Reductive Immobilization of Uranium(VI) by Amorphous Iron Sulfide

Oct 31, 2008 - Environmental Science & Technology 2017, 51 (19) , 10954-10964. ... Vasso G. Alexandratos, Thilo Behrends, and Philippe Van Cappellen ...
0 downloads 0 Views 249KB Size
Download PDF
Environ. Sci. Technol. 2008, 42, 8703–8708

Reductive Immobilization of Uranium(VI) by Amorphous Iron Sulfide BIN HUA AND BAOLIN DENG* Department of Civil and Environmental Engineering, University of Missouri, Columbia, Missouri 65211

Received May 3, 2008. Revised manuscript received September 20, 2008. Accepted September 25, 2008.

Batch experiments were used to evaluate the reductive immobilization of hexavalent uranium (U(VI)) by synthesized, amorphous iron sulfide (FeS) in the anoxic environment. The tests were initiated by spiking 168.0 µM U(VI) to 0.18 g/L FeS suspensions under a CO2-free condition with pH varied from 5.99 to 10.17. The immobilization rate of U(VI) was determined by monitoring the changes of aqueous U(VI) concentration, and the reduction rate of U(VI) associated with FeS was determined by the difference between the total spiked U(VI) and the extractable amount of U(VI) by 25 mM NaHCO3 solution. The results showed that a rapid removal of U(VI) from the aqueous phase occurred within 1 h under all pH conditions accompanied by a simultaneous release of Fe(II), whereas the reduction of U(VI) associated with FeS took hours to over a week for completion. The reduction rate was greatly increased with decreasing pH within the examined pH range. Product analysis by X-ray photoelectron spectroscopy showed the formation of U3O8/U4O9/UO2, polysulfide, and ferric iron.

Introduction Uptake and reduction of the aqueous-phase U(VI) by a variety of minerals has a strong influence on uranium mobility in the subsurface environment, with its immobilization processes often controlled by water chemistry, such as solution pH and redox potential (1-8). In an anoxic environment, the reduction of U(VI) to U3O8/ U4O9/UO2 may occur when certain reductants are present including reduced sulfur and iron species. For example, it has been established that formations of at least some uranium ores are linked to the iron redox cycling (9). Observations on the reduction of U(VI) by Fe(II) as well as sulfide species in the laboratory are not all consistent, and the results depended heavily on the reaction conditions. It is generally agreed that Fe(II)(aq) (“(aq)” denotes species in the aqueous phase) cannot directly reduce U(VI)(aq) under ambient environmental conditions (10), but both sorbed Fe(II) by iron oxide and structural Fe(II) in minerals can reduce U(VI) to U(IV) (10-12). Other studies reported that iron(III) oxides could oxidize U(IV) and generate sorbed Fe(II) species (13, 14). Sulfide ions are found capable of reducing U(VI) (15-18), and the major factors influencing the kinetics of U(VI) reduction include the solution pH (examined from 6.37 to 9.06) and bicarbonate concentration (examined from 0.0 to 30.0 mM) (18). The reduction rates increased with decreasing pH and carbonate concentration, and the reaction products * Corresponding author phone: (573) 882-0075; fax: (573) 8824784; e-mail: [email protected]. 10.1021/es801225z CCC: $40.75

Published on Web 10/31/2008

 2008 American Chemical Society

were primarily nanosized uraninite (UO2) and elemental sulfur. Several minerals containing reduced sulfur species have also been found to reduce U(VI), although the information is mostly qualitative in nature. For instance, reduction of U(VI) by galena and pyrite under anoxic conditions yielded a mixture of U(VI)/U(IV) species and polysulfides (6). Partial reduction of U(VI) was also detected on mackinawite (7). Built on our previous studies on the aqueous-phase U(VI) reduction by sulfide (18), this work aims to examine the interactions of U(VI) and a synthesized, amorphous iron sulfide (FeS) in a CO2-free anoxic environment. While field sites with U contamination are rarely CO2-free, we need to understand the behaviors of U(VI) in this relatively “pure” system first to quantify the effects of CO2 as well as other system constituents (e.g., calcium and iron). Much effort in this study was devoted to identify the reaction products and mechanisms.

Experimental Methods All experiments were conducted under an anoxic condition in a glovebox (Coy Laboratory Products Inc.) with ∼5% H2 balanced with ∼95% N2. Solutions were prepared with deionized and distilled water (DDW; 18.2 MΩ cm, Millipore Co.), pretreated following a reported procedure to minimize the impact of oxygen (18). Glassware was cleaned with 1 N HCl and rinsed with DDW prior to use. U(VI) (uranyl nitrate) standard solution (997 ( 2 mg/L), sodium sulfide, 1,10phenanthroline, and ferrous chloride were purchased from Fisher Scientific Co. All reagents were of ACS reagent grade and were used without further purification. FeS Preparation and Characterization. FeS stock suspension was prepared by reaction of 1.0 M FeCl2 and 1.0 M Na2S (19). The synthesized FeS particles were washed multiple times with DDW until the soluble ferrous iron concentration ([Fe(II)]aq) was less than 10 µM and then stored in a polyethylene bottle. The FeS solid concentration was determined to be 1.80 g of FeS/L by the following procedure: (1) pipet 1.00 mL of suspension into a 40 mL glass vial, (2) dissolve FeS with several drops of 1.0 M HCl, (3) transfer the solution into a 1 L volumetric flask and dilute to the mark with DDW, and (4) determine [Fe(II)]aq by the phenanthroline spectrophotometric method (20). The precision of delivering FeS suspension by pipet was assessed by six replicate tests, in which 1.00 mL of FeS suspension was pipetted into each 40 mL glass vial and analyzed for [Fe(II)]aq, showing a standard deviation of less than 6%. The concentrations of Fe(II) in diluted FeS suspension (0.18 g/L) at pH from 6.01 to 10.10 were determined to evaluate the dissolution of FeS under relevant pH conditions. A portion of the FeS suspension was freeze-dried (4.5 L benchtop freeze-dryer, Labconco Co.) in an airtight container. The obtained FeS particles were stored in the anoxic glovebox prior to crystal structure and specific surface area analyses. To determine the crystal structure, approximately 0.5 g of FeS particles was placed in a sample holder (24.5 mm in diameter and 1 mm in depth) and analyzed on a MiniFlex automated, microprocessor-controlled X-ray powder diffractometer, with a Cu KR X-ray source and a semiconductor detector (operated at 15 mA and 30 kV). The sample was scanned from 20 to 40° (2θ range) in steps of 0.01° and at a scan rate of 2.0 deg/min. The result confirmed that the sample was amorphous iron sulfide (21). The specific surface area was measured on a PMI automated Brunauer-Emmett-Teller (BET) sorptometer (Porous Materials, Inc.) with N2 adsorption at -196 °C, yielding a value of 31 m2/g. VOL. 42, NO. 23, 2008 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

8703

FIGURE 1. Changes of [U(VI)]aq and [Fe(II)]aq with time and pH: (a) U(VI) uptake was almost completed in 1 h for all pH values, (b) the solution pH has a strong impact on [Fe(II)]aq ([U(VI)]0 ) ∼168.0 µM, [FeS] ) 0.18 g/L)), (c) background Fe(II) concentration measured at 1 h of equilibrium time after adjustment of the pH of the 0.18 g/L FeS suspension, (d) [U(VI)]0 versus [Fe(II)]aq at pH 6.90. The [U(VI)]0/[Fe(II)]aq ratio is 0.85 ( 0.03 (within the 90% confidence level). The variation for the tests with [U(VI)]0 ) 168.3 µM is (4.4%. Experiments on U(VI) Uptake and Reduction. Each test began by mixing 1.00 mL of FeS stock suspension (1.80 g of FeS/L), 7.0 mL of DDW, and 2.00 mL of U(VI) solution in a 25 mL brown vial, with the pH of the mixture controlled by preadjusting the pH of DDW. The pH of DDW was altered with HCl/NaOH. At the end of the reaction, the pH was checked again, showing that it never exceeded 0.40 pH unit from the preset values. Mixing was maintained by magnetic stirring (Variomag, Poly 15), and the temperature was controlled at 22 ( 1 °C. The reaction progress was monitored by analyzing [U(VI)]aq (18, 22) and [Fe(II)]aq (20) as a function of time. At each time point, one of the parallel reaction vials was sacrificed to collect 3 mL of supernatant by filtration through a 0.2 µm nylon filter (Fisher Scientific Co.). A 1.00 mL sample of the filtrate was used to determine [U(VI)]aq, and another 1.00 mL sample of the filtrate was used to determine [Fe(II)]aq. The difference between the amount of initially spiked hexavalent uranium ([U(VI)]0) and the aqueous-phase concentration ([U(VI)]aq) was considered the amount associated with FeS ([U]uptake ) [U(VI)]0 - [U(VI)]aq). U(VI) associated with FeS could subsequently be reduced by FeS, a process leading to the immobilization of uranium in the subsurface. We, therefore, devoted significant effort to differentiate the oxidation state of U associated with FeS. To begin with, 10.0 mL of NaHCO3 (50 mM) was added to a reaction vial parallel to the one used for measuring [U(VI)]aq, and then U(VI) in the extractant ([U(VI)]ex) was measured after 2 h of extraction. Our preliminary experiments showed that the extraction efficiency did not change after 30 min of extraction. We considered [U(VI)]ex represented the sum of [U(VI)]aq and [U(VI)]s (i.e., [U(VI)]ex ) [U(VI)]aq + [U(VI)]s). As a result, if no U(VI)(s) were reduced, [U(VI)]ex would always be equal to [U(VI)]0 in the system. Our observation, however, showed that [U(VI)]ex decreased with the reaction time, implying the reduction of U(VI) to uranium in lower oxidation states and formation of products such as uraninite that were not extractable by the bicarbonate solution in the anoxic 8704

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 42, NO. 23, 2008

environment (10). Consequently, the difference between [U(VI)]0 and [U(VI)]ex was considered the amount of reduced U. The possibility of reoxidation of the reduced U on FeS was also investigated by additional experiments. Four reaction vials were prepared at pH 6.90 following the U reduction procedure described above. After 6 h, two vials were sacrificed to measure [U(VI)]ex, showing that U(VI) uptake and reduction was completed within the time period. The other two samples were taken out of the anoxic glovebox and exposed to the ambient air with 0.21 atm of oxygen partial pressure. The samples were mixed by a magnetic stirring bar to facilitate oxygen transfer. After a 24 h exposure, [U(VI)]aq and [U(VI)]s in these samples were measured to determine whether the reduced U was reoxidized. All experimental runs were duplicated. The average results were reported, and the associated error bars in Figures 1 and 2 represent the variations of data from the averages. X-ray Photoelectron Spectroscopy (XPS) analysis. XPS analysis was conducted on a KRATOS model AXIS 165 XPS spectrometer with an argon sputtering KRATOS Minibeam ion source. The sample was prepared by mixing 50 mL of 1683 µΜ U(VI) with 50 mL of 18.0 g/L FeS at pH 6.90 and 5.99 in a 200 mL polyethylene bottle for 24 h. After removal of most of the supernatant, FeS was washed with pretreated DDW three times and freeze-dried in an airtight container for over 24 h. The sample was initially tapped onto a sample supporting plate, which was then placed in a chamber and vacuumed to ∼10-9 Torr. XPS measurements were conducted with monochromatic aluminum at 1486.6 eV, pass energies between 20 and 50 eV. The XPS photo peak energies were calibrated with C 1s, and the data analyses were performed through the use of the software CasaXPS (version 2.3.14).

Results U(VI) Uptake by FeS. As shown in Figure 1a, uptake of U(VI)(aq) by FeS was over 90% within 15 min at pH values

FIGURE 2. Changes of [U(VI)]ex with time at different pH values. The inset highlights [U(VI)]ex changes at the early stage of the reaction in the systems with pH 5.99 and 6.90. The steady decreases in [U(VI)]ex are indicative of the reduction of U(VI)(s). from 5.99 to 8.82 and within 60 min at pH 10.17. After 3 h, U(VI)(aq) was below the detection limit of 0.4 µM for all systems. The uptake of aqueous U(VI) by FeS particles was accompanied by the release of ferrous iron (Fe(II)) into the solution (Figure 1b). The final concentrations of soluble Fe(II) released (i.e., [Fe(II)]aq) at pH 5.99 and 6.90 were 185.5 and 165.8 µM, respectively. It was very interesting to note that these values were very close to the total amount of U(VI) associated with FeS (i.e., ∼168 µM/L). At a higher pH value of 7.99, [Fe(II)]aq jumped from its background value of less than 10.0 to 47.0 µM within 15 min, increased to 56.0 µM in 1 h, and then remained nearly constant. As indicated in Figure 1c, background [Fe(II)]aq did not contribute significantly to the jump of [Fe(II)]aq at pH less than 8. When the pH was at 8.82, [Fe(II)]aq was increased to approximately 18.0 µM within 1 h and then remained constant, while there was no discernible increase in [Fe(II)]aq at pH 10.17 for over 8 h. To establish the quantitative relationship between the uptake of U(VI) and release of Fe(II), we conducted an additional series of experiments at pH 6.90 with increasing [U(VI)]0 from 42.0 to 420.0 µM. We found the amount of [Fe(II)]aq released was directly proportional to the total amount of U(VI) removed from the solution (R2 ) 0.997), with a [U(VI)]0/[Fe(II)]aq ratio of 0.85 ( 0.03 (within the 90% confidence level), as shown in Figure 1d. Reduction Rates of the U(VI) Associated with FeS at Various pH Values. The reduction of the U(VI) associated with FeS was monitored by measuring the concentration of extractable U(VI) ([U(VI)]ex) as a function of time. As illustrated in Figure 2, [U(VI)]ex decreased steadily with time, with the reduction rates strongly dependent on the solution pH. For instance, in the suspensions with pH 5.99 or 6.90, [U(VI)]ex was decreased to less than 50% of [U(VI)]0 after 2 h and near nil after 4 h. At pH 7.99, [U(VI)]ex was decreased by approximately 50% after 4 h and to less than 10.0 µM after 18 h. In comparison, 50% U(VI) reduction took approximately 20 h at pH 8.82 and 110 h at pH 10.17. Potential reoxidation and dissolution of reduced U was assessed by exposing the reduced samples at pH 6.99 to the ambient atmosphere for 6 h under complete mixing conditions. The results showed that 62% of the reduced U was reoxidized, of which 11% was found in the aqueous phase and 89% associated with the solid phase. Reaction Products. The chemical compositions of the solid products were analyzed by XPS, using pure FeS powder as a reference material whenever relevant. The results for the sample prepared at pH 6.90 are presented here, while

the analogous results for the sample prepared at pH 5.99 are presented as Supporting Information (Figure S1). The binding energies of U 4f, S 2p, and Fe 2p were calibrated with that of C 1s (285 eV). As shown by the survey spectrum (Figure 3a), the major identified peaks were for iron, sulfur, carbon, and uranium; the other peaks were for Na, Cl, and O. The XPS spectra of U 4f for the reacted sample at pH 6.90 are shown in Figure 3b, with the FeS-U surface line representing the outer surface scan and the FeS-U interior line the scan after argon etching to remove ∼75 Å of the surface. Scans after argon etching should represent samples free of alterations during sample handling. The spectrum for the outer surface was similar to that of U4O9 or U3O8 with a characteristic peak at 380.8 eV (6, 23). The binding energy for the etched sample was consistent with uraninite (UO2) with a characteristic peak at 380.3 eV (6, 23, 24). In the XPS spectra of S 2p (Figure 3c), FeS-U surface and FeS surface lines represent the outer surface of FeS loaded with U and pure FeS, respectively, and FeS-U interior and FeS interior lines the spectra after slight etching of 75 Å off the surface. No significant difference existed with and without the surface etching for the sample loaded with U and the sample of pure FeS. The S 2p spectra of FeS, both surface and interior, could be assigned to S2- with a characteristic peak at 161.9 eV (25-27). However, after reaction with U, the S 2p spectra showed split peaks centered at 161.4 and 162.4 eV, indicating the presence of polysulfide (26, 27). The XPS spectra of Fe 2p for FeS loaded with U and pure FeS are shown in Figure 3d, with and without surface etching. The FeS surface line exhibited a binding energy of 710.4 eV, and the FeS-U surface line displayed a binding energy of 710.9 eV. The results of the deconvolution analysis (Figure S2, Supporting Information) showed that the FeS surface was composed of Fe(II)-S (710.4 eV) and FeS2 (707.6 eV), while the FeS-U surface comprised Fe(III)-O (713.3 eV), Fe(III)-S/Fe(II)-S (710.7 eV), and FeS2 (707.8 eV).. The spectrum of the FeS-U interior was very similar to that of the FeS interior, indicating there was no major change happening before and after U(VI) uptake (6, 25-27).

Discussion From the above observations, we propose that the reductive immobilization of U(VI) by amorphous FeS proceeds by a two-step mechanism: (step 1) uptake of uranyl ion by FeS, accompanied by the release of Fe(II), (step 2) reduction of U(VI) by sulfide and Fe(II) to form U3O8/U4O9/UO2 on the surface of FeS. VOL. 42, NO. 23, 2008 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

8705

FIGURE 3. XPS spectra for FeS with and without loaded U: (a) FeS-U surface survey (major peaks for Na, Cl, and O not indicated), (b) U 4f spectra, (c) S 2p spectra, (d) Fe 2p spectra. The samples were prepared at pH 6.90. Uptake of U(VI) by FeS. The results in Figure 1b indicate that the uptake of U(VI) from the aqueous phase by FeS is accompanied by the release of Fe(II) under the neutral to acidic pH condition. Additional experiments at pH 6.90 shown in Figure 1d further verify that a linear correlation exists between the amount of U(VI) uptake and Fe(II) release: [U]uptake ) 0.85[Fe(II)]released (R2 ) 0.995). We, therefore, believe the uptake of U(VI) is through an ion exchange mechanism under the experimental condition; i.e., Fe(II) is replaced by U(VI), leading to the observed U(VI) phase transfer (step 1). An alternative explanation for the disappearance of U(VI) from the aqueous phase could also be proposed: U(VI) is precipitated as a mineral such as schoepite. Under the circumneutral condition, however, this hypothesis is not supported by our experiments, which showed no U(VI) precipitation in the FeS-free solution with ∼168 µM U(VI) at pH 6.90. Our observation is consistent with what was reported in the literature (28, 29). Sani et al. (28) found that no U(VI) (∼130 µM) precipitation occurred under the abiotic condition in a pH 7.2, mineral-, CO2-, and O2-free system. Fredrickson et al. (29) also showed that in a system with 125 µM U(VI) at pH 7 (30 mM PIPES buffered), U(VI) primarily existed in solution as UO2OH+(aq) or UO2(OH)2(aq). It is likely that 8706

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 42, NO. 23, 2008

nucleation under the experimental conditions requires a higher U(VI) concentration than the solubility limit of minerals such as schoepite, so no U(VI) precipitation was observed. In comparison with U(VI) sorption at pH 5.99 and 6.90 where the ion exchange mechanism is operative, interpretation of U(VI) uptake at higher pH is confounded by the fact that the release of Fe(II) is much smaller than the stoichiometric amount of U(VI) uptake (Figure 1b). A possible explanation is that the released Fe(II) hydrolyzes to form Fe(OH)2 precipitates, decreasing the soluble Fe(II) that could be detected. We have tested this hypothesis by calculating the equilibrium distribution of Fe(II) as a function of pH at a total Fe(II) concentration of 168.0 µM, using a common speciation program, MINEQL+ (version 4.07). As shown in Figure S3 (Supporting Information), soluble ferrous iron ([Fe(II)]aq) stays constant at 168.0 µM from pH 5.0 to pH 7.6 and then decreases by 3 orders of magnitude when the pH is increased from 7.7 to 8.7. The calculated [Fe(II)]aq values are 30, 0.7, and 4.8 × 10-3 µM at pH 8.0, 8.8, and 10.2, respectively. In comparison, the observed soluble ferrous iron concentrations are 56, 18.9, and 7.9 µM, respectively. Given the detection limit of Fe(II) (0.2 µM) and the variation

of background [Fe(II)]aq (∼10 µM) (Figure 1c), the results implied that the ion exchange process might still occur under the alkaline condition, but lower [Fe(II)]aq is observed because of its partial precipitation. This explanation is not conclusive, however, because we cannot rule out the possibility that partial precipitation of U(VI) occurs in the alkaline solution, without exchange of Fe(II). It appears that the involvement of U(VI) in an ion exchange process is not unique to only amorphous FeS. Our preliminary experiments show that, at pH 6.20, U(VI) uptake by lead sulfide (PbS, synthesized by mixing equal molar concentrations of Pb(NO3)2 and Na2S solutions) is also accompanied by the release of lead(II) and complete reduction within 24 h, 145.0 µM Pb(II) being released with the uptake of 168.0 µM U(VI). Our observations agree with the result by Wersin et al. (6), who found that the sorption of U(VI) by galena was accompanied by the release of lead and partial reduction of U(VI). They also found that [Pb(II)]aq was proportional to the adsorbed [U(VI)], although [Pb(II)]aq is much less than the sorbed [U(VI)]. Reduction of U(VI)(s). We devoted significant effort to quantify the reduction rate of U(VI) by FeS based on measuring the NaHCO3-extractable U on FeS as a function of the reaction time. We considered [U(VI)]ex represented the sum of [U(VI)]aq and [U(VI)]s. Several lines of evidence supported this argument: (i) U(VI) could form strong complexes with carbonate ions (1, 4), leading to the complete desorption of U(VI) from ferrihydrite (30) and hematite (10) under alkaline conditions. (ii) Our preliminary tests showed that no U(VI) uptake by FeS occurred in the presence of 25 mM NaHCO3 during a testing period of 72 h. (iii) While complete uptake of U(VI) by FeS always occurred in less than 1 h of reaction (refer to Figure 1a), over 91% of initially spiked U(VI) could still be extracted at pH 8.82 after 1 h of reaction and over 87% of initially spiked U(VI) extracted at pH 10.17 even after 2 h (as shown in Figure 2), suggesting that U(VI)(s) was extractable by the bicarbonate solution. (vi) Interference from the oxidative dissolution of U(IV) is unlikely, because carbonate does not dissolve UO2 (e.g., uraninite) even at concentrated concentrations (31), and the reoxidation of reduced U could not occur under the anoxic condition for extraction. In the system examined here, potential reductants include Fe(II)(aq), Fe(II)(s), Fe(II)(ad), HS-(aq), and HS-(s). When we assume a first-order kinetics with respect to each species, a generic expression for the total rate of U(VI) reduction can be formulated as -

d[U(VI)]T ) k1[Fe(II)]aq[U(VI)]aq + k2[HS-]aq[U(VI)]aq + dt k3[Fe(II)]ad[U(VI)]aq + k4[HS-]s[U(VI)]aq + k5[Fe(II)]s[U(VI)]s + k6[HS-]s[U(VI)]s (1)

The reduction of U(VI)(aq) by Fe(II)(aq) is thermodynamically unfavorable unless under strong acidic conditions (32). Furthermore, Liger et al. (10) experimentally demonstrated that Fe(II)(aq) could not reduce U(VI)(aq) at pH 7.5. As a result, the first term on the right-hand side (RHS) in eq 1 can be ruled out. Aqueous HS- can reduce U(VI) under certain conditions with a typical half-life of hours (17, 18). However, the second term is unlikely to be a major mechanism here. As shown in Figure 1a, over 90% of U(VI)(aq) has been removed by FeS within 15 min (except for pH 10.17, in which case a slightly longer time was involved), leading to virtually complete transformation of U(VI)(aq) to U(VI)(s). Consequently, the second term as well as the fourth term on the RHS of eq 1 can be omitted too. If the sorbed Fe(II) (Fe(II)ad) had been the major species for the reduction of U(VI) in our system, [Fe(II)]aq would

have been decreased due to the shift of Fe(II) sorption equilibrium. The results in Figure 1b, however, show that [Fe(II)]aq is constant, which does not support this mechanism. As a result, the third term on the RHS can be eliminated. In an experiment to evaluate the influence of sediment on uranium sorption, Liu et al. (33) also did not observe any reduction of aqueous or adsorbed U(VI) by sorbed Fe(II). The relative contribution of the remaining two terms on the RHS of eq 1 cannot be differentiated at this stage. Further studies with extended X-ray absorption fine structure (EXAFS) spectroscopy may help resolve the detailed mechanisms for U(VI) reduction. In summary, this study has demonstrated that the reductive immobilization of U(VI) by synthesized, amorphous iron sulfide follows an ion exchange and reduction processes. The uptake of U(VI) is rapid and insensitive to pH, while reduction strongly depends on the pH, with a half-time ranging from a few hours at circumneutral pH to over 100 h at pH 10.17. Previous studies have shown that iron oxides could be reduced by gaseous hydrogen sulfide to form iron sulfide (FeS); therefore, an FeS-permeable reactive barrier could be created by injecting hydrogen sulfide gas into the contaminated sites with high iron contents (26). According to this research, a permeable reactive barrier of FeS should have a potential to prevent the subsurface migration of U(VI) through reduction. Since FeS could also be produced coupled with microbially mediated sulfate reduction (34, 35), reductive U(VI) immobilization through the abiotic mechanism could occur in principle. The significance of this abiotic contribution, however, cannot be determined at this stage because the concentrations of CO2, calcium, and iron could significantly affect the immobilization process. Future work is needed to evaluate the impact of the complex groundwater and soil constituents on the reduction kinetics.

Acknowledgments This work was supported by the United States Department of Energy under the Environmental Remediation Science Program (ERSP) (Grant No. DE-FG02-03ER63616). The support by Dr. John Yang at the Lincoln University of Missouri to this study is also greatly appreciated.

Supporting Information Available XPS spectra for FeS with and without loaded U prepared at pH 5.99, deconvolution of Fe 2p XPS spectra for FeS with and without loaded U prepared at pH 6.99, and calculated [Fe(II)]aq as a function of pH and its comparison with measured data. This material is available free of charge via the Internet at http://pubs.acs.org.

Literature Cited (1) Langmuir, D. Uranium solution-mineral equilibria at low temperatures with applications to sedimentary ore deposits. Geochim. Cosochim. Acta 1978, 42, 547–569. (2) Hsi, C. D.; Langmuir, D. Adsorption of uranyl onto ferric oxyhydroxides: Application of the surface complexation sitebinding model. Geochim. Cosmochim. Acta 1985, 49, 1931–1941. (3) Ho, C. H.; Doern, D. C. The sorption of uranyl species on the hematite sol. Can. J. Chem. 1985, 63, 1100–1105. (4) Ho, C. J.; Miller, N. H. Adsorption of uranyl species from bicarbonate solution onto hematite particles. J. Colloid Interface Sci. 1986, 110, 165–171. (5) Payne, T. E.; Waite, T. D. Surface complexation modeling of uranium sorption data obtained by isotope exchange techniques. Radiochim. Acta 1991, 52-53, 487–493. (6) Wersin, P., Jr.; Persson, P.; Redden, G.; Leckie, J. O.; Harris, D. W. Interaction between aqueous uranium(VI) and sulfide minerals: Spectroscopic evidence for sorption and reduction. Geochim. Cosmochim. Acta 1994, 58, 2829–2843. (7) Moyes, L. N.; Parkman, R. H.; Charnock, J. M.; Vaughan, D. J.; Livens, F. R.; Hughes, C. R.; Braithwaite, A. Uranium uptake VOL. 42, NO. 23, 2008 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

8707

(8) (9) (10) (11)

(12) (13) (14)

(15)

(16)

(17)

(18) (19) (20) (21) (22)

(23)

8708

from aqueous solution by interaction with goethite, lepidocrocite, muscovite, and mackinawite: An X-ray absorption spectroscopy study. Environ. Sci. Technol. 2000, 34, 1062–1068. Giammer, D. E.; Hering, J. G. Time scales for sorption-desorption and surface precipitation of uranyl on goethite. Environ. Sci. Technol. 2001, 35, 3332–3337. Posey-Dowty, J.; Axtmann, E.; Crerar, D.; Borcsik, M.; Ronk, A.; Woods, W. Releasing rate of uraninite and uranium roll-front ores. Econ. Geol. 1987, 82 (1), 184–194. Liger, E.; Charlet, L.; Van Cappellen, P. Surface catalysis of uranium(VI) reduction by iron(II). Geochim. Cosmochim. Acta 1999, 63, 2939–2955. O’Loughlin, E. J.; Kelly, S. D.; Cook, R. E.; Csencsits, R.; Kemner, K. M. Reduction of uranium(VI) by mixed iron(II)/iron(III) hydroxide (green rust): Formation of UO2 nanoparticles. Environ. Sci. Technol. 2003, 37, 721–727. Missana, T.; Maffiotte, C.; Miguel, G. Surface reactions kinetics between nanocrystalline magnetite and uranyl. J. Colloid Interface Sci. 2003, 261, 154–160. Nevin, K. P.; Lovley, D. R. Potential for nonenzymatic reduction of Fe(III) by electron shuttling in subsurface sediments. Environ. Sci. Technol. 2000, 34, 2472–2478. Istok, J. D.; Senko, J. M.; Krumholz, L. R.; Watson, D.; Bogle, M. A.; Peacock, A.; Change, Y.-J.; White, D. C. In situ bioreduction of technetium and uranium in a nitrate-contaminated aquifer. Environ. Sci. Technol. 2004, 38, 468–475. 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, 82– 87. 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. Cosochim. Acta 1997, 61, 73–81. 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, 2067– 2074. Hua, B.; Deng, B. Kinetics of uranium(VI) reduction by hydrogen sulfide in anoxic aqueous systems. Environ. Sci. Technol. 2006, 40, 4666–4671. Patterson, R. R.; Fendorf, S.; Fendorf, M. Reduction of hexavalent chromium by amorphous iron sulfide. Environ. Sci. Technol. 1997, 31, 2039–2044. APHA. Standard Methods for the Examination of Water and Wastewater, 17th ed.; American Public Health Association: Washington, DC, 1989. JCPDS. Powder Diffraction File Alphabetical Index, Inorganic Phases; JCPDS, International Center for Diffraction Data: Newtown Square, PA, 1987. 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. Allen, G. C.; Crofts, J. A.; Curtis, M. T.; Tucker, P. M.; Chadwick, D.; Hampson, P. J. X-ray photoelectron spectroscopy of some

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 42, NO. 23, 2008

(24)

(25)

(26)

(27)

(28)

(29)

(30)

(31)

(32)

(33)

(34)

(35)

uranium oxide phases. J. Chem. Soc., Dalton Trans. 1974, 1296– 1301. Chadwick, D. Uranium 4f binding energies studies by X-ray photoelectron spectroscopy. Chem. Phys. Lett. 1973, 21, 291– 294. Bukhtiyarova, G. A.; Bukhtiyarov, V. I.; Sakaeva, N. S.; kaichev, V. V.; Zolotovskii, B. P. XPS study of the silica-supported Fecontaining catalysts for deep or partial H2S oxidation. J. Mol. Catal. A 2000, 158, 251–255. 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. Ko, T. H.; Chu, H. Spectroscopic study on sorption of soil hydrogen by means of red soil. Spectrochim. Acta, Part A 2005, 61, 2253–2259. Sani, R. K; Peyton, B. M.; Amonette, J. E.; Geesey, G. G. Reduction of uranium(VI) under sulfate-reducing conditions in the presence of Fe(III)-(hydr)oxides. Geochim. Cosochim. Acta 2004, 68, 2639–2648. Fredrickson, J. K.; Zachara, J. M.; Kennedy, D. W.; Dong, H.; Onstott, T. C.; Hinman, N. W.; Li, S. Biogenic iron mineralization accompanying the dissimilatory reduction of hydrous ferric oxide by a groundwater bacterium. Geochim. Cosmochim. Acta 1998, 62, 3239–3257. Waite, T. D.; Davis, J. A.; Payne, T. E.; Waychunas, G. A.; Xu, N. Uranium(VI) adsorption to ferrihydrite: Application of a surface complexation model. Geochim. Cosmochim. Acta 1994, 58, 5465– 5478. Buck, E. C.; Brown, N. R.; Dietz, N. L. Contaminant uranium phases and leaching at the Fernald site in Ohio. Environ. Sci. Technol. 1996, 30, 81–88. Gu, B.; Liang, L.; Dickey, M. J.; Yin, X.; Dai, S. Reductive precipitation of uranium(VI) by zero-valent iron. Environ. Sci. Technol. 1998, 32, 3366–3373. Liu, C.; Zachara, J.; Zhong, L.; Kukkapdupa, R.; Szecsody, J. E.; Kennedy, D. W. Influence of sediment bioreduction and reoxidation on uranium sorption. Environ. Sci. Technol. 2005, 39, 4125–4133. Wu, W.-M.; Carley, J.; Gentry, T.; Ginder-Vogel, M. A.; Fienen, M.; Mehlhorn, T.; Yan, H.; Carroll, S.; Nyman, J.; Luo, J.; Gentile, M. E.; Fields, M. W.; Hickey, R. F.; Watson, D.; Cirpka, O. A.; Fendorf, S.; Zhou, J.; Kitanidis, P.; Jardine, P. M.; Criddle, C. S. Pilot-scale in situ bioremediation of uranium in a highly contaminated aquifer 2: U(VI) reduction and geochemical control of U(VI) bioavailability. Environ. Sci. Technol. 2006, 40, 3986–3995. Anderson, R. T.; Vrionis, H. A.; Ortiz-Bernad, I.; Resch, C. T.; Long, P. E.; Dayvault, R.; Karp, K.; Marutzky, S.; Metzler, D. R.; Peacock, A.; White, D. C.; Lowe, M.; Lovley, D. R. Stimulating the in situ activity of Geobacter species to remove uranium from the groundwater of a uranium-contaminated aquifer. Appl. Environ. Microbiol. 2003, 69 (10), 5884–5891.

ES801225Z