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Shewanella alga strain BrY and iron (hydr)oxides of varying stabilities results in complete reduction to Cr(III). The maximum sustainable Cr(VI) reduc...
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Environ. Sci. Technol. 2001, 35, 522-527

Iron Promoted Reduction of Chromate by Dissimilatory Iron-Reducing Bacteria BRUCE WIELINGA,† MIDORI M. MIZUBA,‡ COLLEEN M. HANSEL,† AND S C O T T F E N D O R F * ,† Department of Geological and Environmental Sciences, Stanford University, Stanford, California 94305-2115, and Soil Science Division, University of Idaho, Moscow, Idaho 83844-2339

Chromate is a priority pollutant within the U.S. and many other countries, the hazard of which can be mitigated by reduction to the trivalent form of chromium. Here we elucidate the reduction of Cr(VI) to Cr(III) via a closely coupled, bioticabiotic reductive pathway under iron-reducing conditions. Injection of chromate into stirred-flow reactors containing Shewanella alga strain BrY and iron (hydr)oxides of varying stabilities results in complete reduction to Cr(III). The maximum sustainable Cr(VI) reduction rate was 5.5 µg CrVI‚mg-cell-1‚h-1 within ferric (hydr)oxide suspensions (surface area 10 m2). In iron limited systems (having HEPES as a buffer), iron was cycled suggesting it acts in a catalytictype manner for the bacterial reduction of Cr(VI). BrY also reduced Cr(VI) directly; however, the rate of direct (enzymatic) reduction is considerably slower than by Fe(II)(aq) and is inhibited within 20 h due to chromate toxicity. Thus, dissimilatory iron reduction may provide a primary pathway for the sequestration and detoxification of chromate in anaerobic soils and water.

Introduction Chromium contamination of soil and groundwater is a significant problem worldwide. The extensive distribution of this pollutant is due primarily to its use in numerous industrial processes such as metal plating and alloying, leather tanning, and wood treatment. Additionally, ultramafic rocks such as serpentinite are often enriched in chromium and can lead to locally elevated levels in adjacent soils and waters (1-3). In the United States chromium is the third most common pollutant at hazardous waste sites and is the second most common inorganic contaminant, after lead (4). Once released to the environment, chromium is susceptible to oxidation-reduction reactions that dramatically alter its physical and chemical properties. Chromium persists in the environment in two stable oxidation states, Cr(VI) and Cr(III), which have widely contrasting toxicity and transport characteristics. Hexavalent chromium, Cr(VI), partitions weakly onto solids in soils and waters and consequently tends to be mobile in the environment. Additionally, Cr(VI) is acutely toxic (5, 6), subject to biological uptake (7), and is teratogenic and carcinogenic (8, 9). In contrast, trivalent chromium, Cr(III), has a limited hydroxide solubility and * Corresponding author phone: (650)723-5238; fax: (650)725-0979; e-mail: [email protected]. † Stanford University. ‡ University of Idaho. 522

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FIGURE 1. Rate comparison for chromate reduction by chemical and biological constituents. For calculated initial rates, initial [Cr(VI)] ) 500 µM, [Fe(II)] ) 30 µM, [S(-II)] ) 10 µM. Expressions for ferrous iron and sulfide reaction with chromate are from Pettine et al. (31, 32); biological reduction rate is taken from Lovley and Phillips (33). Reaction conditions were based on (i) chromate concentration used by Lovley and Phillips (33) and (ii) common concentrations of dissolved ferrous iron and sulfide observed in natural environments. Biological reduction rates were normalized for a cell density equal to 1 × 109 cells/mL. forms strong complexes with soil minerals making it relatively immobile and less available for biological uptake (10, 11). Furthermore, Cr(III) is essential (at low concentrations) in human nutrition and is only slightly toxic to plants at very high concentrations. The ability of chromium to undergo electron transfer reactions make it particularly problematic with respect to establishing cleanup standards for protecting public health and for remediation purposes (12). Therefore, understanding processes that promote the reduction of Cr(VI) to Cr(III) with the subsequent detoxification and immobilization is of considerable importance. Chromium(VI) can be reduced by biological and chemical means. It is difficult to determine whether reduction of metal contaminants is mediated principally by direct, enzymatic reduction or by chemical reduction. However, it is likely that the specific pathway by which reduction takes place will be defined by the operating reaction kinetics. Published rate expressions indicate that, of the possible chemical reductants in natural environments, ferrous iron and dissolved sulfide will dominate the reduction of chromate (13). By comparing reduction rates resulting from Fe(II) and S(-II) with those reported for direct microbial reduction, we can postulate operating reaction pathways. The normalization of microbial reduction rates determined under different conditions and comparison with chemical reaction rates from laboratory derived rate expressions is admittedly tenuous. Nevertheless, by using D. vulgaris as a model chromate reducer, one notes that the chemical reduction of chromate by ferrous iron is more than 100× faster than the observed biological rate (Figure 1). Thus, it seems likely that chemical reduction will be the major avenue by which chromate is reduced when either ferrous iron or sulfide are present (i.e., anaerobic environments). Generation of reductant pools (e.g., Fe(II) or sulfide), however, depends on microbial activity, because Fe(III) and SO42- reduction occurs primarily via dissimilatory reduction pathways (14, 15). A diverse and widely distributed group of bacteria are able to couple the oxidation of organic com10.1021/es001457b CCC: $20.00

 2001 American Chemical Society Published on Web 12/23/2000

pounds or H2 to the reduction of iron (hydr)oxides (16, 17). The importance of dissimilatory iron-reduction in mediating numerous biogeochemical processes has been the subject of several recent reviews (14, 16, 18, 19). Thus, in many environments where iron reduction is the predominant terminal electron accepting process (TEAP) in microbial respiration, the indirect reduction of chromate through reaction with a respiratory byproduct is likely a dominant reductive pathway as illustrated by reactions 1 and 2. 3

/4C3H5O3- + 3Fe(OH)3 w 3/4C2H3O2- + 3Fe2+ + 3

/4HCO3- + 2H2O + 51/4OH- (1)

3Fe2+ + HCrO4- + 8H2O w 3Fe(OH)3 + Cr(OH)3 + 5H+ (2) The reduction pathway described by reactions 1 and 2 has interesting implications. Ferrous iron produced in reaction 1 is cycled back to Fe(III) in reaction 2 thereby acting as an electron shuttle (a catalytic role) between the bacteria and chromium. This reaction scheme suggests a continual regeneration of the primary terminal electron acceptor. Thus, by cycling minor amounts of iron present in the environment a significant amount of Cr(VI) could potentially be reduced even in systems having limited available Fe. With the rapid cycling of Fe(II) back to Fe(III), evidence for the role of microbial dissimilatory iron reduction in this process, such as the accumulation of reduced iron solids (magnetite, siderite etc.) and high pore water Fe(II) concentrations, would likely be masked. Here we demonstrate the reduction of chromate via a two-step, closely coupled, biotic-abiotic reaction pathway and the cycling of iron during this process.

Experimental Methods We examined microbially induced, ferrous iron mediated reduction of Cr(VI) using Shewanella alga strain BrY (ATCC No. 51181, hereafter referred to as BrY) as a model dissimilatory iron-reducing bacterium (DIRB). BrY is a wellcharacterized, facultative anaerobic bacterium with the demonstrated ability to couple the oxidation of organic acids and H2 to the reduction of Fe(III), Mn(IV), and U(VI) under anoxic conditions (20). BrY was grown to late log phase in Tryptic Soy Broth (TSB, DIFCO, Detroit, MI) at 32 °C and frozen back in 20% glycerol at -80 °C. Seed cultures were started from frozen stocks (1 mL in 100 mL TSB) and grown for 8 h at 32 °C (150 rpm). Cell suspensions were prepared by adding 1 mL of the seed culture to 100 mL of TSB and growing to late log phase (12 h, 32 °C, 150 rpm). Cells were harvested by centrifugation (4500×g, 10 min, 5 °C), washed twice in 100 mL of anaerobic HEPES (30 mM) or bicarbonate buffer (2.5 g NaHCO3 and 2.5 g NaCl per liter, pH 7.0), and resuspended in the same buffer. Cell suspensions were transferred to sterile, anaerobic pressure tubes having a headspace of N2 gas, capped with a thick butyl rubber stopper, and stored on ice for less than 15 min before inoculation of flow-cells. Heat-killed cells were prepared by holding the cell suspension at 80 °C for 20 min. Iron/chromium reduction reactions were performed in polycarbonate stirred-flow reactors (Figure 2) containing 100 mL of either 30 mM HEPES or 10 mM sodium bicarbonate buffered medium which also contained the following ingredients (in g/L): NaCl, 2.5; NH4Cl, 1.5; KCl, 0.1; CaCl2, 0.1; MgSO4, 0.1 and 10 mL each vitamin and mineral solutions (21). Lactate (as sodium lacate) was added as electron donor at a final concentration of 10 mM unless otherwise stated. In experiments testing the ability of BrY to directly reduce Cr(VI) and gain energy for growth in this process, media was amended with 0.6 g/L KH2PO4.

FIGURE 2. Schematic representation of stirred-flow reactor and experimental setup. Hematite (Fe2O3; 8.6 m2/g) and goethite (R-FeOOH; 15 m2/g) were purchased from Strem Chemicals (Newburyport, MA). Ferrihydrite (nominally 5Fe2O3‚9H2O; 300 m2/g) and hydrous ferric oxide (HFO, ∼600 m2) were synthesized by the titration of Fe(NO3)3‚9H2O (0.4 M) with 1 M NaOH to pH 7 as described by Ryden et al. (22). The resulting ferrihydrite was washed with ultrapure water then freeze-dried and ground to a fine powder. Surface areas were calculated using BET isotherms. HFO prepared by this procedure was not freeze-dried and was maintained as an aqueous suspension; surface area was assumed to be ca. 600 m2/g, equivalent to that described by Roden and Zachara (23). As will be shown, reaction rates were independent of surface area, and thus our assumption as to the surface area of HFO is of little consequence (i.e., surface was in excess). Unless otherwise noted, the iron mineral was added to provide an equal surface area of approximately 10 m2. Standard methods for culture of anaerobic bacteria and preparation of anoxic media were used throughout. Media and buffers were made anoxic by boiling and cooling under a stream of O2-free N2 or N2:CO2 (80:20) gas. All reactions were performed in an anaerobic chamber (Coy Laboratories, Inc., Grass Lake, MI) with an N2(90%):H2(10%) atmosphere. Reactions were initiated by adding cell suspension (final cell concentrated ∼ 2 × 107 mL-1) and iron (hydr)oxide to each reactor and bringing the volume to approximately 100 mL with Cr-free medium. A 0.2 µm pore-size filter and filterbacking disk were placed at the top of the chamber, and the reactor was sealed. The filter served to retain the cells and the iron solids in the reactor. Media was pumped into the bottom of the reactor, and effluent was forced out through the filter to the fraction collector (Figure 2). Initial flow rate before Cr addition varied between 10 and 11 mL/h. Following chromium addition, the pump speed was adjusted downward until a steady-state Fe(II) concentration was maintained in the effluent. The flow-rate varied from about 1.5 mL/h to 4.5 mL/h at steady-state conditions depending on both iron mineral and buffer conditions, yielding retention times that ranged form 22 to 67 h. Samples collected were analyzed for soluble iron and chromium as described below. Direct reduction of chromium by BrY was tested in both batch and stirred-flow reactions. Short-term reduction studies were conducted by adding BrY cell suspension (∼2 × 108 cells/mL final concentrated), prepared as described above, to anaerobic pressure tubes containing 10 mL of bicarbonate buffered medium and chromate added to an initial concentration of 0.075 mM. Long-term, direct reduction studies were done under flow conditions in which iron (hydr)oxide was eliminated, and cells were added directly to medium containing either chromate (0.1 mM) or soluble iron pyrophosphate (10 mM, as a positive control). Solution-Phase Analysis. Production of soluble Fe(II) was monitored spectrophotometrically at 562 nm using the ferrozine assay (24). Dissolved Cr(VI) was measured spectrophotometrically at 540 nm by the s-diphenyl carbazide method (25). Total dissolved chromium and iron were determined by flame atomic absorption spectroscopy (AAS) VOL. 35, NO. 3, 2001 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 3. Temporal changes in Fe(II)(aq) concentration detected in flow-reactor effluent for different iron minerals: (A) goethite; (B) hematite; (C) ferrihydrite; (D) hydrous ferric oxide. Conditions: 30 mM HEPES buffered medium; pH 7; 10 mM lactate added as elect. or by inductively coupled plasma (ICP) optical emission spectroscopy. Solid-Phase Analysis. At the conclusion of each experiment, solids were collected for Cr and Fe speciation and quantification. Total Cr and Fe in the solid-phase were determined by microwave assisted total acid digestion and ICP analysis. A 5 mL aliquot of well-mixed suspension was combined with 4 mL of concentrated, trace metal-grade nitric acid and 1 mL of concentrated hydrofluoric acid and digested following EPA Method No. 3052. The oxidation state of Cr in the solid-phase was determined using X-ray absorption near-edge structure (XANES) spectroscopy as described previously (26), which was conducted at the Stanford Synchrotron Radiation Laboratory (SSRL).

Results and Discussion The ability of BrY to reduce Cr(VI) via Fe(II) produced during iron respiration was tested using four environmentally relevant ferric (hydr)oxides. Soluble Fe and Cr were monitored and speciated throughout the reaction. Prior to the introduction of Cr(VI) to the reactor (0 to 4 h), iron reduction rates specific for each system were calculated from the increase in effluent Fe(II) over time (Figure 3). Chromate was then added to the media reservoirs and pumped continuously into the reaction cell at a rate equal to onethird the calculated rate of Fe(II) production. When chromate amended media enters the reactors, the concentration of Fe(II)(aq) rapidly decreases (Figure 3A-D). In a control to which Cr(VI) is not added to the medium, dissolved Fe(II) concentrations also decline briefly and then continue to increase, at a reduced rate, before reaching a plateau at 40 h (Figure 3A). While the decrease in aqueous Fe(II) in the No-Cr control appears to coincide with the addition of chromate to the other systems, it actually occurs about 2 h after Cr addition and corresponds to a reduction in pumping rate. The decrease in the rate of iron accumulation is expected, due to poisoning of the hydroxide surface area as Fe(II) is adsorbed. Chromium(VI) is not detected in the effluent of 524

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systems having active iron reduction. Inlet Cr(VI) concentration and delivery rates were adjusted until the concentration of Fe(II)(aq) in the effluent was approximately constant over time (i.e., until a steady-state of dissolved Fe(II) is reached). This phase of the reaction is represented by the horizontal regression lines in Figure 3A-D and was maintained for a period of time that ranged between 40 and 100 h. We assume that the steady-state results from Fe(II) produced by BrY at a rate equal to which it is being reoxidized by Cr(VI), as described in reactions 1 and 2. There are two factors that complicate this assumption. First, it is difficult to monitor the concentration of Fe(II) and Cr(VI) present in the solid-phase during the reaction. However, the concentration of Fe(II)(aq) and 0.5 N HCl extractable Fe(II) at the end of the experiment were equivalent, suggesting that Fe(II) was not appreciably sequestered as a surface phase. Additionally, neither Fe(II) nor Cr(VI) were detected in the solid-phase using XANES spectroscopy (data not shown). A second complication is the possibility that BrY could reduce chromate directly. In batch reactions, when BrY is added to bicarbonate buffered medium (ca. 1 × 108 cells/mL) containing 0.075 mM chromate, Cr(VI) is removed from solution in approximately 0.6 h. When the systems are spiked with additional chromate (0.10 mM), Cr(VI) is again removed from solution in ca. 1.5 h. Initial Cr(VI) reduction rates calculated under these conditions averaged 1 × 10-4 M/h, in close agreement with values previously published for bacterial chromate reduction (13), and considerably slower than those observed for Cr(VI) reduction by Fe(II) (27). The slower reaction kinetics observed for direct Cr(VI) reduction by BrY in conjunction with the dramatic decrease in effluent Fe(II) strongly suggests that Fe(II) is the primary reductant under these conditions. The ability of BrY to continue reducing Cr(VI) directly for prolonged periods was also evaluated in stirred reactors. When BrY cell suspension is added to a reactor with an initial Cr(VI) concentration of 0.1 mM, chromate is completely reduced within the first 30 min (Figure 4). There is no

FIGURE 4. Soluble Cr(VI) in reactor effluent versus time. Systems were run without iron mineral to test the direct reduction of Cr(VI) by BrY. Initial Cr(VI) concentration in reactors was 0.1 mM, and the concentration of Cr(VI) in the reservoir was 0.2 mM.

TABLE 1. Averaged Cr(VI) Inlet Concentrations and Steady-State Cr(VI) Reduction Rates for Each Iron Mineral Testeda

iron mineral

Cr(VI) inlet concn (mg/L)

max. sustainable reduction rate (µg CrVI‚mg-cell-1‚h-1)

hematite goethite ferrihydrite hydrous ferric oxide

10.1 ((1.8) 12.7 ((7.3) 9.8 ((2.1) 12.0 ((1.2)

2.7 ((0.3) 3.4 ((1.9) 1.6 ((0.1) 1.8 ((0.2)

a Conditions: 30 mM HEPES buffered medium, pH 7.0 with 10 mM sodium lactate as electron donor. Values represent the mean (( SD); for goethite (n)3), for all other minerals (n)2).

detectable change in soluble Cr(VI) when heat-killed BrY cells are added. Injecting 0.2 mM Cr(VI) leads to an immediate increase in Cr(VI) concentration within the effluent of sterile reactors. In contrast, Cr(VI) is completely removed from solution in live reactors during the first 20 h of the experiment; after this initial reaction period, however, effluent Cr(VI) increases. The rate of increase in the live system after 20 h is comparable to that observed in sterile reaction, indicating that direct reduction of chromate by BrY has ceased. Thus, prolonged reduction of Cr(VI) in the presence of iron is likely the result of Fe(II) catalyzed reduction. An equally important finding of these studies is the difference in cell viability for the chromium versus iron pyrophosphate reactions. At the conclusion of the experiment, aliquots from each system were 10-fold serially diluted and plated on Trypticase Soy Agar medium (incubated aerobically at 25 °C). The iron pyrophosphate amended system contained 5 × 10 7 CFU/ mL, while no viable cells were recovered from chromium amended reactors. Hence, it appears that reduction of iron also provides BrY with an indirect resistance mechanism against Cr(VI) toxicity. Initial iron reduction rates, and consequently the ability to catalyze chromate reduction, are similar among all iron minerals tested. Inlet Cr(VI) concentrations along with maximum sustainable Cr(VI) reduction rates, averaged for each mineral, are listed in Table 1. These findings are consistent with the observations of Roden and Zachara (23) that the rate and extent to which bacteria reduce iron (hydr)oxides is primarily controlled by mineral surface area, which is constant in these reactions. In addition, Fe(III) produced during the reduction of Cr(VI) (reaction 2) will in large part precipitate as an amorphous hydroxide and be amenable to further microbial reduction. Thus, while the ferric iron initially present may be in a crystalline form, the reaction should

lead to the formation of a high surface area, amorphous mineral. It is also possible, and probable, that Fe(III) formed is incorporated into a mixed (CrxFe1-x)(OH)3 solid, as previously observed during reduction of Cr(VI) by amorphous iron sulfide (26). The availability of Fe(III) for microbial reduction in this solid is unknown. However, comparison of solids from experiments in which iron cycles (see discussion below) with solids in which iron is not indicates that further reduction of the mixed hydroxides takes place. Preliminary analysis of X-ray absorption spectra, which will be expanded upon subsequently, suggests that a mixed Cr-Fe hydroxide is formed and that the ratio of Cr:Fe changes with time. Data presented here, however, suggests that even crystalline iron oxides should be potent promoters of Cr(VI) reduction under iron-reducing conditions. When HFO is used as the iron phase (Figure 3D), the total amount of Cr(VI) reduced to Cr(III) is 7.9 × 10-5 ((8.5 × 10-6) mols which would require that 2.4 × 10-4 ((2.8 × 10-5) mols of Fe(II) be oxidized. Total iron initially available for reduction in both systems was 1.6 × 10-4 mols. Based on the total iron added, iron is thus cycled approximately 1.5 times. Mass balance calculations of chromium recovered in the aqueous and solid-phase accounted for between 85 and 90% of the total Cr(VI) added. Aggregation during the experiments made collection of solids difficult owing to an accumulation of solids on the stirring mechanism; thus, it was anticipated that solid phase analysis would underestimate total Cr sequestered in that fraction. If soluble Fe(II) removed from the system in the effluent is considered, iron is cycled approximately twice during this relatively short experiment. Thus, the cycling of iron implicit in reactions 1 and 2 is observed. Because solid-phase iron remaining at the end of the experiment is in the oxidized (ferric) state, it appears that Fe-cycling could continue unabated. In practice this may not hold true since the rate of iron reduction appears to decrease at the end of the reaction. The diminished reduction rate may result from a decrease in total Fe via the loss of Fe(II)(aq) from the flow cell, passivation of the ferric iron surface, or decreased reactivity of a mixed (Cr,Fe)(OH)3 phase. It is also possible that because these experiments were conducted under nongrowth conditions (e.g. no phosphate), the viable population of cells may have declined. Canfield et al. (28) demonstrate the importance of Fe redox cycling for organic matter mineralization in coastal sediments showing that, on average, each Fe atom is oxidized and reduced between 100 and 300 times before being buried in the sediment, indicating that Fe cycling similar to that demonstrated in our laboratory experiments also occurs in natural environments. The ability of iron to be cycled and act as a catalyst for Cr(VI) reduction should be considered when accessing the capacity of soils or waters to effectively reduce Cr(VI) to Cr(III). While Cr(VI) is completely reduced to Cr(III) in active iron-reducing reactors, its precipitation is limited in these systems (Figure 5). HEPES buffer (at 30 mM) and lactate (at 10 mM) prevent the precipitation of Cr(III) at pH 7, presumably through complexation. Total chromium in the effluent of HFO systems, expressed as a percentage of the influent concentration (Figure 5), is 85%, 70%, and 96% for the two viable reactors, and the abiotic control, respectively. However, total chromium and Cr(VI) are equivalent in the sterile control, whereas no Cr(VI) is measured in the effluent of live reactions. An average of 46% of the total chromium entering iron-reducing cells is sequestered versus 26% in the control. Difficulty in removing Cr(III) from organic-rich, microbiological media has been previously encountered (29). To determine if Cr(III) would likely precipitate and be sequestered under conditions similar to those encountered in most natural environments, reactions with HFO were performed in 20 mM bicarbonate buffered medium. Here, VOL. 35, NO. 3, 2001 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 5. Difference between inlet Cr(VI) concentration (C0) and total dissolved chromium in reactor effluent (CT) as a function of time for systems containing HFO. Conditions: 30 mM HEPES buffered medium; pH 7; 10 mM lactate as electron donor.

chromium. In bicarbonate-buffered reactions, no chromium is detected in the effluent of active iron-reducing reactors (Figure 6B). Thus, all of the chromium entering these systems is retained in the solid-phase compared to only about 45% in the HEPES-buffered reactions. Furthermore, XANES analysis verified that all the chromium present was in the reduced, Cr(III) state. It is possible that precipitation of siderite occurs and may sequester Cr(III) in the bicarbonate buffered systems. However, Cr(III) has a low affinity for carbonate phases, and our preliminary structural analysis indicates the formation of Cr-Fe hydroxides. As in the HEPES-buffered reactions, effluent Cr(VI) and Crtotal in the abiotic controls was equivalent and approached the inlet Cr concentration after ca. 80 h (Figure 6B). Urrutia et al. (30) have recently shown that the extent of iron reduction is enhanced by the presence of Fe(II) complexants such as NTA. Their work supports the contention that compounds which retard the sorption of Fe(II) on oxide and/or cell surfaces promotes and extends iron reduction. The increased solubility of Cr and Fe in the presence of HEPES presumably had the same affect, delaying the sorption of Cr(III) and Fe(II) to the iron hydroxide and/or cell surface thereby extending iron reduction. In summary, the microbial reduction of iron (hydr)oxides, which varied over a wide range of stabilities, promotes reduction of Cr(VI) to Cr(III). Reduction and immobilization of chromate is the result of a coupled, two-step, bioticabiotic reaction pathway in which Fe(II) produced during iron respiration catalyzes the reduction of Cr(VI). In the presence of HEPES buffer, iron is cycled almost 2 times indicating that it is essentially acting as an electron shuttle between the bacterium and chromium. In the absence of organic buffers, reduction also resulted in the total sequestration of Cr(III) in the solid-phase. Thus, attenuation of chromate in saturated, subsurface environments may be in large part attributable to dissimilatory iron reduction. In addition, the capacity for soils to reduce and immobilize Cr(VI) could be dramatically underestimated if this bioticabiotic process is not appreciated. It is therefore feasible that the stimulation of DIRB may provide a useful tool for the in situ stabilization and detoxification of chromium.

Acknowledgments This work was supported by the Department of Energy’s NABIR program (grant number DE-FG03-97ER62481). We thank Dr. Guangchao Li and Benjamin Bostick for their technical support with solution and solid analysis and three anonymous reviewers for thoughtful comments.

Literature Cited FIGURE 6. Temporal changes in (A) Fe(II)(aq) and (B) Cr(VI)(aq) resulting from reactors containing HFO and 20 mM sodium bicarbonate buffered medium. Lactate is present as the electron donor at an initial concentration of 5 mM. a lactate concentration of 5 mM and HFO surface areas of 10 m2 and 20 m2 were used. Differences between the HEPES and bicarbonate buffered systems are noted with respect to iron reduction (Figure 6A and Figure 3D). In HEPES buffered systems, soluble Fe(II) levels remained stable for approximately 60 h, and levels drop only slightly after 168 h (Figure 3D). In contrast, Fe(II) levels in the bicarbonate buffered systems stay constant for only about 20 and 40 h in systems containing 10 m2 and 20 m2 HFO, respectively (Figure 6A). Additionally, Fe(II)(aq) levels decline rapidly after about 50 h of operation. The most significant difference between the two systems, however, is the sequestration of 526

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(1) )Nriagu, J. O. In Chromium in the Natural and Human Environments; Nriagu, J. O., Nieboer, E., Eds.; Wiley: New York, 1988; Vol. 20, p 81. (2) Robertson, F. N. Ground Water 1975, 13, 516. (3) Armienta, M. A.; Rodriguez, R.; Quere, A.; Juarez, F.; Ceniceros, N.; Aguayo, A. Int. J. Environ. Anal. Chem. 1993, 54, 1. (4) National Research Council. Environmental Epidemiology, Vol. 1, Public Health and Hazardous Wastes; National Academy Press: Washington, DC, 1991; p 108. (5) Babich, H.; Schiffenbauer, M.; Stotzky, G. Bull. Environ. Contam. Toxicol. 1982, 28, 452. (6) Beyersmann, D.; Koster, A.; Buttner, B.; Flessel, P. Toxicol. Environ. Chem. 1984, 8, 279. (7) Turner, M. A.; Rust, R. H. Soil Sci. Soc. Am. Proc. 1971, 35, 755. (8) Abbasi, S. A.; Soni, R. Int. J. Environ. Stud. 1984, 23, 131. (9) Yassi, A.; Nieboer, E. In Chromium in the Natural and Human Environments; Nriagu, J. O., Nieboer, E., Eds.; Wiley and Sons: New York, 1988; Vol. 20, p 443. (10) Sass, B. M.; Rai, D. Inorg. Chem. 1987, 26, 2228. (11) Fendorf, S. E.; Fendorf, M.; Sparks, D. L.; Gronsky, R. J. Colloid Interface Sci. 1992, 153, 37.

(12) James, B. R.; Petura, J. C.; Vitale, R. J.; Mussoline, G. R. J. Soil Contam. 1997, 6, 569. (13) Fendorf, S.; Wielinga, B. W.; Hansel, C. M. Int. Geol. 2000, 42, 691. (14) Lovley, D. R. Microbiol. Rev. 1991, 55, 259. (15) Lovley, D. R.; Phillips, E. J. P.; Lonergan, D. J. Environ. Sci. Technol. 1991, 25, 1062. (16) Lovley, D. R. Annu. Rev. Microbiol. 1993, 47, 263. (17) Coates, J. D.; Phillips, E. J. P.; Lonergan, D. J.; Jenter, H.; Lovley, D. R. Appl. Environ. Microbiol. 1996, 62, 1531. (18) Nealson, K. H.; Saffarini, D. Annu. Rev. Microbiol. 1994, 48, 311. (19) Fredrickson, J. K.; Gorby, Y. A. Curr. Opin. Biotechnol. 1996, 7, 287. (20) Caccavo, F. J.; Blakemore, R. P.; Lovley, D. R. Appl. Environ. Microbiol. 1992, 58, 3211. (21) Balch, W. E.; Fox, G. E.; Magrum, L. J.; Woese, C. R.; Wolfe, R. S. Microbiol. Rev. 1979, 43, 260. (22) Ryden, J. C.; Syers, J. K.; Tillman, R. W. J. Soil Sci. 1987, 38, 211. (23) Roden, E. E.; Zachara, J. M. Environ. Sci. Technol. 1996, 30, 1618. (24) Stookey, L. L. Anal. Chem 1970, 42, 779. (25) Bartlett, R.; James, B. J. Environ. Qual. 1979, 8, 31.

(26) Patterson, R. R.; Fendorf, S.; Fendorf, M. Environ. Sci. Technol. 1997, 31, 2039. (27) Pettine, M.; L., D. O.; Campanella, L.; Millero, F. J.; Passino, R. Geochim. Cosmochim. Acta 1998, 62, 1509. (28) Canfield, D. E.; Thamdrup, B.; Hansen, J. W. Geochim. Cosmochim. Acta 1993, 57, 3867. (29) Komori, K.; Rivas, A.; Toda, K.; Ohtake, H. Appl. Microbiol. Biotechnol. 1990, 33, 117. (30) Urrutia, M. M.; Roden, E. E.; Zachara, J. M. Environ. Sci. Technol. 1999, 33, 4022. (31) Pettine, M.; Millero, F. J.; Passino, R. Mar. Chem. 1994, 46, 335. (32) Pettine, M.; Barra, I.; Campanella, L.; Millero, F. J. Water Res. 1998, 32, 2807. (33) Lovley, D. R.; Phillips, E. J. P. Appl. Environ. Microbiol. 1994, 60, 726.

Received for review July 6, 2000. Revised manuscript received October 25, 2000. Accepted November 8, 2000. ES001457B

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