Reoxidation of Reduced Uranium with Iron(III) (Hydr)Oxides under

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Environ. Sci. Technol. 2005, 39, 2059-2066

Reoxidation of Reduced Uranium with Iron(III) (Hydr)Oxides under Sulfate-Reducing Conditions R A J E S H K . S A N I , † B R E N T M . P E Y T O N , * ,† ALICE DOHNALKOVA,‡ AND JAMES E. AMONETTE‡ Department of Chemical Engineering, Center for Multiphase Environmental Research, Washington State University, Dana Hall Room 118, Pullman, Washington 99164-2710, and Fundamental Sciences Directorate, Pacific Northwest National Laboratory, 902 Battelle Boulevard, Richland, Washington 99354-0999

In cultures of Desulfovibrio desulfuricans G20 the effects of iron(III) (hydr)oxides (hematite, goethite, and ferrihydrite) on microbial reduction and reoxidation of uranium (U) were evaluated under lactate-limited sulfate-reducing conditions. With lactate present, G20 reduced U(VI) in both 1,4-piperazinediethanesulfonate (PIPES) and bicarbonate buffer. Once lactate was depleted, however, microbially reduced U served as an electron donor to reduce Fe(III) present in iron(III) (hydr)oxides. With the same initial amount of Fe(III) (10 mmol/L) for each iron(III) (hydr)oxide, reoxidation of U(IV) was greater with hematite than with goethite or ferrihydrite. As the initial mass loading of hematite increased from 0 to 20 mmol of Fe(III)/L, the rate and extent of U(IV) reoxidation increased. Subsequent addition of hematite [15 mmol of Fe(III)/L] to stationary-phase cultures containing microbially reduced U(IV) also resulted in rapid reoxidation to U(VI). Analysis by U L3-edge X-ray absorption near-edge spectroscopy (XANES) of microbially reduced U particles yielded spectra similar to that of natural uraninite. Observations by high-resolution transmission electron microscopy, selected area electron diffraction, and energy-dispersive X-ray spectroscopic analysis confirmed that precipitated U associated with cells was uraninite with particle diameters of 3-5 nm. By the same techniques, iron sulfide precipitates were found to have a variable Fe and S stoichiometry and were not associated with cells.

such a process makes reductive precipitation more desirable. In situ microbial reduction of U(VI) to form U(IV) precipitates of uraninite may be an attractive alternative strategy for remediation of U-contaminated subsurface environments (5-7). It is vital, therefore, to identify and characterize processes that control the stability of U to restrict its environmental risk. Dissimilatory metal-reducing bacteria (DMRB) can precipitate U in the subsurface through enzymatic reduction. One group of DMRB is the sulfate-reducing bacteria (SRB), which may play a very important role in U(VI) reduction processes in subsurface soils and groundwaters (8). In addition to U(VI), SRB can enzymatically reduce Cr(VI), Mn(VI), Tc(VII), and Fe(III). Furthermore, microbial reduction of iron(III) (hydr)oxides, either enzymatically or indirectly by H2S production (9-11), can strongly influence the geochemistry of anaerobic soil and sedimentary environments, as well as the persistence and mobility of various types of organic and inorganic contaminants (12, 13). Using SRB, researchers have shown U(VI) reduction in the absence of soil minerals, under nongrowth conditions (14 and references therein). However, knowledge of U(VI) reduction by SRB under microbial growth conditions is limited (5). In situ stimulation of anaerobic microbial metaltransformation processes may be an effective treatment alternative to immobilize heavy metals and radionuclides. More research is needed, however, to better understand interactions of SRB, metal and radionuclide contaminants, and mineral phases in the subsurface. Recent results with Desulfovibrio desulfuricans G20 suggest that, for iron(III) (hydr)oxides (hematite, goethite, and ferrihydrite) with equal surface area, sulfate and U(VI) were reduced concomitantly (5). In lactate-limited systems, after depletion of all lactate, soluble U(VI) concentrations increased and appeared to depend on the concentration of Fe(III) present in the system. On the basis of these results, we hypothesized that microbially reduced U can act as an electron donor to reduce Fe(III) present in iron(III) (hydr)oxides and that the amount of U(IV) reoxidation depends on the form, as well as the concentration, of Fe(III) present. To test these hypotheses, laboratory experiments were performed to evaluate the effects of redox-sensitive and -insensitive aquifer materials (hematite, goethite, ferrihydrite, and quartz) commonly found in subsurface soils and sediments on the microbial reduction of U(VI) under lactatelimited, sulfate-reducing conditions.

Experimental Methods Introduction Uranium (U) is the most common radionuclide in soils, sediments, and groundwater at U.S. Department of Energy sites and therefore is of particular environmental concern (1). The most common form of U in groundwater is typically U(VI), which is present either as the uranyl cation, schoepite, or as anionic carbonate complexes (2, 3). Removal of U(VI) from solution can occur by sorption, precipitation as a U(VI) compound, or reductive precipitation as a U(IV) compound. Although sorption of U(VI) can control aqueous concentrations of U under oxidizing conditions (4), the reversibility of * Corresponding author phone: (509)335-4002; fax: (509)335-4806; e-mail: [email protected]. † Washington State University. ‡ Pacific Northwest National Laboratory. 10.1021/es0494297 CCC: $30.25 Published on Web 03/04/2005

 2005 American Chemical Society

Bacteria, Materials, and Minerals. Desulfovibrio desulfuricans G20 (subsequently referred to as G20) used in the study was a gift of J. Wall, University of MissourisColumbia (Columbia, MO), derived from D. desulfuricans G100A (15). G20 was maintained in a lactate-C medium modified to metal toxicity medium (MTM) (16). Uranium was purchased as UO2Cl2‚3H2O from Bodman Industries (Aston, PA). The iron(III) (hydr)oxides hematite (R-Fe2O3), goethite (R-FeOOH), and ferrihydrite (Fe5O7OH‚4H2O) were synthesized as described by Schwertmann and Cornell (17). A Coulter SA 3100 BET analyzer (by N2 sorption) was used to determine the specific surface areas (18) of the hematite, goethite, ferrihydrite, and quartz (R-SiO2, 212-300 µm), which were found to be 30, 52, 180, and 0.02 m2/g, respectively. Due to potential phase modifications that can occur during autoclaving, the iron(III) (hydr)oxides were not autoclaved (13). Prior to use, the iron(III) (hydr)oxides were heat-treated in presterilized VOL. 39, NO. 7, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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culture tubes by incubation in an oven at 80 °C for 16 h. No measurable change in specific surface area occurred as a result (5). U(VI) Reduction Experiments with Minerals. Washed cells for inoculation were prepared as described previously (5). Volumes of 120 mL of lactate-limited MTM with lactate and sulfate at 30 and 20 mM, respectively, with 1,4-piperazinediethanesulfonate (PIPES) or bicarbonate buffer each at 30 mM (pH 7) in 150-mL serum bottles were autoclaved. With each set of experiments, mineral- and U(VI)-free controls were used. A filtered (0.2 µm, Gelman Acrodisc) stock solution (42 mM) of UO2Cl2‚3H2O was aseptically added to the serum bottles to give the desired U(VI) concentrations. Samples were withdrawn aseptically to measure initial concentrations of U(VI), lactate, and sulfate. Heat-treated hematite, goethite, and ferrihydrite at equal Fe(III) concentration (10 mmol/L) were added to the serum bottles containing U(VI) and MTM with PIPES or bicarbonate buffer. In separate experiments, hematite at different Fe(III) mass loading (030 mmol/L) was also added to serum bottles containing MTM with PIPES buffer. In addition, each serum bottle was supplemented with 1.52 g of quartz. The serum bottles containing PIPES or bicarbonate buffer were then purged with ultrapure N2(g) or O2-free N2/CO2 (80:20), respectively, for 10 min, sealed with butyl rubber septa, capped and crimped with aluminum seals, and pressurized at 68.9 kPa (10 lb/in.2) above atmospheric pressure. Serum bottles were incubated at 25 °C on an orbital shaker (Lab-Line Instruments, Inc.) operated at 125 rpm. Samples were collected over a 5-day period to monitor establishment of U(VI) equilibrium between the soluble and solid phases under abiotic conditions. After abiotic sorption of U(VI) onto the minerals, washed cells of G20 were injected into the serum bottles to give a final concentration of 3 mg/L cell protein. After inoculation, serum bottles were again incubated at 25 °C on an orbital shaker at 125 rpm. Periodically, 0.3-mL samples were aseptically removed by syringe and needle and analyzed for soluble concentrations of U and sulfide. In addition, at the end of the each experiment, soluble concentrations of lactate, sulfate, and acetate were determined. Uranium Reoxidation Experiments. G20 was grown with U(VI) (180 µM) in 12 serum bottles containing lactate-limited MTM in bicarbonate buffer (30 mM, pH 7). After the reduction of U(VI), three serum bottles were amended with heat-treated hematite [15 mmol of Fe(III)/L] and three bottles were amended with quartz (hematite-free controls) in a glovebox (Model 1025, Forma Scientific Inc., OH) under an O2-free atmosphere. In addition to hematite-free controls, heat-killed cell controls were also included. In separate experiments, nanoparticles of U(IV) were synthesized as described by Beyenal et al. (19). In a glovebox, chemically precipitated U(IV) nanoparticles (0.8 µmol of U) were suspended with either 15 mM FeCl3 or 15 mmol of Fe(III)/L as hematite in serum bottles containing 10 mL of bicarbonate buffer (30 mM, pH 7). Hematite-, FeCl3-, and U(IV)-free controls were also used. Periodically, 0.3-mL samples were aseptically removed by syringe and needle and analyzed for soluble U concentrations. Final Fe(II) concentrations were determined spectrophotometrically (Milton Roy Co. Spectronic Genesys 5, Rochester, NY) by the ferrozine method (20). Analytical Methods: (A) Soluble Uranium. To monitor abiotic U(VI) sorption onto minerals, samples were taken before inoculation and centrifuged at 5000g for 5 min in a microcentrifuge. The supernatant was filtered (0.2 µm, Gelman Acrodisc) and U(VI) concentrations were measured as previously described (5, 21) with a kinetic phosphorescence analyzer (KPA-11, Chemcheck Instruments Inc., Richland, WA). Calibration was done with UO2Cl2‚3H2O solutions of 0-160 nM. Because the KPA-11 allows detection of U(VI) concentrations as low as 0.04 nM with a precision of (5%, the estimated 2060

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detection limit in this study for 1000-fold dilutions was 40 nM. After the appearance of biotic activity, samples were filtered (0.2 µm) and analyzed immediately for U(VI). (B) Anions and Soluble Sulfide. Samples for lactate, sulfate, and acetate were filtered (0.2 µm), and concentrations were determined on a Dionex DX-500 ion chromatograph (IonPac AS11-HC 4-mm column, conductivity detection) eluted with a NaOH gradient (1-100 mM) with an estimated error of (5%. The detection limit was 3 mg/L for each anion. Bacterial activity was monitored by measurement of biogenic sulfide formation in treatments (22). The soluble sulfide concentrations in filtered (0.2 µm) samples with zinc acetate added (10% w/v) were determined spectrophotometrically, by the methylene blue method (Hach Co., Loveland, CO). The detection limit for sulfide was estimated to be 3 nM. X-ray Absorption Near-Edge Spectroscopy of Microbially Reduced U. At the end of the experiments, the resulting black precipitates from G20-treated hematite (25 mmol/L) and U(VI) were separated from the aqueous phase by centrifugation at 10000g for 10 min under anoxic conditions. After the supernatant was decanted, the black precipitates were dried in a glovebox (Model 1025, Forma Scientific Inc., OH) under an O2-free N2/CO2/H2 (90:5:5) atmosphere. Dried solid subsamples were packed between Kapton windows in a multispecimen Al sample holder modified to fit into a Lytle detector chamber. To avoid oxidation, the specimenmounting process was carried out in a glovebox, and the sample holder was placed inside a BBL Gas-Pak (Becton Dickinson, Franklin Lakes, NJ) container for shipping and storage at the synchrotron before analysis. In addition to black precipitates, an aqueous-phase sample was filtered twice through a 0.2-µm pore-size membrane filter. The resulting filtered solution was again filtered through a 1 kDa cutoff filter membrane (0.6-nm pore size, Gelman Acrodisc), and the membranes were dried and placed in the sample holder as described above. XANES spectra at the U L3-edge were collected in fluorescence mode at the Pacific Northwest Consortium Collaborative Access Team (PNC-CAT) bending-magnet beamline (20-BM-B) located at the Advanced Photon Source (Argonne, IL). Fluorescence data were collected by use of a scintillation detector with a single-channel analyzer to define the energy window. A strontium filter was placed between the sample and the detector to minimize background from the incident beam. Natural uraninite (nominally UO2) and UO2Cl2‚3H2O samples served as U(IV) and U(VI) reference standards, respectively, and were analyzed in transmission mode. During the analysis, the Lytle detector chamber was continuously flushed with He(g) to prevent oxidation of the samples by air. The resulting spectra were normalized for comparative purposes. Transmission Electron Microscopy. The whole embedding procedure, as well as thin sectioning, was conducted in a glovebox (Ar/H2 95:5; Coy Laboratory Products, Inc.). The visible black precipitates resulting from G20-treated hematite (25 mmol/L) and U(VI) were briefly fixed in 2.5% glutaraldehyde and washed in anoxic deionized water followed by a gradual ethanol dehydration series and infiltration in LR White embedding resin. Cured blocks were sectioned to 70 nm on an ultramicrotome (Leica Ultracut UCT), and sections were mounted on 200-mesh copper grids coated with Formvar support film sputtered with carbon. Sections were examined on a JEOL 2010 high-resolution transmission electron microscope (HR-TEM) equipped with a LaB6 filament operating at 200 kV with resolution of 0.19 nm. Elemental analysis was performed on an Oxford energy-dispersive X-ray spectroscopy (EDS) system equipped with a SiLi detector coupled to the TEM, and spectra were analyzed with ISIS software. Images were digitally collected and analyzed on Gatan’s Digital Micrograph. The d-spacings obtained from

FIGURE 1. (A) Concentration profiles of soluble U(VI) during abiotic sorption onto hematite, goethite, ferrihydrite, and quartz in medium containing PIPES buffer, and reduction by D. desulfuricans G20. (B) Concentration profiles of soluble sulfide during the growth of D. desulfuricans G20. the selected area electron diffraction (SAED) ring patterns were evaluated by Desktop Microscopist (Lacuna) software. Statistical Analysis. Each set of experiments was carried out in duplicate or triplicate and repeated. In each experiment, duplicate treatment profiles were similar in soluble U(VI) and sulfide concentrations; however, the length of the lag time was somewhat variable. One-way analysis of variance (ANOVA) was used to determine if there were statistically significant differences in G20 lag times among controls and treatments containing U with and without iron(III) (hydr)oxides. The threshold level of statistical significance for this study was R ) 0.05.

Results Reduction of U(VI) in the Presence of Iron(III) (hydr)Oxides. Reduction of U(VI) under lactate-limited conditions was tested with 10 mmol/L Fe(III) for each iron(III) (hydr)oxide. As shown in Figure 1A, the first 5 days of the experiments consisted of abiotic equilibration of U(VI) in MTM containing PIPES buffer (30 mM, pH 7) with iron(III) (hydr)oxides (termed “minerals” hereafter) and quartz. During abiotic incubation, significant amounts of U(VI) sorbed to the minerals with greater sorption to ferrihydrite than to hematite and goethite. In the mineral-free controls, U(VI) concentration decreased only slightly during the first 2 days of incubation. After 5 days, four of the five treatments and an additional treatment [U(VI)- and mineral-free control] were inoculated with washed G20. U(VI) concentrations decreased gradually, but once growth began, U(VI) concentrations decreased sharply. The lag times before initiation of growth were longer with higher aqueous U(VI) concentrations. In contrast to the mineral-free control and ferrihydrite treat-

FIGURE 2. Concentration profiles of soluble U(VI) during abiotic sorption onto hematite at different concentrations in medium containing PIPES buffer, and reduction by D. desulfuricans G20. ments, after active U(VI) reduction was complete, the soluble U(VI) concentrations in the hematite and goethite treatments gradually increased. As an indicator of biotic activity, concentrations of sulfide were monitored (Figure 1B). During growth and active reduction of U(VI), the concentration of sulfide increased. When minerals were present, however, the concentrations of sulfide decreased after the active growth phase, probably as a result of reaction with Fe(II) or solid-phase Fe(III). No decrease in sulfide concentration was observed in the U(VI)and mineral-free treatments. During the course of the experiments, the color of the solids changed to black, suggesting formation of iron sulfides and uraninite. Since hematite yielded the greatest reoxidation of U, the effects of different mass loadings of hematite [0-30 mmol of Fe(III)/L] on the profiles of soluble U(VI) concentrations in PIPES buffer were examined (Figure 2). As before, trends in the U(VI) concentration profiles showed that the lag period between inoculation and the onset of growth depended on the soluble concentration of U(VI). With hematite at 0 and 5 mmol Fe/L, no increase in U(VI) concentration was observed after active U(VI) reduction. With hematite at 10-20 mmol of Fe/L, however, the soluble U(VI) concentrations increased after the minimum was reached, and the amount of the increase was correlated with the amount of hematite initially present (Figure 2A). In contrast, at the highest levels of hematite, the soluble U(VI) concentrations increased for a period of 15 days, and then again decreased to 20 and 30 µM, respectively (Figure 2B). During the course of the experiment, the concentration profiles for sulfide were also monitored and were similar to those shown in Figure 1B (data not shown). The experiments with 10 mmol/L Fe(III) each for hematite, goethite, and ferrihydrite in a lactate-limited medium were VOL. 39, NO. 7, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 4. (A) TEM image of an unstained thin section of D. desulfuricans G20 treated with U(VI) and hematite (25 mmol/L) showing an oblique section of a bacterium. (B-E) Evident precipitates of biogenic uraninite associated with the cell surfaces were confirmed by EDS and SAED. Copper background signal originates from the Cu grid.

FIGURE 3. (A) Concentration profiles of soluble U(VI) during abiotic sorption onto hematite, goethite, ferrihydrite, and quartz in medium containing bicarbonate buffer, and reduction by D. desulfuricans G20. (B) Concentration profiles of soluble sulfide during the growth of D. desulfuricans G20. repeated in a bicarbonate buffer (30 mM, pH 7). In some respects, trends in the U(VI) concentration profiles (Figure 3A) were similar to those observed for the medium containing PIPES buffer (Figure 1A). There were three exceptions, however. First, during the first 5 days of abiotic incubations, significantly less sorption of U(VI) to minerals occurred. Second, after inoculation, no lag times of G20 activity were observed. Third, after active U(VI) reduction, with all mineral treatments, the soluble U(VI) concentrations increased gradually over time. The concentration profiles of soluble sulfide followed the same trends (Figure 3B) seen with the PIPES buffer system (Figure 1B). Taken together, Figures 1-3 suggest that U(IV) may possibly act as an electron donor to reduce Fe(III) present in iron(III) (hydr)oxides. Characterization of Microbially Reduced U: (A) XANES Spectra. The U L3-edge XANES spectra for U(VI) treated with G20 and hematite (25 mmol/L) were similar to that of natural uraninite (data not shown) but had distinct features, notably a much more intense and slightly broader white line (the intense absorption peak located at or just above the absorption edge, i.e., at about 17 125 eV for the U L-edge). These features were more vivid in the spectrum of aqueous reduced uranium filtered on a 1-kDa cutoff membrane. These spectra were easily distinguishable from that of an untreated natural uraninite by having slightly higher edge energy, a slightly less intense white line just above the edge, and a greater intensity in the region 15-35 eV above the edge. The unique features of the biotically reduced U spectra, notably the intense white line, most likely indicate that the reduced-U particles are very small, on the order of a few nanometers. The small particle size apparently eliminates some of the destructive interference in the electron waves giving rise to the absorption spectrum by increasing the fraction of surface 2062

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FIGURE 5. Reoxidation of abiotically generated U(IV) nanoparticles by FeCl3 and hematite in bicarbonate buffer (30 mM, pH 7) under anaerobic conditions. U atoms and thereby decreasing the number of U nearest neighbors (23). (B) Transmission Electron Microscopy, Energy-Dispersive X-ray Spectroscopy, and Selected Area Electron Diffraction Pattern. Figure 4A shows a TEM image of an unstained thin section of a G20 cell harvested at the end of the experiment obtained from incubation of G20 with U(VI) and hematite (25 mmol/L). To identify the elemental composition of the precipitates in the periplasm, several EDS spectra were obtained. The spectra of randomly selected precipitates within the cell membrane contained exclusively U and O peaks, with no detectable Fe or S concentrations (as shown in Figure 4B-E). High-resolution transmission electron microscopic images and SAED patterns from these cell-associated precipitates showed that the diameters of individual uraninite particles were in the range of 3-5 nm and occurred as both discrete and aggregated particles (data not shown). Reoxidation of Reduced U by FeCl3 and Hematite. Figures 1-3 suggest that, in PIPES or bicarbonate buffer, U(IV) can potentially act as an electron donor to reduce Fe(III) present in iron(III) (hydr)oxides. For confirmation of reoxidation of reduced U by Fe(III), two types of experiments were performed. First, nanoparticles of U(IV) (0.8 µmol) (19) were suspended with either 15 mM FeCl3 or 15 mmol Fe(III)/L hematite in serum bottles containing 10 mL of bicarbonate buffer (30 mM, pH 7) under anoxic conditions. Figure 5 shows that, in contrast to hematite treatments, U(IV) nanoparticles

FIGURE 6. Soluble U(VI) concentration profiles comparing reoxidation by addition of hematite [15 mmol of Fe(III)/L] in the presence and absence of SRB activity in MTM containing bicarbonate buffer (30 mM, pH 7). were completely oxidized by anoxic Fe(III) chloride within 24 h. The stoichiometry of U(IV) consumption and Fe(II) production in the FeCl3 treatment was consistent with the following equation:

U(IV) + 2Fe(III) / U(VI) + 2Fe(II)

(1)

In a second experiment, after the reduction of U(VI) by G20, hematite [15 mmol of Fe(III)/L] was added to the serum bottles with and without SRB activity. Figure 6 shows that in the treatments with SRB activity, added hematite reoxidized microbially reduced U, whereas in the treatments with no added hematite, U remained reduced. It can also be seen from Figure 6 that hematite added in heat-killed cell controls also reoxidized a small amount of U. These results are in contrast to abiotic conditions, where no reaction between hematite and U(IV) was observed (Figure 5). No reoxidation of U was observed in any Fe(III)-free control under biotic or abiotic conditions.

Discussion The results provide evidence that D. desulfuricans G20 reduced U(VI) to form nanoparticulate uraninite, which was observed both in the cell periplasm and in the bulk liquid. Further, results indicate that the type and amount of iron(III) (hydr)oxide present and the type of pH buffer (PIPES or bicarbonate) had significant effects on the rate of U(VI) reduction. After depletion of lactate required for microbial U(VI) reduction, the rate and extent of U(IV) reoxidation were dependent on the type and amount of iron(III) (hydr)oxide present. These results suggest that application of longterm bioimmobilization of U must consider additional complexity in processes and reactions, since sulfate and SRB are commonly found in many U-contaminated aquifers. Removal of Aqueous-Phase U(VI). Prior to inoculation, some removal of U(VI) from the aqueous phase occurred as a result of adsorption to, or possibly precipitation onto, the iron(III) (hydr)oxide. In the presence of bicarbonate buffer, the sorption of U(VI) onto the minerals was significantly lessened, as observed by others (13, 24, 25), with U(VI) likely kept in solution through the formation of negatively charged uranyl-carbonate complexes such as UO2(CO3)22- and UO2(CO3)34- (13). In the presence of PIPES buffer, however, significantly more adsorption to iron(III) (hydr)oxide occurred, with most of the adsorption of U(VI) occurring within 1 day of U(VI) addition. An apparent equilibrium between the solution and solid phase was achieved within 5 days (Figure 1), and the data corroborate with previous reports

(24, 26, 27) showing less U(VI) sorption to ferrihydrite than hematite and goethite per unit surface area. After abiotic adsorption, experiments were inoculated with G20 and were monitored for aqueous U(VI) and sulfide concentrations. In bicarbonate buffer, the presence or absence of iron(III) (hydr)oxides had no effect on observed lag time for U(VI) reduction by G20 (Figure 3). Our results were consistent with other literature (28-30) that also show no effect of U(VI) on DMRB lag time when a bicarbonate buffer (30 mM, pH 7) was used. Spear et al. (31) reported a slight increase in the lag time (about 3 h) with 1 mM U(VI) in an experiment that contained 20 mM bicarbonate buffer. In contrast, our results with PIPES buffer show that with increased soluble U(VI) concentration, G20 lag times significantly increased (up to 17 days) as compared to U(VI)free controls (p-values of 0.000 33 to 0.008 431). The length of lag time in PIPES buffer was somewhat variable, but similar variability in lag times of D. desulfuricans and Shewanella oneidensis MR-1 has been observed for metals such as Cu(II), Cr(VI), Ni(II), Pb(II), and Zn(II) (5, 16, 32-35). The decrease in lag time in bicarbonate buffer, as compared to PIPES buffer, may be the result of lower U(VI) toxicity to G20 due to the formation of uranyl carbonate complexes. U(VI) was removed from solution by precipitation as a U(IV) solid phase. High-resolution TEM revealed the finegrained nature of the U precipitates and yielded lattice-fringe images having d-spacings 0.164, 0.193, 0.273, and 0.316 nm (data not shown). Theses d-spacings were consistent with UO2(s) (JCPDS 41-1442) and comparable to those previously obtained for U reduced by D. desulfuricans ATTC 29577 and Geobacter metallireducens strain GS-15 (6, 54). The individual particle diameters were in the range of 3-5 nm and were typically observed to be aggregated particles. Most of these aggregates occurred in association with cell surfaces. Our results corroborate those of Suzuki et al. (55, 56), who also observed nanometer-sized particles of UO2 directly associated with microbial cells resulting from bacterial U(VI) reduction. With an iron-reducing bacterium, Fredrickson et al. (44) showed that U reduced by Shewanella putrefaciens strain CN32 could be found in the periplasmic space. These observations suggest that most of the uranium reductase activity is associated with the periplasmic space. Uranium Reoxidation. Figures 1-3 show that U(VI) concentrations reached a minimum and then, in the presence of iron(III) (hydr)oxides, began to slowly increase over time. The only exception to this observation was the ferrihydrite experiments with PIPES buffer (Figure 1A). We hypothesize that the observed increase in U(VI) concentration could be the result of three potential mechanisms. The first is that an increase in soluble carbonate concentration stemming from G20 activity, as shown in eq 2, led to increased desorption of U(VI) from mineral and microbial surfaces as carbonate complexes:

2CH3CHOHCOO- + SO42- + H+ f 2CH3COO- +

2CO2 + 2H2O + HS- (2)

This mechanism would have an insignificant influence in 30 mM bicarbonate buffer (Figure 3) but could play a role in the PIPES-buffered systems. Second, the increase could be sorbed U(VI), released when Fe(III) was reduced by biogenic H2S (36) or, less likely, by enzymatic reduction. Third, the increase in soluble U(VI) concentrations could be the result of reoxidation of U(IV) to U(VI), as shown in eq 1, by Fe(III) remaining in the system. Mechanisms 1 and 2 require desorption of U(VI) from a mineral or microbial solid phase to increase the aqueous U(VI) concentration. It can be seen in the bicarbonatebuffered systems (Figure 3A), especially for hematite, that VOL. 39, NO. 7, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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the increase in soluble U(VI) from the minimum concentration at 9-19 days, was about 34 µM. Even if all the initially adsorbed U(VI) was mobilized due to mechanisms 1 and 2, our results indicate that, in addition to desorption, reoxidation of U(IV) to U(VI) was partly responsible for the observed increases in U(VI) concentrations (Figure 3A). The data in Figure 6 in which hematite was added after U(VI) had been reduced clearly show that the observed increases in soluble U(VI) are not the result of a desorption-based mechanism for U(VI) regeneration. These results, analogous to Figures 1-3, support the hypothesis that Fe(III) was responsible for reoxidation of U(IV) to U(VI). Hematite gave the largest increase in U(VI) concentration (Figure 1), so it was examined in more detail by mass loadings that ranged from 0 to 30 mmol of Fe(III)/L (Figure 2). In Figure 2, results show that with Fe(III) in hematite ranging from 0 to 20 mmol/L, two general trends were observed. It can be seen that, for 0 and 5 mmol/L, no increase in U(VI) concentration was observed during the experiments. For concentrations of 10 mmol/L and above, U(VI) concentrations in solution increased significantly after a minimum concentration was achieved. These results suggest that, even under highly reducing conditions, U(IV) may act as an electron donor (by eq 1) to reduce available Fe(III). Since Fe(II) reacts quickly with H2S to form iron sulfide precipitates, examination of eq 1 shows that, in the absence of H2S, Fe(II) would accumulate. This accumulation may affect U(IV) oxidation through the reversible reaction (eq 1) locally approaching equilibrium concentrations or through the passivation of the iron(III) (hydr)oxide surface, as has been seen by others in iron-reducing systems (37-39). Others have reported related observations under different conditions (40-42). Nevin and Lovley (42) showed that addition of U(VI) to bicarbonate-buffered cell suspensions of G. metallireducens in the presence of synthetic poorly crystalline iron(III) oxide stimulated the reduction of iron(III) oxide. Their results suggested that, under anaerobic bicarbonate-buffered conditions, reduced uranium [U(IV)] acted as an electron shuttle for synthetic, poorly crystalline iron(III) oxide. Finneran et al. (43) observed that addition of Fe(II) to a reduced Geobacter cell suspension stimulated the reoxidation of biologically reduced U; they speculated that it might be due to biological Fe(II) oxidation coupled to nitrate reduction, followed by chemical oxidation of U(IV) by the freshly precipitated iron(III) (hydr)oxide. A similar effect has been reported for Mn(IV)-mediated U(IV) oxidation during Mn(IV) reduction by S. putrefaciens strain CN32 (44). Similar mechanisms may be at work under sulfate-reducing conditions, where biologically reduced U may act as an electron shuttle for the unreacted Fe(III) in mineral treatments (Figures 1-3), with the amount of U(IV) reoxidation depending on both the mass and type of iron(III) (hydr)oxides present. At the end of the experiments, the molar lactate/sulfate utilization ratio averaged 2.11 ( 0.07, in close agreement with the theoretical value of 2 shown in eq 2 (45). Also, the molar lactate consumption to acetate production ratio averaged 0.98 ( 0.03, in excellent agreement with the theoretical value of 1 from eq 2. These results clearly suggest that sulfate reduction proceeded according to eq 2 in the presence of iron(III) (hydr)oxides and that most of the Fe(III) was reduced indirectly by hydrogen sulfide. Although SRB are known to enzymatically reduce Fe(III) (10, 46, 47), such activity appeared to be low in these tests. Figures 1B and 3B indicate that, as expected on the basis of crystallinity, the reactivity of sulfide was initially greater (less sulfide accumulation) with ferrihydrite as compared to the more crystalline structures of hematite and goethite. With an equal amount of Fe(III) present, sulfide was eventually depleted in the goethite and hematite experiments; however, with ferrihydrite, a significant amount of soluble sulfide 2064

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remained in solution (Figures 1B and 3B). These results suggest that the stoichiometry of the secondary iron sulfides (and perhaps polysulfides) that formed depended on the iron(III) (hydr)oxide present. The reaction stoichiometry of Fe and biogenic sulfide is complex and rarely homogeneous. Neal et al. (48), with hematite and D. desulfuricans Essex 6, observed a variety of intermediate S species including S2-, S22-, and Sn2-. Similar S species may also have formed in our experiments as was seen from the variable stoichiometry of Fe and S observed in precipitates (data not shown). The reduction of U by dissolved S(-II) species such as HS- and H2S does not proceed readily in a high-strength bicarbonate buffer (49, 50). However, reactions of U(VI) with reduced-S minerals such as pyrite (FeS2) and galena (PbS) or with Fe(II) sorbed to iron(III) (hydr)oxides and green rusts (mixed ferrous/ferric hydroxide) can result in U reduction and the formation of mixed-valence U oxide compounds (13, 23, 51-53). In contrast, our results suggested that the reduced-S minerals (likely different forms of iron sulfides) associated with SRB activity did not prevent the reoxidation of U(IV) by Fe(III). Finally, the reoxidation of U(IV) nanoparticles by Fe(III) was further studied with FeCl3 and hematite (Figures 5 and 6). Under abiotic conditions, FeCl3 reoxidized U(IV) within 24 h; however, hematite did not. In contrast to the abiotic system, hematite reoxidized U(IV) when sulfate-reducing activity was present. These findings suggest that sulfatereducing activity may make Fe(III) available from the mineral surface to react with microbially reduced U(IV). At this time it is unclear whether nanocrystalline UO2(s) leaves the cell to reduce Fe(III) at the mineral surface or if an electronshuttle compound is involved. Further research is necessary to clarify these reaction mechanisms. The results presented have strong implications for field application of in situ biological reduction of U(VI). After the depletion of an electron-donating substrate, reactions that make Fe(II) unavailable for iron(III) (hydr)oxide surface passivation or the reaction shown in eq 2 may lead to U(IV) reoxidation. For example, acetate was injected in a uraniumcontaminated aquifer located in Rifle, Colorado, to stimulate in situ activity of mixed DMRB. After U concentrations reached a minimum, an increase in U concentration in the groundwater was observed, which was highly correlated with measured SRB activity (8). On the basis of our results, we speculate it may be possible that sulfate-reducing activity made Fe(III) from soil minerals more available to reoxidize microbially reduced U. Our results suggest that, even in an anoxic aquifer, reduction of all available Fe(III) will likely improve the long-term stability of biogenically immobilized U(IV). Implications are that, with biological immobilization of U, care should be taken to maintain the organic substrate in excess throughout the system, such that the aquifer is electron-acceptor-limited, rather than electron-donorlimited, at least until available Fe(III) is depleted from the system. Failing to maintain sufficient substrate concentrations until available Fe(III) is reduced could lead to unfavorable consequences for the long-term stability of immobilized uranium.

Acknowledgments We gratefully acknowledge the financial support provided by the Natural and Accelerated Bioremediation Research program (NABIR) within the Office of Biological and Environmental Research (OBER), U.S. Department of Energy (DOE) (Grants DE-FG03-98ER62630/A001 and DE-FG0301ER63270). The support of the Center for Multiphase Environmental Research and the Department of Chemical Engineering also contributed significantly to this research. The HR-TEM work was performed in the Environmental Molecular Sciences Laboratory, a national scientific user

facility sponsored by DOE/OBER and located at the Pacific Northwest National Laboratory (PNNL). The PNNL is operated for DOE by Battelle Memorial Institute under Contract DE-AC06-76RL0 1830. Use of the Advanced Photon Source was supported by DOE’s Office of Science, Office of Basic Energy Sciences, under Contract W-31-109-Eng-38.

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Received for review April 15, 2004. Revised manuscript received January 13, 2005. Accepted January 14, 2005. ES0494297