions of U end members. It was speculated that U would first be sorbed as UVI and would slowly reduce to UIV. Before analysis, the samples were centrifuged, the supernatants were pipetted off, and the remaining solids were washed with distilled H2O (to dilute and remove any remaining U solution species; U solution analysis of these washing solutions revealed that on average less than 1% of solid-phase U was removed by this step). The solids were dried under vacuum, homogenized, and transferred to 500-µL cuvettes for XAS analysis. Data were collected for both XANES and EXAFS analyses. In addition to the six experimental samples, a number of U standards {0.1 mol · L-1 solutions of UO2(CO3)22-, UO2(CO3)34-, UO2(NO3)2, and UO2(OAc); solid phases of U3O8(s), poorly crystallized or amorphous UO2(s), uraninite UO2(c), and UVI adsorbed onto goethite [R-FeOOH(s)] in the absence and presence of bicarbonate} were also analyzed. Each sample or standard was measured at least three times, and the spectra were averaged and dead-time-corrected using SixPACK (version 0.53) software (21) running under the IFEFFIT program suite (22). XANES and EXAFS spectra were treated by the Athena and Artemis programs (23) to subtract the background from the spectra, normalize the spectra, and extract EXAFS spectra. The program FEFF8 (24, 25) was used to construct theoretical models on the basis of the crystallographic atomic positions. More details can be found in the Supporting Information.

Results and Discussion Batch Isotherm Experiments with Pyrite, FeS2: Solution Analysis. The solution-phase U concentrations ([U]aq) of the FeS2-containing batch isotherm series after 120 days of equilibration time are depicted in Figure 1. Figure 1a shows the SBCW series (absence of organics), and Figure 1b shows the RBCW series (∼70 mg · L-1 dissolved organic carbon). Tables SI1 and SI2 containing all experimental data can be found in the Supporting Information, section 1. The circled data points represent the results from U solution analysis of the samples used for XAS measurements (samples S1 and S2, 120-day equilibration time). The final pH equaled 8.4 ( 0.1 for SBCW systems and 8.9 ( 0.3 for RBCW systems. The final Eh values in the SBCW systems ranged from -20 to -180 mV vs SHE, and those in the RBCW systems ranged from -170 to -290 mV vs SHE. In the systems containing SBCW (Figure 1a), two samples with initial U concentrations ([U]init) of 10-8 and 10-7 mol · L-1 had final solution-phase U concentrations ([U]aq) below the analytical limit of detection (LOD, 0.1 ppb or 4 × 10-10 mol · L-1). The two samples with initial U concentrations of 10-6 and 10-5 mol · L-1 have final solution-phase U concentrations of about (4-8) × 10-9 mol · L-1. Ultrafiltration does not influence these concentrations, which was observed before (7). The sample with the highest initial U concentration (10-4 mol · L-1) showed a much higher final solution-phase U concentration (9 × 10-7 mol · L-1), which decreased by nearly 1 order of magnitude upon ultrafiltration at a molecular-weight cutoff (MWCO) of 30000. XAS sample S1 showed comparable behavior. In the RBCW samples (Figure 1b), natural U present in Boom Clay piezometer water imposes a background concentration of ∼0.5 ppb (or 2 × 10-9 mol · L-1, measured in a “blank” FeS2 sample containing RBCW electrolyte). The sample containing the lowest initial U concentration of 10-8 VOL. 44, NO. 11, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 1. Plots of solution-phase U concentration ([U]aq, mol · L-1) versus solid-phase U concentration ([U]solid,, mol · kg-1) in batch systems containing ∼5 g · L-1 FeS2 and ∼10-2 mol · L-1 NaHCO3: (a) SBCW and (b) RBCW series. The circled data points represent the results from samples used for XAS analysis (samples S1 and S2, 120-day equilibration time, 20 g · L-1 FeS2). mol · L-1 had a solution-phase U concentration around 0.5 ppb. The highest measured solution-phase U concentration in the present setup equaled 5.8 × 10-6 mol · L-1, which is 1 order of magnitude higher than measured before (7). Comparison of the SBCW and RBCW systems indicates that the presence of dissolved organic carbon apparently decreases the uptake of U by FeS2, in line with previous observations (15). No consistent influence of ultrafiltration was observed in the RBCW systems. Only for the systems with the lowest initial U concentration and for XAS sample S2 did a decrease in DOC results in a substantial decrease of U, presumably due to interactions between UIV and the humic substances. Indeed, complexation of UVI with DOC could occur (26, 27), but as shown before (7), complexation with Boom Clay humic substances is negligible under these conditions and could account for only 99% accuracy and that the experimental samples could be subdivided into two groups, with the presence of DOC being the most influential parameter. Based on these results, the maximum number of standards used to reconstruct the six samples during LCF, was constrained to 3. Using the set of standard spectra, it was possible to accurately reconstruct the XANES spectra of the experimental samples (Table 1 and Figure 3). In all cases, best fits were obtained with only two standards: the spectrum of UO2(c) and the spectrum of the UVI-goethite surface complex. In Figure 2, it is already apparent that the six U-FeS2 samples and these two U standards cross at an isosbestic point at ∼17180 eV. Recently, FeIII-O environment was detected using XPS on UVI-reacted pyrite (6), which confirms the possibility of a UVI-goethite-like surface complex in our systems. Target transformation (i.e., reconstruction) of these two spectra using the main eigenvectors from the PCA also appeared to be successful (see Supporting Information). The UVI/UIV ratio increased from 0 in sample S1 to more than 3 in sample S5 VOL. 44, NO. 11, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 3. XANES spectra (open symbols) and LCF reconstructions (blue lines) of samples S1-S6; standards that gave best fits were UO2(c) and a UVI -goethite surface complex (red symbols). (Table 1). Therefore, all of the FeS2(s) samples showed evidence of UVI reduction, and in the sample with longest equilibration time and in the absence of DOC, this reduction was near 100%. Shorter equilibration times and the presence of DOC decreases the amount of UVI reduced. Information on the local environment of the uranium atoms in the U-FeS2 samples is provided by analysis of the EXAFS data (except for sample S6 whose EXAFS spectrum exhibited too much noise to be useful for analysis), shown in Figure 4. The EXAFS data of all five samples can again be reconstructed using only two significant components extracted by PCA (see Supporting Information). Similar to the XANES spectra, the EXAFS spectra could be grouped depending on the presence of DOC in the respective samples (Table SI7, Supporting Information). This is also apparent in the Fourier transformed spectra (Figure 4b): the first shell in samples S2 and S4 (presence of DOC) appears to be split, but not in the spectra for S1, S3 and S5 (absence of DOC). Also, the height of the peak at ∼3.5 Å (indicated by a vertical arrow in Figure

4b) is consistently higher in samples S1, S3, and S5, compared to samples S2 and S4. The EXAFS spectra of the experimental samples were compared to several different theoretical models. The models involved contributions from different coordinating atoms (O, C, U, and Fe) at a variety of distances and in a variety of combinations, including the crystalline structures of uraninite, schoepite, uranylhydroxide, U3O8, U4O9, and so on. All details concerning the modeling of EXAFS spectra are given in the Supporting Information. Sample S1 was fit first in accordance with XANES results, showing that this sample’s spectrum closely resembles that of UO2(c). A best fit was obtained with a “nanoparticulate” UO2 model (11, 13, 36, 44). Although this fit was qualitatively and quantitatively not bad, the UO2 model alone could not account for all of the features present in the spectra. Also, the remainder of the U-FeS2 samples could not be accurately reconstructed using only one crystal structure. Because none of the structures individually provided a satisfactory fit to the data of the five experimental samples, the EXAFS spectra were then modeled with mixtures of two uranium species, in accordance with the results of XANES analysis. As a fitting strategy, we used the same number of neighboring atoms, interatomic distances, and Debye-Waller factors as obtained in the fit of UO2 and UVI-goethite surface complex end members. The amplitude reduction factor (S02 parameter) was used as a scaling factor for the proportions of UIV and UVI compounds in the samples. Although the sum of the two S02 values does not necessarily add up to 1, this is a reasonable approach, at least for the UVI surface complex, whose separate spectrum was fit with a S02 value equal to 1. Fits for U-FeS2 spectra were obtained using k-weighting of both 2 and 3, with the R range constrained from 1 to 4.8. As can be observed from Figure 4a-c and the tables and figures presented in the Supporting Information, the combination of a nanoparticulate UO2 and a uranyl structure, linked to a surface complex on FeOOH, provided good fits to all data sets. The different parameters are quite consistent throughout the experimental samples. Based on the S02 value of the uranyl unit, the ratio of UIV/UVI in the samples can be

FIGURE 4. (A) U LIII-edge EXAFS Χ(k) · k3 spectra and (B) magnitude and (C) imaginary part of the Fourier transform for U-FeS2 samples S1-S5. Open symbols represent experimental data; solid lines are best fits. 4214

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estimated (Table 1). It appears that, within the error margins, this ratio is quite consistent with the results from LCF of the XANES spectra. Thus, from XAS analysis, it appears that, in the U-FeS2 batch experiment, both UIV and UVI species are observed. Although the solution phase might be dominated by UVI in all systems, the relative amounts of the two oxidation states in the solid phase seem to be largely determined by the presence of DOC, which tends to stabilize the UVI oxidation state. The impact of humic substances on the reduction and/ or oxidation of U and on the complexation of reduced UIV in natural systems is yet poorly understood. It was observed that UIV-humic substance complexes, formed during bioreduction, also greatly increased the reoxidation rate to UVI (because of the prevention of the precipitation of reduced UIV), even compared to the presence of 0.03 mol · L-1 NaHCO3 solution (41). When humics were added separately to reduced UIV solids (obtained by the reduction of UVI without humics), the oxidation rates of UIV were comparable with those added with water or 0.03 mol · L-1 NaHCO3 (41). Dissolution experiments in the presence of humic acids then again showed that the presence of HCO3- increased the proportion of UVI compared to UIV in the solutions (30). These results are in line with past studies (7, 45) and with the present study. As a conclusion, it can be stated that, in U-FeS2 batch systems, at least partial reduction of UVI to UIV occurs, forming a uraninite-like structure (possibly UO2 nanoparticulates). In the solution phase, no solubility limit was observed, and the solution-phase U concentration increased in the presence of DOC. The amount of UIV in the solid phase was the highest in systems with SBCW (absence of DOC) background electrolyte, and tended to increase with increasing equilibration time. However, in all samples, both UVI and UIV coordination environments were detected. The influence of HCO3- and DOC in the batch systems is quite complex and probably two-fold: although promoting oxidation of the FeS2 surface (and therefore reduction rates for UVI) because of complexation with Fe2+, both ligands also have a positive influence on the stabilization of UVI, either by direct complexation (in the case of HCO3-) or by prevention of UO2(s) precipitation (in the case of DOC). The coexistence of multiple U valence states in (near-)natural systems exhibiting excess reducing capacity renders a scientifically sound appraisal of U migration rates under repository conditions a very complex task.

Acknowledgments This work was performed in close cooperation with, and with the financial support of, ONDRAF/NIRAS, the Belgian Agency for Radioactive Waste and Fissile Materials, as part of the program on geological disposal of high-level/longlived radioactive waste that is carried out by ONDRAF/NIRAS and of the EC in the framework of the 6FP FUNMIG project. NOW/FWO Vlaanderen and the European Synchrotron Radiation Facility provided facilities and financial support for performing measurements at DUBBLE and ROBL, respectively. We thank S. Nikitenko and A. Scheinost for their help during XAFS measurements and A. Scheinost for his kind help in U spectra interpretation. Three anonymous reviewers are also thanked for their fruitful suggestions to improve this article.

Supporting Information Available Details on experimental data, solution analysis of batch isotherm experiments, and X-ray absorption near-edge spectroscopy (XANES) and extended X-ray absorption fine structure (EXAFS) methods and analysis. This material is available free of charge via the Internet at http://pubs.acs.org.

Literature Cited (1) De Craen, M.; Wang, L.; Van Geet, M.; Moors, H. Geochemistry of the Boom clay pore water at the Mol site; SCK•CEN Scientific Report BLG-990; SCK•CEN: Mol, Belgium, 2004. (2) 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. (3) Aubriet, H.; Humbert, B.; Perdicakis, M. Interaction of U(VI) with pyrite, galena and their mixtures: A theoretical and multitechnique approach. Radiochim. Acta 2006, 94, 657–663. (4) Eglizaud, N.; Miserque, F.; Simoni, E.; Schlegel, M.; Descostes, M. Uranium(VI) interaction with pyrite (FeS2): Chemical and spectroscopic studies. Radiochim. Acta 2006, 94, 651–656. (5) Scott, T. B.; Riba Tort, O.; Allen, G. C. Aqueous uptake of uranium onto pyrite surfaces; reactivity of fresh versus weathered material. Geochim. Cosmochim. Acta 2007, 71, 5044–5053. (6) Descostes, M.; Schlegel, M. L.; Eglizaud, N.; Descamps, F.; Miserque, F.; Simoni, E. Uptake of uranium and trace elements in pyrite (FeS2) suspensions. Geochim. Cosmochim. Acta 2010, 74, 1551–1562. (7) Dele´caut, G. The geochemical behaviour of uranium in the Boom clay. Ph.D. Dissertation, Universite´ Catholique de Louvain-LaNeuve, Louvain-La-Neuve, Belgium, 2004. (8) Maes, N.; Wang, L.; Dele´caut, G.; Beauwens, T.; Van Geet, M.; Put, M.; Weetjens, E.; Marivoet, J.; van der Lee, J.; Warwick, P.; Hall, A.; Walker, G.; Maes, A.; Bruggeman, C.; Bennett, D.; Hicks, T.; Higgo, J.; Galson, D. Migration case study: Transport of radionuclides in areducingclaysediment(TRANCOM-II);FinalReport,EUR21022EN; European Commission: Luxembourg, 2004. (9) Liger, E.; Charlet, L.; Van Cappellen, P. Surface catalysis of uranium(VI) reduction by iron(II). Geochim. Cosmochim. Acta 1999, 63, 2939–2955. (10) Behrends, T.; Van Cappellen, P. Competition between enzymatic and abiotic reduction of uranium(VI) under iron reducing conditions. Chem. Geol. 2005, 220, 315–327. (11) O’Loughlin, E. J.; Kelly, S. D.; Cook, R. E.; Csencsits, R.; Kemner, K. M. Reduction of Uranium(VI) by Mixed Iron(II)/Iron(III) Hydroxide (Green Rust): Formation of UO2 Nanoparticles. Environ. Sci. Technol. 2003, 37, 721–727. (12) Ilton, E. S.; Boily, J.-F.; Buck, E. C.; Skomurski, F. N.; Rosso, K. M.; Cahill, C. L.; Bargar, J. R.; Felmy, A. R. Influence of dynamical conditions on the reduction of UVI at the magnetitesolution interface. Environ. Sci. Technol. 2010, 44, 170–176. (13) O’Loughlin, E. J.; Kelly, S. D.; Kemner, K. M. XAFS investigation of the interactions of UVI with secondary mineralization products from the bioreduction of FeIII oxides. Environ. Sci. Technol. 2010, 44, 1656–1661. (14) Guillaumont, R.; Fangha¨nel, T.; Fuger, J.; Grenthe, I.; Neck, V.; Palmer, D. A.; Rand, M. H. Update on the Chemical Thermodynamics of Uranium, Neptunium, Plutonium, Americium and Technetium; OECD Nuclear Energy Agency/Elsevier: Amsterdam, 2003. (15) Lenhart, J. J.; Honeyman, B. D. Uranium(VI) sorption to hematite in the presence of humic acid. Geochim. Cosmochim. Acta 1999, 63, 2891–2901. (16) Lovley, D. R.; Phillips, E. J. P.; Gorby, Y. A.; Landa, E. R. Microbial reduction of uranium. Nature 1991, 350, 413–415. (17) Lovley, D. R.; Phillips, E. J. P. Reduction of uranium by Desulfovibrio desulfuricans. Appl. Environ. Microbiol. 1992, 58, 850–856. (18) Luo, J.; Weber, F.-A.; Cirpka, O. A.; Wu, W.-M.; Nyman, J. L.; Carley, J.; Jardine, P. M.; Criddle, C. S.; Kitanidis, P. K. Modeling in-situ uranium(VI) bioreduction by sulfate-reducing bacteria. J. Contam. Hydrol. 2007, 92, 129–148. (19) Boivin-Jahns, V.; Ruimy, R.; Bianchi, A.; Daumas, S.; Christen, R. Bacterial diversity in a deep-subsurface clay environment. Appl. Environ. Microbiol. 1996, 62, 3405–3412. (20) Descostes, M.; Mercier, F.; Beaucaire, C.; Zuddas, P.; Trocellier, P. Nature and distribution of chemical species on oxidized pyrite surface: Complementarity of XPS and nuclear microprobe analysis. Nucl. Instrum. Methods B 2001, 181, 603–609. (21) Webb, S. SixPACK (Sam’s Interface for XAS Analysis Package); Stanford Synchrotron Radiation Source: Stanford, CA, 2007. Available at http://ssrl.slac.stanford.edu/∼swebb/sixpack.htm (accessed Feb 23, 2009). (22) Newville, M. IFEFFIT: Interactive XAFS analysis and FEFF fitting. J. Synchrotron Radiat. 2001, 8, 322–324. (23) Ravel, B.; Newville, M. ATHENA, ARTEMIS, HEPHAESTUS: Data analysis for X-ray absorption spectroscopy using IFEFFIT. J. Synchrotron Radiat. 2005, 12, 537–541. VOL. 44, NO. 11, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

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(24) Ankudinov, A. L.; Ravel, B.; Rehr, J. J.; Conradson, S. D. Realspace multiple-scattering calculation and interpretation of x-rayabsorption near-edge structure. Phys. Rev. B 1998, 58, 7565– 7576. (25) Rehr, J. J.; Albers, R. C. Theoretical approaches to X-ray absorption fine structure. Rev. Mod. Phys. 2000, 72, 621–654. (26) Higgo, J. J.; Kinniburgh, D.; Smith, D.; Tipping, E. Complexation of Co2+, Ni2+, UO22+ and Ca2+ by humic substances in groundwaters. Radiochim. Acta 1993, 61, 91–103. (27) Czerwinski, K. R.; Buckau, G.; Scherbaum, F.; Kim, J. I. Complexation of the uranyl ion with aquatic humic acid. Radiochim. Acta 1994, 65, 111–119. (28) Ranville, J.; Hendry, M. J.; Reszat, Th. N.; Xie, Q.; Honeyman, B. D. Quantifying uranium complexation by groundwater dissolved organic carbon using asymmetrical flow field-flow fractionation. J. Contam. Hydrol. 2007, 91, 233–246. (29) Warwick, P.; Evans, N.; Hall, A.; Walker, G.; Steigleder, E. Stability constants of U(VI) and U(IV)-humic acid complexes. J. Radioanal. Nucl. Chem. 2005, 266, 179–190. (30) Luo, W.; Gu, B. Dissolution and mobilization of uranium in a reduced sediment by natural humic substances under anaerobic conditions. Environ. Sci. Technol. 2009, 43, 152–156. (31) De Craen, M.; Van Geet, M.; Wang, L.; Put, M. High sulphate concentrations in squeezed Boom Clay pore water: Evidence of oxidation of clay cores. Phys. Chem. Earth 2004, 29, 91–103. (32) Descostes, M.; Beaucaire, C.; Mercier, F.; Savoye, S.; Sow, J.; Zuddas, P. Effect of carbonate ions on pyrite (FeS2) dissolution. Bull. Soc. Ge´ol. Fr. 2002, 173 (3), 265–270. (33) Parkhurst, D. L.; Appelo, C. A. J. User’s Guide to PHREEQC (Version 2)sA Computer Program for Speciation, Batch-Reaction, One-Dimensional Transport, and Inverse Geochemical Calculations; Water-Resources Investigations Report 99-4259; U.S. Geological Survey, U.S. Department of the Interior: Denver, CO, 1999. (34) Gimenez, J.; de Pablo, J.; Torrero, M. E.; Casas, I. Conceptual and mathematical model for the UO2(s) dissolution in brines under different redox conditions. Radiochim. Acta 1997, 78, 21–25. (35) Pierce, E. M.; Icenhower, J. P.; Serne, R. J.; Catalano, J. G. Experimental determination of UO2(cr) dissolution kinetics:

4216

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 44, NO. 11, 2010

(36)

(37)

(38)

(39)

(40)

(41) (42) (43)

(44) (45)

Effects of solution saturation state and pH. J. Nucl. Mater. 2005, 345, 206–218. Kelly, S. D.; Kemner, K. M.; Carley, J.; Criddle, C.; Jardine, P. M.; Marsh, T. L.; Phillips, D.; Watson, D.; Wu, W.-M. Speciation of uranium in sediments before and after in situ biostimulation. Environ. Sci. Technol. 2008, 42, 1558–1564. Ulrich, K.-U.; Singh, A.; Schofield, E. J.; Bargar, J. R.; Veeramani, H.; Sharp, J. O.; Bernier-Latmani, R.; Giammar, D. Dissolution of biogenic and synthetic UO2 under varied reducing conditions. Environ. Sci. Technol. 2008, 42, 5600–5606. Ro¨llin, S.; Spahiu, K.; Eklund, U. B. Determination of dissolution rates of spent fuel in carbonate solutions under different redox conditions with a flow-through experiment. J. Nucl. Mater. 2001, 297, 231–243. Schofield, E.; Veeramani, H.; Sharp, J.; Suvorova, E.; BernierLatmani, R.; Mehta, A.; Stahlman, J.; Clark, D. L.; Conradson, S. D.; Webb, S.; Bargar, J. R. Structure of biogenic uraninite produced by Shewanella oneidensis, strain MR-1. Environ. Sci. Technol. 2008, 42, 7898–7904. Wan, J.; Tokunaga, T. K.; Brodie, E.; Wang, Z.; Zheng, Z.; Herman, D.; Hazen, T. C.; Firestone, M. K.; Sutton, S. R. Reoxidation of bioreduced uranium under reducing conditions. Environ. Sci. Technol. 2005, 39, 6162–6169. Gu, B.; Yan, H.; Zhou, P.; Watson, B. D. Natural humics impact uranium bioreduction and oxidation. Environ. Sci. Technol. 2005, 39, 5268–5275. Hudson, E. A.; Rehr, J. J.; Bucher, J. J. Multiple-scattering calculations of the uranium L(3)-edge x-ray absorption nearedge structure. Phys. Rev. B 1995, 52 (19), 13815–13826. Kelly, S. D.; Hesterberg, D.; Ravel, B. Analysis of soils and minerals using X-ray absorption spectroscopy. In Methods of Soil Analysis. Part 5: Mineralogical Methods; Ulery, A. L., Drees, R., Eds.; Soil Science Society of America: Madison, WI, 2008; pp 387-494. Suzuki, Y.; Kelly, S. D.; Kemner, K. M.; Banfield, J. F. Radionuclide contaminationsNanometre-size products of uranium bioreduction. Nature 2002, 419 (6903), 134–134. Cachoir, C.; Lemmens, K.; Van den Berghe, S.; Van Iseghem, P. UO2 dissolution in Boom Clay conditions. J. Nucl. Mater. 2003, 321, 49–59.

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