Reduction of Uranium (VI) by Soluble Iron (II) Conforms with

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Reduction of Uranium(VI) by Soluble Iron(II) Conforms with Thermodynamic Predictions Xin Du,† Benjaporn Boonchayaanant,† Wei-Min Wu,† Scott Fendorf,‡ John Bargar,§ and Craig S. Criddle*,† †

Department of Civil and Environmental Engineering, Stanford University, Stanford, California 94305, United States Department of Environmental Earth System Sciences, Stanford University, Stanford, California 94305, United States § Stanford Synchrotron Radiation Laboratory (SSRL), Stanford Linear Accelerator Center (SLAC) National Accelerator Laboratory, Stanford, California 94305, United States ‡

bS Supporting Information ABSTRACT: Soluble Fe(II) can reduce soluble U(VI) at rapid rates and in accordance with thermodynamic predictions. This was established by initially creating acidic aqueous solutions in which the sole oxidants were soluble U(VI) species and the sole reductants were soluble Fe(II) species. The pH of the solution was then increased by stepwise addition of OH
, thereby increasing the potential for electron transfer from Fe(II) to U(VI). For each new pH value resulting from addition of base, values of ΔG for the Fe(II)-mediated reduction of U(VI) were calculated using the computed distribution of U and Fe species and possible half reaction combinations. For initial conditions of pH 2.4 and a molar ratio of Fe(II) to U(VI) of 5:1 (1 mM Fe(II) and 0.2 mM U(VI)), ΔG for U(VI) reduction was greater than zero, and U(VI) reduction was not observed. When sufficient OH
was added to exceed the computed equilibrium pH of 5.4, ΔG for U(VI) reduction was negative and soluble Fe(II) species reacted with U(VI) in a molar ratio of ∼2:1. X-ray absorption near-edge structure (XANES) spectroscopy confirmed production of U(IV). A decrease in pH confirmed production of acidity as the reaction advanced. As solution pH decreased to the equilibrium value, the rate of reaction declined, stopping completely at the predicted equilibrium pH. Initiation of the reaction at a higher pH resulted in a higher final ratio of U(IV) to U(VI) at equilibrium.

’ INTRODUCTION Mining and refining of uranium for weapons and fuel have led to widespread groundwater contamination, and the challenge of protecting groundwater resources from uranium contamination is likely to grow; a 2008 forecast of global uranium production (existing, committed, planned, and prospective) projected a ∼40% increase between 2010 and 2020, from 86 720 tonnes/ year to 122 620 tonnes/year.1 Because uranium is both toxic and bioaccumulative,2 strategies are needed to decrease its bioavailability and to prevent its transport in water.3,4 In oxidizing environments, uranium exists primarily as hydroxyl and carbonate complexes of uranyl,5 either dissolved in solution or adsorbed to surfaces,6
8 and as precipitated mineral phases. In reducing environments, U exists as sparingly soluble forms, such as uraninite (UO2).9
11 Bioremediation of soluble U(VI) by reduction to sparingly soluble U(IV) was first proposed in the early 1990s12,13 and subsequently tested under field conditions.3,4,14
17 Bacteria that reduce U(VI) include dissimilatory iron-reducing bacteria (DIRB), such as Shewanella sp.,12 Geobacter spp.,16 Anaeromyxobacter spp.,18 various Clostridia,19 and sulfate-reducing bacteria (SRB), such as Desulfovibrio spp.20,21 and Desulfosporosinus spp.22 SRB and DIRB also mediate indirect uranium(VI) reduction via sulfide and Fe(II) species, respectively.11,17,23
25 U(VI) reduction by hydrogen sulfide r 2011 American Chemical Society

can explain the observed rates of U(VI) reduction for Desulfovibrio aerotolerans, a SRB isolated from a U(VI)-contaminated site.26 Many U-contaminated sediments are rich in iron,27,28 and iron can significantly impact U mobility.29 Ferrous iron solids known to reduce U(VI) include mixed Fe(II)
Fe(III) green rust,11 FeS minerals,25 siderite,17 and magnetite.30 X-ray photoelectron spectroscopy (XPS) studies have shown that FeS and magnetite reduce U(VI) to UO2, U3O8, and U4O9.25,31 Using extended X-ray absorption fine structure (EXAFS) analyses, Kelly et al. (2008) discovered a shift in the speciation and valence of sedimentassociated uranium before and after in situ bioreduction. In the bioreduced sediment, U(IV) was present as U(IV)
O
Fe, implicating Fe in the removal of U(VI)(aq).29 This observation was confirmed in static microcosms,31 where U(VI) species initially associated with C- and P-containing ligands were transformed to Feassociated U(IV) and uraninite. In the present study, we focus on the capacity for soluble Fe(II) species to reduce soluble U(VI) species. Solution chemistry controls the direction of these electron transfer reactions.5 At pH 7 and U(VI) and Fe(II) concentrations of 2  10
6 and Received: July 17, 2009 Accepted: April 14, 2011 Published: May 09, 2011 4718

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7 M, respectively, the Fe(III)/Fe(II) and U(VI)/U(IV) couples have similar reduction potentials. Small changes in solution chemistry reverse the direction of coupled reactions, switching iron species from U(VI) reductants to U(IV) oxidants.5 While high levels of carbonate favor formation of soluble uranyl carbonate complexes, a fact exploited for in situ leach mining of uranium, they also change the reduction potentials of Fe(III)/ Fe(II) and U(VI)/U(IV) couples, facilitating Fe(III) oxidation of U(IV). Studies that use high levels of high carbonate to extract U(VI) from a solid sample containing Fe(III) thus likely also measure U(IV). This is important for experiments designed to evaluate Fe(II) reduction of U(VI), where the reaction products are Fe(III) and U(IV) species. In the present work, we evaluated U(VI) reduction in two well-defined reaction systems, an iron
uranium (IU) system and an iron
uranium
carbonate (IUC) system, and we avoided use of carbonate extraction. Contrary to an earlier report,33 we find that the U
Fe system is under classical thermodynamic control with rapid reduction of soluble U(VI) by soluble Fe(II) species when thermodynamic conditions are favorable.

’ MATERIALS AND METHODS Stock Solutions. Stock solutions were prepared in 160 mL anaerobic serum bottles with oxygen-free water under a helium headspace, using analytical-grade chemicals or better: UO2Cl2, 10 mM; FeCl2, 20 mM; NaHCO3, 1 M; HCl, 0.5 M; NaOH, 0.2 M. All experiments were carried out in an anaerobic glovebox (Coy Laboratory Products Inc.) containing a He headspace. Iron(II) chloride tetrahydrate (99.99%) purity purchased from Signma-Aldrich was used for preparation of stock solutions. Uranyl chloride was obtained from Johnson Matthey, Ward Hill, MA. Uranium(VI) Reduction in the Iron
Uranium (IU) System. Serum bottles sealed with butyl rubber stoppers were prepared with 90 mL of helium-sparged oxygen-free deionized water under a helium headspace and stored inside an anaerobic glovebox containing a helium atmosphere. The following volumes of stock solutions were added: 2 mL of the UO2Cl2 stock, 5 mL of the FeCl2 stock, and 1 mL of the HCl stock. Samples were withdrawn by syringe for measurement of initial concentrations of Fe(II) and U (VI). The measured initial values were 1.0 mM for Fe(II) and 0.21 mM for U(VI). Initial solution pH was then increased in a stepwise manner by adding defined aliquots (0.02
0.5 mL) of NaOH stock. After each addition and mixing, a sample was extracted by syringe and transferred to a plastic screw-cap anaerobic centrifuge tube and the pH measured. The pH range was 2.36
10.77. Duplicate samples were prepared at each pH level. At pH > 8, a green precipitate appeared and bottles were shaken well to obtain a well-mixed sample. After collection of 14 samples, each at progressively higher pH levels, samples were centrifuged for 4 min at 1400 rpm in a glovebox. Centrate was removed by pipet and transferred to glass tubes containing 10 mL of anaerobic acidifed deionized water (pH 2.1) and then assayed for Fe(II) and U(VI). One sample with an initial pH of 6.52 (5.45 after reaction) was stored for analysis by X-ray adsorption near-edge structure (XANES) spectroscopy. Solids from another sample at the highest pH level (10.77), with no U(VI) in solution, were transferred to a plastic screw-cap anaerobic centrifuge tube, sealed, and stored inside a second anaerobic bottle for XANES analysis. Uranium(VI) Reduction in the Iron
Uranium
Carbonate (IUC) System. The same procedure described above was employed

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for the IUC system except a syringe was used to add 0.5 mL of NaHCO3 stock solution to the initial mixture. The pH range was 2.42
10.73. All samples were prepared in duplicate at each pH level. One sample with an initial pH of 6.39 (6.28 after reaction) was stored for analysis by XANES spectroscopy. As before, when a green precipitate appeared (pH > 6), the serum bottles were shaken well to ensure a well-mixed sample. After collection of 19 samples, each at a progressively higher pH level, all samples were centrifuged for 4 min at 1400 rpm in the glovebox. Centrate was transferred by pipet to glass tubes containing 10 mL of anaerobic acidifed deionized water (pH 2.1) for measurement of Fe(II) and U(VI). After measurement of aqueous U(VI) concentrations at each pH level, solids from the sample at pH 10.73 (no U(VI) in solution) were transferred to a plastic screw-cap anaerobic centrifuge tube, sealed, and stored inside a second anaerobic bottle for XANES spectroscopic analysis. Kinetic Measurements of Uranium Reduction. To assess the kinetics of U(VI) reduction, solution pH was adjusted from 2.40 to 4.10, 6.19, and 8.79 in the IU system and from 2.40 to 4.03, 6.22, and 8.93 in the IUC system. For both systems, the initial pH of ∼2.4 resulted from addition of 1 mM ferrous iron and 0.2 mM U(VI). The IUC system also contained 5 mM NaHCO3. Samples for U(VI) measurement were prepared as described above. Analytical Methods. Soluble U(VI) was assayed by spectrofluorometry after correcting for Fe(II) quenching.34 Centrate from centrifuged samples was transferred to glass tubes containing 10 mL of anaerobic acidifed deionized water (pH 2.1) and assayed for Fe(II). U(VI) standard curves were prepared at the same pH and levels of Fe(II) measured in the centrate. Sample unknowns and standards were transferred to cuvettes, supplemented with Uraplex (Chemcheck Instruments, Richland, WA), and assayed for U(VI) on a spectrofluorometer (Jobin Yvon, Inc., Edison, NJ) at wavelengths of 515.4 nm in emission acquisition mode and 280.0 nm in excitation acquisition mode. The U(VI) detection limit was 0.01 mg/L. Method accuracy was confirmed with a TJA IRIS Advantage/1000 Radial ICAP inductively coupled plasma optical emission spectrometer (ICP-OES) equipped with a solid state CID detector. The ICP-OES uses three wavelengths to measure fluorescence, enabling compensation for quenching, but was less sensitive (110 ppb vs 0; see Supporting Information). Other experiments described in ref 33 were subject to analytical problems due to fluorescence quenching by iron and/or sampling protocols that would not have detected a rapid decrease in U(VI). The present study avoided such concerns. Our results appear consistent with evidence of Fe-associated U(IV) sequestration within static microcosms31 the reported formation of Fe-associated U(IV)

Figure 3. Kinetics of uranium(VI) reduction in the IU and IUC systems: (A) Changes in U(VI) concentration in the IU system as a function of pH, for initial pH of values of 4.10 (line I), 6.19 (line II), and 8.79 (line III); (B) changes in measured pH in the IU system over time for an initial pH of 6.19; (C) changes in ΔG for different U(VI) species in the IU system over a pH range of 5.2
6.2, with initial concentrations of 0.2 mM for U(VI) and 1 mM for Fe(II); (D) changes in U(VI) concentration in the IUC system over time for initial pH values of 4.03 (line IV), 6.22 (line V), and 8.93 (line VI). 4723

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3. Review of Reference 33: Liger E.; Charlet L.; Cappellen P. V. Surface catalysis of uranium (VI) reduction by iron (II). Geochim. Cosmochim. Acta 1999, 63, 2939-2955. This material is available free of charge via the Internet at http:// pubs.acs.org.

’ AUTHOR INFORMATION Corresponding Author

*Phone: 650-723-9032; fax: 650-725-3164; e-mail: criddle@ stanford.edu.

’ ACKNOWLEDGMENT This work was supported the U.S. Department of Energy Environmental Remediation Sciences Program under grant DOEAC05-00OR22725. Figure 4. Normalized U LIII-edge XANES spectra for IU and IUC samples (lines) and fits (dotted) listed in Table 3. Vertical line gives the position of the absorbance intensity maximum for U(IV).

Figure 5. Diagram of pe
pH in the IU system for Fe2þ/Fe(III) and U(IV)/U(VI). Red lines illustrate pe
pH relationships for Fe(OH)3/ Fe2þ couples at aqueous-phase Fe2þ concentrations of 100 mM (line I), 1 mM (line II), and 0.01 mM III (line III). Blue lines illustrate the pe
pH relationships for U(VI)/U(IV) couples at aqueous-phase U(VI) concentrations of 20 mM (line IV), 0.2 mM (line V), and 0.002 mM (line VI).

within a bioreduced region of the subsurface at a field site in Oak Ridge, TN,32 and open new avenues for U(VI) sequestration research.

’ ASSOCIATED CONTENT

bS

Supporting Information. The Supporting Information contains the following materials: 1. TABLE S1. Examples of half-reactions for uranium reduction and the corresponding Standard Gibbs Free Energy Change (IUC system) at pH= 4.5 7.0 and 9.0. 2. TABLE S2. Half-Reactions with FeCO3 as the electron donor (pH =6.0, 7.0, 9.5).

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