Kinetics of Uranium(VI) Reduction by Hydrogen ... - ACS Publications

Jun 24, 2006 - University of Missouri-Columbia, Columbia, Missouri 65211,. Department of ... University of Wisconsin, Madison, Wisconsin 53706, and...
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Environ. Sci. Technol. 2006, 40, 4666-4671

Kinetics of Uranium(VI) Reduction by Hydrogen Sulfide in Anoxic Aqueous Systems BIN HUA,† HUIFANG XU,‡ J E F F T E R R Y , § A N D B A O L I N D E N G * ,† Department of Civil and Environmental Engineering, University of Missouri-Columbia, Columbia, Missouri 65211, Department of Geology and Geophysics, University of Wisconsin, Madison, Wisconsin 53706, and Department of Biological, Chemical, and Physical Sciences, Illinois Institute of Technology, Chicago, Illinois 60616

Aqueous U(VI) reduction by hydrogen sulfide was investigated by batch experiments and speciation modeling; product analysis by transmission electron microscopy (TEM) was also performed. The molar ratio of U(VI) reduced to sulfide consumed, and the TEM result suggested that the reaction stoichiometry could be best represented by UO22+ + HS- ) UO2 + S° + H+. At pH 6.89 and total carbonate concentration ([CO32-]T) of 4.0 mM, the reaction took place according to the following kinetics: -d[U(VI)]/dt ) 0.0103[U(VI)][S2-]T0.54 where [U(VI)] is the concentration of hexavalent uranium, and [S2-]T is the total concentration of sulfide. The kinetics of U(VI) reduction was found to be largely controlled by [CO32-]T (examined from 0.0 to 30.0 mM) and pH (examined from 6.37 to 9.06). The reduction was almost completely inhibited with the following [CO32-]T and pH combinations: [(g15.0 mM, pH 6.89); (g4.0 mM, pH 8.01); and (g2.0 mM, pH 9.06) ]. By comparing the experimental results with the calculated speciation of U(VI), it was found that there was a strong correlation between the measured initial reaction rates and the calculated total concentrations of uranium-hydroxyl species; we, therefore, concluded that uranium-hydroxyl species were the ones being reduced by sulfide, not the dominant U-carbonate species present in many carbonate-containing systems.

Introduction Uranium (U) contamination has been detected at numerous U.S. Department of Energy waste sites (1). It is also found at other locations such as the agriculture evaporation ponds (3 µg to 22 mg/L) of the San Joaquin Valley, California (2) and various U mine tailings (3). In aquatic systems, uranium exists primarily as complexed, sorbed, or precipitated uranyl (UO22+) carbonate and/or hydroxide species (4). The mobility of U in aquatic systems is largely determined by its speciation and interaction with other aquatic constituents, including microorganisms, natural organic matter, inorganic ions, and mineral surfaces. Such interactions may result in U adsorption, redox transformation, and precipitation, all of which are relevant to many geological and environmental processes. * Corresponding author phone: (573) 882-0075; fax: (573) 8824784; e-mail: [email protected]. † University of Missouri, Columbia. ‡ University of Wisconsin, Madison. § Illinois Institute of Technology, Chicago. 4666

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Adsorption of U(VI) onto mineral surfaces has been well explored as one of the potential approaches for U(VI) immobilization (3, 5, 6). Adsorption, however, could be greatly influenced by the capacity of adsorbents, competition from other species for the surface sites, and reaction with U-complexing agents. For instance, U(VI) adsorption by iron oxyhydroxides depended strongly upon carbonate concentration and pH because of their effects on U(VI)-carbonate complex formation (7). Biotic or abiotic reduction of U(VI) to uraninite (UO2), which has an extremely low solubility [Ksp ) 10-60.6 (4)], could be an effective approach for uranium immobilization. Some microorganisms, including Fe(III)-reducing Geobacter metallireducens (8), sulfate-reducing Desulfovibrio vulgaris (9), and sulfate-reducing Desulfovibrio desulfuricans strain G20 (10), are known to reduce soluble U(VI) to insoluble U(IV). In heavily contaminated sites, however, biotic reduction of heavy metals and radionuclides could be limited by the lack of microorganisms and suitable electron donors, by the presence of competitive electron acceptors, by the presence of heavy metals toxic to microorganisms, and/or by unfavorable pH conditions (11, 12). Chemical reduction of U(VI) could be coupled with various reduced iron and sulfur species, including mixed ferrous/ ferric hydroxides (13), Fe(II) sorbed on hematite surfaces (14), zerovalent iron (15), and various dissolved and solid sulfide species (4, 6, 16-18). There have been, however, conflicting reports on U(VI) reduction kinetics and products. Wersin et al. (17) and Livens et al. (18) provided spectroscopic evidence showing the sorption and reduction of U(VI) by sulfide minerals. Laboratory experiments also showed formation of precipitates upon bubbling H2S gas into 5-10 × 10-2 g/L of uranyl solution (free from CO2) within several hours (19), as well as removal of aqueous U(VI) through both enzymatic and chemical reduction that was associated with sulfate-reducing biofilms in the absence of bicarbonate (20). In contrast, uranium in anoxic Black Sea water was present as U(VI), rather than the seemingly favored U(IV) species, even though the sulfide concentration was as high as 400 µM (21). Another perplexing observation is the production of multiple reaction products. For example, uranium in sulfidic sediments from the San Joaquin Valley of California was determined to be pitchblende (U3O8) by X-ray absorption near edge structure spectroscopy (XANES) (22), whereas precipitates produced by aqueous sulfide reduction in the laboratory were identified to be uraninite by X-ray diffraction (19), as well as by XANES and transmission electron microscopy (TEM) (20). While the detailed kinetics and mechanism of U(VI) reduction by sulfides have not been fully elucidated, available literature has, nevertheless, overwhelmingly suggested that reductive U immobilization by sulfide species could take place under reducing environmental conditions. A contaminant immobilization/site remediation technology, in situ gaseous reduction, has been proposed for the remediation of soils contaminated with redox-sensitive metals (e.g., Cr, U, and Tc) by hydrogen sulfide injection, and has been fieldtested for Cr (23). The technology could similarly be applied for reductive U immobilization. A better understanding of the kinetics and mechanism of U(VI) reduction by sulfide is needed to apply such reduction-based technologies. The information could also provide insight into other geological and environmental processes such as U ore formation and its biogeochemical cycling. The objectives of this study were to (i) determine the reaction stoichiometry and the major reaction products of 10.1021/es051804n CCC: $33.50

 2006 American Chemical Society Published on Web 06/24/2006

aqueous U(VI) reduction by H2S; (ii) investigate the kinetics; and (iii) evaluate the effects of carbonate and pH on the reduction process. Particular effort was made to correlate the initial rates of U reduction and its speciation, aiming to identify the dominant U species involved in the redox transformation.

Materials and Methods Chemicals. Uranyl nitrate standard solution (997 ( 2 mg/L), sodium sulfide (Na2S‚xH2O), sodium bicarbonate, and N,Ndimethyl-p-phenylenediamine (DPD) were from Fisher Scientific. 2-(2-thiazolylazo)-p-cresol (TAC), 4-(2-hydroxyethyl)-1-piperazineethanesulfonic acid (HEPES), and piperazine-N,N′-bis(2-ethanesulfonic acid) (PIPES) were from Sigma-Aldrich. Tris(hydroxymethyl)aminomethane (TRIS) was from BioRad Laboratory. All reagents were ACS reagent grade and used as received, except that sodium sulfide crystals were rinsed with distilled and deionized water (18 MΩ‚cm, Millpore, DDW) to remove any oxidized surface layer. Experimental Approaches. Experiments were performed in an anoxic gas chamber (Coy Laboratory Products) with ∼5% H2 and ∼95% N2. To minimize the ambient concentrations of CO2 and H2S in the chamber, 2 L of 2 M NaOH and 2 L of 15 mM Ag2SO4 solutions were stored inside the chamber in open containers for the gas absorption. To eliminate CO2 and O2 in DDW, the pH of the water was first adjusted to ∼3.0 with HCl to expel any dissolved CO2; then brought back to 7.0 ( 0.1 with freshly prepared NaOH solution in the chamber. DDW was open to the anoxic atmosphere for at least 24 h prior to experiments, so any residual oxygen would be removed from solution by outgassing and subsequent reaction with H2 in the glovebox. Stock solutions of NaHCO3 and U(VI) were also stored in the chamber. Preliminary experiments showed that phosphate and HEPES buffers were not suitable for use, as both interfered with the spectrophotometric determination of U(VI). Consequently, PIPES and TRIS were selected, as neither formed complexes with uranyl ion (24), nor showed interference with U(VI) analysis at a level of 12.5 mM. Experimental runs began by transferring specific amounts of pH- buffered Na2S solution and NaHCO3 solution into a set of 20 mL-glass vials; then, U(VI) stock solution was added to initiate the reaction. Controls with only Na2S or U(VI) were also conducted to assess the reactant stability. Total volume of solution was set to be 20 mL so there was no headspace in the experimental vials. For each kinetic series, one vial was sacrificed to measure the initial pH. Other vials were mixed by magnetic stirring at 22 ( 2 °C. At each selected time point, a reaction vial was sacrificially sampled by filtering 1-2 mL of solution through a 0.2 µm nylon filter (Fisher Scientific) for U(VI) analysis. A series of experiments were conducted to investigate the effect of pH on U(VI) reduction by sulfide. Solution pH ranged from 6.37 to 9.06, controlled by PIPES buffer for pH values below 7.5 and by TRIS buffer for higher pH. Total carbonate concentration ([CO32-]T) was varied from 0 to 30.0 mM. Selected experiments were run in duplicate, and the average values were reported as shown in Figures 3 and 5 in the next section. Analytical Methods. U(VI) was determined spectrophotometrically following the procedure prescribed by Teixeira et al. (25). U(VI) ions and TAC formed a blue complex that was stable for at least 3 h. The absorbance of the solution was determined at 588 nm in a 1-cm cell using a spectrophotometer. Beer’s law was obeyed up to 500 µM, with a detection limit of 0.4 µM. Selected samples were also analyzed by KPA-11 (Chemchek Instruments), which is specific for aqueous U(VI) with a detection limit as low as 0.4 × 10-4 µM, based on a time-resolved kinetic phosphorescence measure-

ment. Data obtained by these two methods were consistent, with a difference always less than 8%. Sulfide concentration in the aqueous phase was analyzed spetrophotometrically through its reaction with a DPD reagent, yielding methylene blue (26, 27). The sulfide stock solution was calibrated by the standard iodometric method (27). Transmission Electron Microscopy. Solid reaction products were examined by high-resolution transmission electron microscopy (HRTEM) and X-ray energy-dispersive spectroscopy (EDS) on a Tecnai F30 field emission gun scanning transmission electron microscope (EM Vision 4.0). The precipitates were prepared by mixing 70 mL of 360 µM U(VI) with 70 mL of 2.0 mM sulfide. pH and [CO32-]T for the reaction were set at 6.89 and 4.0 mM, respectively. After 12 h, the supernatant was decanted and the precipitates were rinsed by DDW several times and transferred to a 25 mL amber glass vial. The samples were stored under the anoxic condition for 2 days prior to TEM analysis by depositing a drop of suspension onto holey carbon-coated Cu grids. Drying on the grids took only several minutes in the ambient atmosphere, and past experience showed no U(IV) oxidation following this procedure.

Results and Discussion Control experiments were performed to check the reactant stability. The first one showed that 168.3 µM U(VI) in a system with 4.0 mM total carbonate (pH 6.89) was stable for at least 24 h. The other system containing 2.0 mM sulfide showed a slight decrease in sulfide concentration within 6 h at pH 6.89, but the decrease was always less than 10% in the time frame used for kinetics measurements. Reaction Stoichiometry and Product Identification by TEM. The molar ratio of U(VI) consumed to the amount of sulfide reacted was determined by simultaneously measuring the concentrations of both reactants as a function of time, sacrificing a testing vial at each time point. All vials initially contained 376.0 µM [U(VI)]0, 498.0 µM [S2-]T, 4.0 mM [CO32-]T, as well as 12.5 mM PIPES buffer (pH 6.89). The results indicated that U(VI) consumption was directly proportional to sulfide consumption. The slope of the best fitting line was 0.95 ( 0.11 (95% confidence level, Supporting Information: Figure S1), suggesting that the stoichiometric ratio of U(VI) to sulfide was 1:1, i.e., it took one mole of sulfide to reduce one mole of U(VI). Precipitates formed from the reaction between U(VI) and sulfide were composed of uranium and oxygen (Figure 1a) and had a uniform particle diameter of approximately 3 nm (Figure 1b). The HRTEM image and electron diffraction pattern (Figure 1c) indicated that the uranium product was crystalline uraninite. We also tried to examine the oxidation product of sulfide by TEM but failed to detect a strong signal of sulfur species; it was likely that the washing process implemented to obtain clean TEM samples removed sulfur residuals. Different products were observed from U(VI) reduction by sulfide in various environments. Wersin et al. (17) reported that the products of U(VI) reduction after exposure to galena or pyrite were some U3O8-type compounds, based on Fourier transformed infrared analyses. Moyes et al. (6) and Livens et al. (18) also proposed the formation of a U3O8 phase on mackinawite surfaces. Using XANES, Duff et al. (22) found that uranium in sulfidic sediments was approximately 25% U(IV) and 75% U(VI). Thus, the precipitates could be pitchblende (U3O8) formed according to the following stoichiometry:

12UO2(CO3)34- + HS- + 47H+ ) 4U3O8(s) + SO42- + 36CO2(g) + 24H2O (1) VOL. 40, NO. 15, 2006 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 1. Results from TEM and energy dispersive spectroscopy shows (a) the product is primarily composed of uranium and oxygen (the peaks of Cu and C were from carbon-coated Cu-grids); (b) the average diameter of the precipitate particles is ∼3 nm; and (C) the HRTEM and electron diffraction patterns are consistent with the product being nanocrystalline uraninite. The first three rings of the diffraction pattern are 111, 200, and 220 reflections of uraninite. Thermodynamic analyses (4, 16), however, suggested that the reduction of U(VI) by aqueous sulfide with uraninite as the proposed reaction product was also highly favorable:

4UO2(CO3)34- + HS- + 15H+ ) 4UO2(s) + SO42- + 12CO2(g) + 8H2O (2) Uraninite formation from aqueous U(VI) reduction has, in fact, been observed in laboratory systems with hydrogen sulfide: in a CO2-free system with pH 6.0-6.5 and [U(VI)]0 of 5∼10 × 10-2 g/L, 10-20 nm hollow globules were precipitated from the solution and were identified to be uraninite by X-ray diffraction (19). A recent study by Beyenal et al. (20) similarly found that the reduction product of 140 µM U(VI) by 1.0 mM sulfide at pH 7.0 was uraninite by the U L3-edge XANES spectra and TEM with selected area electron diffraction. Their analysis was conducted under a strictly anoxic condition. Our study using TEM and EDS clearly shows that U(VI) reduction by hydrogen sulfide produces crystalline uraninite (Figure 1), agreeing with the X-ray diffraction analysis by Kochenov et al. (19) and XANES and TEM by Beyenal et al. (20). Since the stoichiometric ratio of U(VI) reduced to sulfide consumed is 1:1, the result implies that the product of sulfide oxidation is elemental sulfur. While elemental sulfur was not directly detected in this study, sulfide oxidation by chromate was found to produce elemental sulfur nanoparticles under comparable experimental conditions (28, 29). A reaction stoichiometry consistent with these experimental observations is as follows:

UO22+ + HS- ) UO2 + S0 + H+

(3)

Reaction Orders. To obtain the reaction orders with respect to both U(VI) and sulfide, a set of tests were conducted with [U(VI)]0 being fixed at ∼168.3 µM, but [S2-]T varied from 0.83 to 6.83 mM, always significantly higher than [U(VI)]0. Solution pH was controlled at 6.89 and [CO32-]T was at 4.0 4668

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FIGURE 2. ln([U(VI)]/[U(VI)]0) versus time plots showing the reaction is pseudo-first order with respect to U(VI) (reaction conditions: [U(VI)]0 ) 168.3 µM; pH 6.89; [CO32-]T ) 4.0 mM; temperature ) 22 °C. R2 represents correlation coefficient). mM. As shown in Figure 2, ln[U(VI)] versus time plots are mostly linear, indicating a pseudo-first order dependence on U(VI). Therefore, at a given pH and [CO32-]T, the reaction kinetics could be expressed as the following:

d[U(VI)] ) k0[U(VI)][S2-]xT dt

-

(4)

where k0 is the rate constant and x the reaction order in sulfide. This kinetic expression, however, should only be considered as the first approximation of the kinetic process, because some deviation of the plots from linearity was visible at low initial U(VI) concentrations. The derivation could result from catalysis of reaction products, similar to the accelerated Cr(VI) reduction by sulfide due to catalysis of elemental sulfur nanoparticles (29), but more detailed studies are needed to fully assess the phenomenon.

FIGURE 3. Percentage U(VI) reduction ([U(VI)]r/[U(VI)]0 % as a function of [CO32-]T, with [U(VI)]r denoting the reduced amount of U(VI) after 1 h of reaction. It decreased dramatically with increasing [CO32-]T and pH. The data at pH 6.89 were the averages of duplicated tests. The curves were not the best fit of the data but the calculated U(VI) speciation as discussed in the text (initial reactant concentrations: [U(VI)]0 ) 168.3 µM; [S2-]T ) 2.0 mM). Since [S2-]T was significantly higher than [U(VI)]0 for all tests and the stoichiometric ratio of the two reactants was close to 1, [S2-]T was near constant during the reaction. As a result, the slopes of the lines in Figure 2 should be proportional to [S2-]T, leading to a linear relation between ln(-slope) and ln[S2-]T. The reaction order with respect to sulfide obtained from this relationship was 0.54 ( 0.27 (at 90% confidence level) and the corresponding rate constant, k0, was 0.0103 ( 0.0032 (at 90% confidence level). The kinetic equation under the experimental conditions, therefore, is as follows:

d[U(VI)] ) 0.0103[U(VI)][S2-]T0.54 dt

-

(5)

Effects of [CO32-]T on U(VI) Speciation and Reactivity. The effect of [CO32-]T on U(VI) reduction was examined in systems with 168.3 µM [U(VI)]0 and 2.0 mM [S2-]T by measuring percentage U(VI) reduction after exactly 1 h of reaction, represented by [U(VI)]r/[U(VI)]0, where [U(VI)]r denotes the reduced amount after 1 h. [CO32-]T was varied from 0.0 to 30.0 mM in three sets of experiments with pH fixed at 6.89, 8.01, or 9.06, respectively. As shown by Figure 3, increasing [CO32-]T dramatically decreased [U(VI)]r/[U(VI)]0 and the effect was more pronounced at higher pH. At pH 6.89, over 96% U(VI) was reduced at 1.0 mM [CO32-]T, 64% U(VI) was reduced at 4.0 mM [CO32-]T, and only 17% U(VI) was reduced at 6.0 mM [CO32-]T. Similarly, at pH 8.01, when [CO32-]T was set at 1.0 mM, 80% U(VI) was removed from solution in 1 h; after [CO32-]T was increased to 4.0 mM, the reduction of U(VI) was minimal. The effect of carbonate was more dramatic at pH 9.06: even when [CO32-]T was as low as 1.0 mM, only 18% U(VI) was precipitated; as [CO32-]T was increased to 2.0 mM, no reaction was observed in 1 h of testing period. The coordination chemistry of U(VI) is known to be complex. Depending on the solution pH and [CO32-]T, there could have been numerous hydroxyl and/or carbonate species present in the aqueous phase. Reactivity of U(VI) toward its reduction by sulfide is expressed here in a generic form based on all existing U(VI) species:

d[U(VI)] + ) (k1[UO2+ 2 ] + k2[UO2OH ] + dt + k3[(UO2)2(OH)2+ 2 ] + k4[(UO2)3(OH)5 ]

-

+ k5[UO2CO3] +

k6[UO2(CO3)22 ]

+

2- x k7[UO2(CO3)43 ])[S ]T (6)

where k1-k4 are the kinetic constants corresponding to various hydroxyl species, and k5-k7 are the constants for various carbonate species. The reaction order is assumed to be one for all U(VI) species and fractional order (x) for total sulfide, similar to what was observed (eq 3). To determine which terms (or U(VI) species) in eq 6 are most important for U(VI) reduction, we have calculated U(VI) species distribution under various conditions comparable to initial reactant concentrations and pH used in our experiments, utilizing a speciation software, MINEQL+ (version 4.07), and associated equilibrium constants (SI: Table S1). The hypothesis is that by comparing the initial rates and U(VI) speciation under a wide variety of pH and concentration conditions, the major terms responsible for U(VI) reduction could be identified. The calculation was performed without taking ionic strength into consideration. The ionic strength in all experiments varies based on the amounts of buffers, carbonate, sulfide, and U(VI) used, but it is always below 0.05 M and should not significantly affect the results. The calculated initial species distribution of U(VI) in a system with [U(VI)]0 ) 168.3 µM under various carbonate concentrations is presented in Figure 4. Solution pH is fixed at 6.89. The results show that as [CO32-]T is increased from 4.0 mM to 15.0 mM, the calculated [UO2(CO3)34-]/[U(VI)]0 and [UO2(CO3)22-]/[U(VI)]0 would increase from 2.2 to 8.7% and 87.9 to 88.8%, respectively (Figure 4a). Therefore, should these two U(VI)-carbonate complexes be the major reacting species, the rate of U(VI) reduction by aqueous sulfide would increase or at least remain constant with increasing [CO32-]T. The experimental results, however, indicated that the values of [U(VI)]r/[U(VI)]0 after 1 h reaction with sulfide decreased from 64 to 8% with the corresponding increase of [CO32-]T (Figure 3). Thus, both UO2(CO3)22- and UO2(CO3)34- could be disregarded as the major species for U(VI) reduction. Another U(VI)-carbonate species, UO2CO3(aq), is also unlikely to be the major species by similar reasoning: when [CO32-]T is increased from 1.0 to 4.0 mM, the calculated [UO2CO3(aq)]/ [U(VI)]0 would increase from 7.7 to 9.8% (Figure 5a), but the observed ratio of [U(VI)]r/[U(VI)]0 after 1 h reaction decreased from 96 to 64% with the same magnitude of [CO32-]T increase (Figure 3). On the other hand, increasing [CO32-]T from 1.0 mM to 30.0 mM results in a more than 2 orders of magnitude decrease in calculated concentrations of U(VI)-hydroxyl complexes (Figure 4b). The trend is very similar to what was observed in the experiments (Figure 3). We, therefore, believe that the major species responsible for U(VI) reduction by sulfide are U(VI)-hydroxyl complexes and eq 6 could be simplified to the following:

d[U(VI)] ) dt 2+ + (k1[UO2+ 2 ] + k2[UO2OH ] + k3[(UO2)2(OH)2 ]

-

+ k4[(UO2)3(OH)5+])[S2-]XT

(7)

Because it is difficult to unambiguously determine the individual rate constants, k1, k2, k3, and k4, using our experimental data sets, we shall use the total concentration of all U(VI)-hydroxyl complexes, (∑[U(VI)]-hydroxyl)), in the interpretation of experimental results under various pH and [CO32-]T conditions below. The total concentration is defined as the following:

∑ [U(VI) - hydroxyl] ) ([UO

2+ 2 ] + +

[UO2OH ] + 2 × [(UO2)2(OH)2+ 2 ]

+ 3 × [(UO2)3(OH)2+ 5 ]) VOL. 40, NO. 15, 2006 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 4. Variations of U speciation with increasing [CO32-]T. (a) U(VI)-carbonate complexes; (b) U(VI)-hydroxyl complexes (pH 6.89; [U(VI)]0 ) 168.3 µM).

FIGURE 5. (a) Decrease of [U(VI)] with time at various pHs. (b) Initial rates of U(VI) reduction and calculated changes of ∑[U(VI)hydroxyl] as a function of pH under comparable conditions. The data at pH 6.89 were the averages of duplicated tests (initial reactant concentrations: [U(VI)]0 ) 168.3 µM; [S2-]T ) 2.0 mM; [CO32-]T ) 4.0 mM). The effect of increasing [CO32-]T on ∑[U(VI)]-hydroxyl] at three pHs is illustrated in Figure 3 (the curves). ∑[U(VI)]hydroxyl] decreases dramatically with increasing [CO32-]T, which is very similar to the abrupt decreases of the observed [U(VI)]r/[U(VI)]0 with the increasing [CO32-]T (Figure 3, the dots). Under each given pH, there exists an overlapped range of [CO32-]T, in which the dramatic decreases of ∑[U(VI)]hydroxyl] correlate with the abrupt rate decreases of U(VI) reduction. At pH 6.89, when [CO32-]T is raised from 4.0 to 30.0 mM, calculated ∑[U(VI)]-hydroxyl] deceases from 1.4 × 10-7 M to 1.3 × 10-9 M; in comparison, [U(VI)/U(VI)0] decreases from 64 to 9%. At pH 8.01, increasing [CO32-]T from 1.5 to 4.0 mM results in a decrease of ∑[U(VI)]-hydroxyl] from 1.4 × 10-7 to 4.1 × 10-9 M; the experimentally measured decrease in [U(VI)]r/[U(VI)]0 is from 68 to 0.6%. Similarly at 4670

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pH 9.06, when [CO32-]T is raised from 0.2 to 2.0 mM, ∑[U(VI)]-hydroxyl] deceases from 1.3 × 10-8 to 2.0 × 10-9 M; the observed decrease of [U(VI)]r/[U(VI)]0 is 54 to 2%. Effect of pH at Fixed [CO32-]T on U(VI) Speciation and Reactivity. The effect of pH on U(VI) reduction at fixed [CO32-]T of 4.0 mM is shown in Figure 5a. With increasing pH from 6.37 to 9.06, the rate of U(VI) reduction was significantly reduced. At pH 6.37 and 6.89, the reductions were near completion within 2 h; while at pH 7.31, only 25% U(VI) was reduced in the same time period. At pH 8.01 or 9.06, no significant change in [U(VI)] was observed. This result agreed with the experiments on the effect of [CO32-]T (Figure 3), where significant U(VI) reduction was observed at pH 8.01 and 9.06 only when [CO32-]T was much lower than 4.0 mM. The change of ∑[U(VI)-hydroxyl] with pH was computed under [U(VI)]0 at 168.3 µM and [CO32-]T at 4.0 mM (Figure 5b). As pH increases, ln ∑[U(VI)-hydroxyl] is decreased in a near linear fashion. The initial rates of U(VI) reduction at various pHs, which were calculated from the experimental data in Figure 5a, are indicated by symbols in Figure 5b. Comparing ∑[U(VI)-hydroxyl] with the initial rates of U(VI) reduction at various pHs explicitly shows a near direct correlation between ∑[U(VI)-hydroxyl] and the initial rates of U(VI) reduction. The result supports our premise that the reactive species in U(VI) reduction by sulfide is through U(VI)hydroxyl species. Experiments examining the effect of increasing initial U(VI) concentrations were also conducted at fixed ratio of [U(VI)]0 to [CO32-]T (Supporting Information: Figure S3). When the increase in [U(VI)]0 was accompanied by increasing [CO32-]T, the reaction occurred quickest at the lowest [U(VI)]0. This result appeared to be counterintuitive at a first glance, but it agreed with U(VI)-hydroxyl species being the reactive species for U(VI) reduction and illustrated the importance of carbonate in controlling the rate of U(VI) reduction by sulfide. In summary, aqueous U(VI) reduction by sulfide occurs with a stoichiometric ratio of nearly 1:1 and yields nanosized uraninite as product. The rate of U(VI) reduction is greatly influenced by solution pH and [CO32-]T. Most of aqueous U(VI) could be reductively precipitated within several hours under neutral to weakly acidic pH at [CO32-]T up to 4.0 mM. At lower [CO32-]T, U(VI) reduction could occur quickly even at pH as high as 9. Since the total carbonate in many aqueous environments is around 1 mM at pH∼7.0 (30), the results of this study indicates that U(VI) could be reduced by sulfide in those environments with relatively low carbonate. Reductive immobilization of uranium by hydrogen sulfide, therefore, should be applicable for soil and water remediation, similar to reductive chromium immobilization (23, 31, 32).

The rate of U(VI) reduction is proportional to the total concentration of U(VI)-hydroxyl species; thus, the influence of pH and [CO32-]T on the rate of U(VI) reduction can be predicted by speciation calculation. Such information is critical to the design of the uranium remediation system based on sulfide reduction and to understanding U redox transformation and biogeochemical cycles in the environment. Further study is needed to elucidate the effect of other subsurface constituents, such as natural organic matter and phosphate, on the U(VI) reduction process. Potential transport and reoxidation of uraninite nanoparticles should also be addressed to assess the long-term U stability following its reductive immobilization.

Acknowledgments This work was supported by the United States Department of Energy under the Environmental Management Science Program (EMSP) (grant no. DE-FG02-03ER63616). TEM imaging by H.X. was supported by the Graduate School of the University of Wisconsin. We thank Professors Silvia Jurisson and Judy Wall at the University of Missouri-Columbia and Dr. Edward C. Thornton at Pacific Northwest National Laboratory for their valuable support and comments on this work.

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Supporting Information Available

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Figures S1-S3 and Table S1. This material is available free of charge via the Internet at http://pubs.acs.org.

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Literature Cited (1) Riley, R. G.; Zachara, J. M.; Wobber, F. J. Chemical Contaminants on DOE Lands and Selection of Contaminant Mixtures for Subsurface Science Research; U.S. Department of Energy: Washington, DC, 1992. (2) Bradford, G. R.; Bakhtar, D.; Westcot, D. Uranium, vanadium, and molybdenum in saline waters of California. J. Environ. Qual. 1990, 19, 105-108. (3) Giammar, D. Geochemistry of Uranium at Mineral-Water Interfaces: Rates of Sorption- Desorption and DissolutionPrecipitation Reactions. Ph.D. Dissertation. California Institute of Technology, 2001. (4) Langmuir, D. Uranium solution-mineral equilibria at low temperatures with applications to sedimentary ore deposits. Geochim. Cosmochim. Acta 1978, 42, 547-569. (5) Waite, T. D.; Davis, J. A.; Fenton, B. R.; Payne, T. E. Approaches to modeling uranium(VI) adsorption on natural mineral assemblages. Radiochim. Acta 2000, 88, 687-693. (6) Moyes, L. N.; Parkman, R. H.; Charnock, J. M.; Vaughan, D. J.; Livens, F. R.; Hughes, C. R.; Braithwaite, A. Uranium uptake from aqueous solution by interaction with goethite, lepidocrocite, muscovite, and mackinawite: an X-ray aborption spectroscopy study. Environ. Sci. Technol. 2000, 34, 1062-1068. (7) Wazne, M.; Korfiatis, G. P.; Meng, X. Carbonate effects on hexavalent uranium adsorption by iron oxyhydroxide. Environ. Sci. Technol. 2003, 37, 3619-3624. (8) Lovley, D. R.; Phillips, E. J. P.; Gorby, Y. A.; Landa, E. R. Microbial reduction of uranium. Nature 1991, 350, 413-416. (9) Lovley, D. R.; Widman, P. K.; Woodward, J. C.; Phillips, E. J. P. Reduction of uranium by cytochrome c3 of Desulfovibrio vulgaris. Appl. Environ. Microbiol. 1993, 59, 3572-3576. (10) Payne, R. B.; Casalot, L.; Rivere, T.; Terry, J.; Larsen, L.; Giles, B. J.; Wall, J. D. Interaction between uranium and the cytochrome c3 of Desulfovibrio desulfuricans strain G20. Arch. Microbiol. 2004, 181, 398-406. (11) Brooks, S. C.; Fredrickson, J. K.; Carroll, S. L.; Kennedy, D. W.; Zachara, J. M.; Plymale, A. E.; Kelly, S. D.; Kenneth M. Kemner, K. M.; Fendorf, S. Inhibition of bacterial U(VI) reduction by calcium. Environ. Sci. Technol. 2003, 37, 1850-1858. (12) Wielinga, B.; Bostick, B.; Hansel, C. M.; Rosenzweig, R. F.; Fendorf, S. Inhibition of bacterially promoted uranium reduction: ferric (hydr)oxides as competitive electron acceptors. Environ. Sci. Technol. 2000, 34, 2190-2195. (13) O’Loughlin, E.; Kelly, S. D.; Cook, R. E.; Csencsits, R.; Kemner, K. M. Reduction of uranium(VI) by mixed iron(II)/iron(III)

(23)

(24)

(25)

(26)

(27)

(28)

(29)

(30)

(31)

(32)

hydroxide (green rust): formation of UO2 nanoparticles. Environ. Sci. Technol. 2003, 37, 721-727. Liger, E.; Charlet, L.; Cappellen, P. V. Surface catalysis of uranium(VI) reduction by iron(II). Geochim. Cosmochim. Acta 1999, 63, 2939-2955. Gu, B.; Liang, L.; Dickey, M. J.; Yin, X.; Dai, S. Reduction precipitation of uranium(VI) by zero-valence iron. Environ. Sci. Technol. 1998, 32, 3366-3373. Barnes, C. E.; Cochran, J. K. Uranium geochemistry in estuarine sediments: controls on removal and release process. Geochim. Cosmochim. Acta 1993, 57, 555-569. Wersin, P.; Hochella, M. F., Jr.; Persson, P.; Redden, G.; Lechie, 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. Livens, F. R.; Jones, M. J.; Hynes, A. J.; Charnock, J. M.; Mosselmans, J. F. W.; Hennig, C.; Steele, H.; Collison, D.; Vaughan, D. J.; Pattrick, R. A. D.; Reed, W. A.; Moyes, L. N. X-ray ansorption spectroscopy studies of reactions of technetium, uranium and neptunium with mackinawite. J. Environ. Radioact. 2004, 74, 211-219. Kochenov, A. V.; Korolev, K. G.; Dubinchuk, V. T.; Medvedev, Y. L. Experimental data on the conditions of precipitation of uranium from aqueous solutions. Geochem. Int. 1978, 14, 8287. Beyenal, H.; Sani, R. K.; Peyton, B. M.; Dohnalkova, A. C.; Amonette, J. E.; Lewandowski, Z. Uranium immobilization by sulfate-reducing biofilms. Environ. Sci. Technol. 2004, 38, 20672074. Anderson, R.; Fleiser, M. Q.; LeHuray, A. P. Concentration, oxidation state, and particulate flux of uranium in the Black Sea. Geochim. Cosmochim. Acta 1989, 53, 2215-2224. Duff, M. C.; Amrhein, C.; Bertsch, P. M.; Hunter, D. B. The chemistry of uranium in evaporation pond sediment in the San Joaquin Valley, California, USA, using X-ray fluorescence and XANES techniques. Geochim. Cosmochim. Acta 1997, 61, 7381. Thornton, E. C.; Amonette, J. E. Hydrogen sulfide gas treatment of Cr(VI)-contaminated sediment samples from a plating-waste disposal site-implications for in-situ remediation. Environ. Sci. Technol. 1999, 33, 3, 4096-4101. Huang, F. Y. C.; Brady, P. V.; Lindgren, E. R.; Guerra, P. Biodegradation of uranium-citrate complexes: implications for extraction of uranium from soils. Environ. Sci. Technol. 1998, 32, 379-382. Teixeira, L. S. G.; Costa, A. C. S.; Ferreira, S. L. C.; Freitas, M. L.; Carvalho, M. S. Spectrophotometric determination of uranium using 2-(2-thiazolylazo)-p-cresol(TAC) in the presence of surfactants. J. Braz. Chem. Soc. 1999, 10 (6), 519-522. APHA, AWWA, WPCE Standard Methods for the Examination of Water and Wastewater, 15th ed; American Public Health Association: Washington, DC, 1985 Allen, H. E.; Fu, G.; Deng, B. Analysis of acid-volatile sulfide (AVS) and simultaneously extracted metals (SEM) for the estimation of potential toxicity in aquatic sediments. Environ. Toxicol. Chem. 1993, 12, 1441-1453. Kim, C.; Zhou, Q.; Deng, B.; Thornton, E.; Xu, H. Chromium(VI) reduction by hydrogen sulfide in aqueous media: stoichiometry and kinetics. Environ. Sci. Technol. 2001, 35, 2219-2225. Lan, Y.; Deng, B.; Kim, C.; Thornton, E. C.; Xu, H. Catalysis of elemental sulfur nanoparticles on chromium(VI) reduction by sulfide under anaerobic conditions. Environ. Sci. Technol. 2005, 39, 2087-2094. Van Geen, A.; Robertson, A. P.; Leckie, J. O. Complexation of carbonate species at the goethite surface: implications for adsorption of metal ions in natural waters. Geochim. Cosmochim. Acta 1994, 58, 2073-2086. Cantrell, K. J.; Yabusaki, S. B.; Engelhard, M. H.; Mitroshkov, A. V.; Thornton, E. C. Oxidation of H2S by iron oxides in unsaturated conditions. Environ. Sci. Technol. 2003, 37, 2192-2199. Hua, B.; Deng, B. Influences of water vapor on Cr(VI) reduction by gaseous hydrogen sulfide. Environ. Sci. Technol. 2003, 37, 4771-4777.

Received for review September 11, 2005. Revised manuscript received May 11, 2006. Accepted May 12, 2006. ES051804N

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