Inhibition of U(VI) Reduction by Synthetic and Natural Pyrite

Article pubs.acs.org/est

Inhibition of U(VI) Reduction by Synthetic and Natural Pyrite Zhuanwei Yang,† Mingliang Kang,*,†,‡ Bin Ma,† Jinglin Xie,† Fanrong Chen,‡ Laurent Charlet,§ and Chunli Liu*,† †

Beijing National Laboratory for Molecular Sciences, Fundamental Science Laboratory on Radiochemistry & Radiation Chemistry, College of Chemistry and Molecular Engineering, Peking University, Beijing, 100871, China ‡ CAS Key Laboratory of Mineralogy and Metallogeny, Guangzhou Institute of Geochemistry, Chinese Academy of Sciences, Guangzhou, 510640, China § Environmental Geochemistry Group, ISTerre, University of Grenoble I, 38041 Grenoble, France S Supporting Information *

ABSTRACT: Reductive precipitation is an effective method of attenuating the mobility of uranium (U) in subsurface environments. The reduction of U(VI) by synthetic and naturally occurring pyrite was investigated at pH 3.0−9.5. In contrast to thermodynamic calculations that were used to predict UO2(s) precipitation, a mixed U(IV) and U(VI) product (e.g., U3O8/U4O9/U3O7) was only observed at pH 6.21−8.63 and 4.52−4.83 for synthetic and natural pyrite, respectively. Under acidic conditions, the reduction of UO22+ by surface-associated Fe2+ may not be favored because the mineral surface is nearly neutral or not negative enough. At high pH, the sorption of negatively charged U(VI) species is not favored on the negatively charged mineral surface. Thus, the redox reaction is not favored. Trace elements generally contained within the natural pyrite structure can affect the reactivity of pyrite and lead to a different result between the natural and synthetic pyrite. Because UO2(s) is extremely redox-sensitive toward U(VI), the observed UO2+x(s) phase reduction product indicates a surface reaction that is largely controlled by reaction kinetics and pyrite surface chemistry. These factors may explain why most laboratory experiments have observed incomplete U(VI) reduction on Fe(II)-bearing minerals.

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Fe(III) hydroxides,8 Fe(II) sorbed on hematite9 or clay,10 magnetite,11,12 and iron sulfides.13−15 Pyrite (FeS2) is the Earth’s most widespread and abundant sulfide mineral and occurs in a range of geological environments, such as hydrothermal, sedimentary, and igneous settings.16 In addition, pyrite occurs in granite and claystone, which are widely considered as host rocks for potential nuclear waste repositories.17,18 Small amounts of pyrite (∼1 wt %) are responsible for the low redox potential of the geologic environments that are located where radioactive waste storage is planned.17,19 Many studies have investigated the reduction of U(VI) by pyrite.20−25 However, partially reduced species (e.g., U3O8 and U4O9) are often reported as reduced uranium end products, although UO2(s) is the thermodynamically most stable form of uranium in pyrite-containing systems. In previous studies,20−22,25 a considerable amount of residual unreduced U(VI) was still present in the supernatant when the solid was sampled for characterization. Therefore, incomplete reduction potentially resulted from the relatively low reactivity of the naturally occurring pyrite that was used in these studies. Similar discrepancies have been observed for the reduction of

INTRODUCTION Mining and refinement of uranium (U) for weapons and fuel has led to widespread environmental contamination.1−3 Because uranium is toxic and bioaccumulative,4 strategies are needed that decrease its bioavailability and prevent its transport in water. In addition, it is important to understand the environmental behavior of uranium from aspects of the geologic disposal and storage of high-level radioactive wastes and spent fuel because the oxidation of UO2(s), the matrix material of nuclear fuel, will cause leaching of fission products and transuranium actinides. Uranium is redox-sensitive and can occur in several oxidation states. The most stable uranium oxidation states are U(IV) and U(VI), which are the most stable and prevalent species in the ambient environment. In reducing environments, uranium exists in insoluble forms, such as uraninite (UO2). Under oxidizing conditions, uranium exists as the uranyl cation (UO22+), which forms strong complexes with carbonate and other ligands in aqueous solutions and contributes to high uranium mobility.5,6 Therefore, uranium redox reactions are important for assessing the risks of a nuclear waste repository, remediating a site contaminated with uranium, forming uranium ore deposits, and global uranium cycling.7 Due to the reducing capacity of ferrous iron, numerous studies have investigated the abiotic sorption and reduction of uranyl on various Fe(II)-bearing minerals, such as Fe(II)/ © 2014 American Chemical Society

Received: Revised: Accepted: Published: 10716

May 6, August August August

2014 22, 2014 22, 2014 22, 2014

dx.doi.org/10.1021/es502181x | Environ. Sci. Technol. 2014, 48, 10716−10724

Environmental Science & Technology

Article

Table 1. Solution Data for U(VI) Reacted with Synthetic (I ∼ IV) and Natural (V ∼ IX) Pyritea ID

time d

pH

[U]tot (mM)

RSDb %

[Fe]tot (mM)

RSDb %

[S]tot (mM)

RSDb %

c

I

0 4 9 12 17

3.07 3.06 3.09 3.06 3.08

0.182 0.180 0.177 0.176 0.174

1.09 0.54 0.30 0.29 0.45

0.005 0.007 0.005 0.005 0.005

1.33 0.48 0.66 0.61 1.61

im imc imc imc imc

II

0 4 9 12 17

4.55 4.29 4.33 4.29 4.40

0.174 0.164 0.162 0.161 0.161

0.38 0.34 0.43 0.18 0.49

0.003 0.003 0.003 0.003 0.003

1.14 0.99 0.36 0.94 1.91

imc imc imc imc imc

III

0 4 9 12 17

8.52 6.31 ↓ 7.02d 6.21 ↓ 7.19d 6.71 7.16

0.001 0.002 0.002 0.002 0.001

11.45 2.26 2.40 1.57 1.69

0.002 0.002 0.004 0.003 0.002

0.67 0.84 0.49 0.54 1.96

imc imc imc imc imc

IV

0 4 9 12 17

9.48 8.93 8.52 8.49 8.59

0.001 0.007 0.014 0.014 0.014

6.69 0.76 0.51 0.55 0.69

0.002 0.002 0.002 0.002 0.002

1.29 2.21 1.56 1.50 1.96

imc imc imc imc imc

V

0 4 7 12 34 110

3.02 nde 3.05 nde 3.50 3.86

0.167 0.195 0.194 0.195 0.188 0.185

2.51 0.52 0.57 0.18 0.38 0.24

0.008 0.008 0.009 0.011 0.056 0.058

0.89 0.13 0.84 0.49 0.43 0.19

imc imc imc imc imc 0.018

12.72

0 4 7 12 34 110

4.52 nde 4.48 nde 4.74 4.83

0.165 0.186 0.188 0.182 0.171 0.015

6.72 0.77 0.36 0.56 0.97 0.46

0.005 0.006 0.006 0.007 0.031 0.029

10.99 0.30 0.92 0.68 0.49 0.64

imc imc imc imc 0.010 0.053

8.34 2.45

0 4 7 12 34 110

6.48 nde 6.42 nde 7.14 8.06

0.005 0.003 0.008 0.007 imc imc

56.32 11.00 7.56 2.37

0.003 0.002 0.003 0.003 0.006 0.004

18.97 0.55 0.73 1.43 2.93 0.84

imc imc imc imc imc 0.021

7.49

0 4 7 12 34

8.52 nde 5.77 nde 8.05

0.001 0.001 0.004 0.006 0.001

3.45 0.50 25.12 1.45 3.82

0.003 0.002 0.003 0.004 0.004

6.46 0.49 5.46 1.40 0.29

imc imc imc imc 0.058

1.68

0 4 7 12 34

9.49 nde 8.54 nde 8.60

0.002 0.007 0.013 0.017 0.017

11.65 30.99 45.92 23.64 13.44

0.003 0.005 0.003 0.004 0.002

1.80 17.45 9.22 19.05 1.44

imc imc imc imc 0.059

28.14

VI

VII

VIII

IX

For reactors I, II, V, and VI, 2 × 10−4 M UO2(NO3)2 was added as one dose. However, for reactors III, IV, VII, VIII, and IX, 2 × 10−5 M UO2(NO3)2 was added in 40 doses (i.e., 5 × 10−7 M each time) with a time interval of ∼2.0 h between each addition. In reactors I and II, ∼3.0 g·L−1 of synthetic pyrite was used, and ∼1.0 g·L−1 was used in reactors III and IV. In reactors V ∼ IX, 10.0 g·L−1 of natural pyrite was used. bRelative standard deviation. cBelow detection limit. dpH was adjusted using 0.1 M NaOH. eNot determined. a

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3.07−3.08, 4.55−4.40, 8.52−7.16, and 9.48−8.59 (labeled as I, II, III, and IV, respectively) for U(VI) reduction by synthetic pyrite. Before adding U(VI) for the reaction, the pyrite was further washed with 0.2 M HCl. A set of preliminary experiments was also performed (SI Text S1). For the reaction between U(VI) and natural pyrite, the reaction pH values were 3.02−3.86, 4.52−4.83, 6.48−8.06, 8.52−8.05, and 9.49−8.60 and were labeled as V, VI, VII, VIII, and IX, respectively. A 0.05 M HAc-NaAc solution was used for the reaction in reactors V and VI. Before the addition of U(VI), the pyrite powder was equilibrated with 50 mL of a 0.01 M NaCl solution for 48 h. To avoid the surface precipitation of schoepite (UO3·2H2O) (SI Figure S3), UO2(NO3)2 was added progressively under neutral to alkaline conditions. The experimental conditions are summarized in Table 1 and SI Table S1. All of the reactors were placed in a rotary tumbler at 30 rpm. The solution pH was regularly measured and allowed to vary (unless noted otherwise). A Mettler Toledo LE438 pH electrode connected to a FE20−FiveEasy Plus pH meter was used for the pH measurement. Three pH buffer solutions were used to calibrate the pH electrode. At defined time intervals, a 5 mL aliquot of suspension was sampled from each reactor and passed through a 0.22 μm pore size filter membrane (Millipore, Massachusetts). The total uranium, iron, and sulfur concentrations in the filtrate were analyzed by ICP-OES with a Leeman Laboratories Prodigy apparatus. Spectroscopic and Microscopic Analyses. X-ray photoelectron spectroscopy (XPS) was used to characterize the uranium speciation on the pyrite surface. A small amount of the reacted solid was collected from the reactor through a narrow tube and dried in the glovebox with silica gel. The dried samples were mounted on the sample holder in the glovebox and brought to the XPS facility in an O2-free glass bottle. To avoid any oxidation, the samples were rapidly transferred (99%. Anhydrous FeCl3 and NaHS·xH2O were purchased from the Alfa Aesar chemical company (Massachusetts). Deionized water (18.2 MΩ·cm) was boiled and purged with high-purity nitrogen gas while cooling before placing it in a glovebox. All of the other chemicals that were used in this study were analytical grade. Microsized pyrite was synthesized by reacting 0.05 M FeCl3 with 0.15 M NaHS following the method indicated in a recent study.32 An X-ray diffraction (XRD) analysis confirmed the product of FeS2, and no other phases were detected (SI Figures S1a,b). The specific surface area was 1.77 m2·g−1 according to the BET N2-absorption method. Commercial natural pyrite with a particle size of 0.06−0.19 in. was purchased from Alfa Aesar. These pyrite grains were ground to a fine powder with an agate mortar before transferring to the O2-free glovebox (O2 < 1 ppm). Next, the ground pyrite was washed several times with 0.2 M HCl, degassed water, and acetone before drying and storing under anoxic conditions. The specific surface area of the ground pyrite was 0.19 m2·g−1. The XRD characterization confirmed the product of pyrite, and no other phases were detected (SI Figure S1c). A scanning electron microscopy-energy dispersive spectrometer (SEM-EDS) analysis (JEOL 6301F) was performed on the sample, and it only detected sulfur and iron, with an average S/Fe ratio of 2.07. Inductively coupled plasma-optical emission spectrometry (ICP-OES) analysis after aqua regia dissolution indicated that the As (1.50), Ni (0.11), Cu (0.13), and Zn (0.09) impurities in the pyrite were generally under the detection limit (the values in parentheses are the detection limits in units of ppb). However, ∼0.44 wt % Co was detected (i.e., FeS2.07Co0.0089). The UO2 reference was prepared by using the hydrothermal method developed by Wang et al. (2008).33 To avoid any surface oxidation, the collected UO2 product was immediately transferred to the O2-free glovebox (O2 < 1 ppm) for washing and drying. UO3 and U3O8 references were prepared by heating UO2(NO3)2·6H2O to 550 and 1000 °C, respectively, in the air.34 The XRD analyses indicated that these synthetic U references were composed of the target products, and no other phases were detected (SI Figures S2a−c). Sorption Experiments. Sorption experiments were performed in an O2-free glovebox (O2 < 1 ppm) with a N2 atmosphere in 50 mL serum bottles. The solution pH was

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RESULTS AND DISCUSSION Aqueous Phase Analyses. The U(VI) concentrations did not show a significant decrease throughout the monitored reaction time at pH 3.06−4.55 (reactors I and II) when reacted with synthetic pyrite (Table 1). In addition, the concentrations of aqueous iron remained low and constant, and the concentration of dissolved sulfur remained below the detection limit. The slight U(VI) decrease may be ascribed to the physical adsorption or measurement error. Consequently, the solution results demonstrated that either the anticipated redox reaction

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Figure 1. Fitted Fe(2p), S(2p), and U(4f) XPS spectra for fresh and U(VI)-reacted synthetic pyrite. The U(4f) spectrum at the top right corner is for the U3O8 reference. The reaction time was 17 days.

between U(VI) and synthetic pyrite did not occur under acidic conditions, or the reaction was too slow to be detected. When pyrite was reacted at higher pH, aqueous U(VI) remained at ∼0.014 mM in reactor IV, whereas it remained as low as ∼0.001 mM in reactor III. Because 0.020 mM U(VI) was added in these reactors, reactor III may undergo reductive precipitation of U(VI). However, under alkaline conditions, precipitation of aqueous U(VI) as schoepite (UO3·2H2O) was spontaneous (See SI Figure S3).35 Similar results were also observed in the preliminary experiments (SI Text S1). Thus, the solution results alone could not identify the occurrence of a redox reaction without ambiguity. The solution results of U(VI) reacted with natural pyrite are shown in Table 1. Similar to the synthetic pyrite, no evidence for decreasing U(VI) concentrations occurred at pH < 4.0 in reactor V, although the reaction was investigated for over 110 days. In contrast, the U(VI) concentration decreased to 0.015 mM at pH 4.52−4.83 in reactor VI after 110 days of reaction. At the higher pH of reactors VII and VIII, the aqueous U(VI) and iron concentrations were extremely low throughout the reaction time, which implied that the redox reaction or the formation of schoepite controlled the U(VI) concentrations. Corresponding with the relatively high solubility of schoepite at pH > 8.6 (SI Figure S3), the aqueous U(VI) concentration remained at ∼0.017 mM in reactor IX, which suggested that no discernible redox reaction occurred. XPS Surface Features. XPS was used to characterize the speciation of uranium on the pyrite surface. The surface uranium was undetectable in the samples from reactors I, II, V, and IX, which corresponded to the constant U(VI) concentration in solution (Table 1). The XPS spectra of

U(4f), Fe(2p3/2), and S(2p) for the samples from reactors III, IV, VI, VII, and VIII are shown in Figures 1 and 2. The U(4f) spectra were fit with a spin−orbit splitting of 10.89 eV, a fixed intensity ratio of 4:3 (4f 7/2:4f5/2), and identical fwhm values (full width at half of maximum) for the doublet peaks. With respect to synthetic pyrite, the sample III could be fit with two components, the (4f 7/2) peak located at 380.2 and 381.5 eV that correspond to U(IV) and the unreduced U(VI) (SI Figure S4), respectively. These two contributions were consistent with the U3O8 reference, which is a mixed valent oxide of U(IV) and U(VI) that has two peaks at 380.3 and 381.5 eV (Figure 1), and indicated that the U(VI) was partially reduced. However, we could not exclude the possibility that the reaction products were composed of reduced forms together with schoepite precipitation. Sample IV was fit by only one contributor at 381.7 eV, which indicated unreduced U(VI). Thus, the slight decrease in the U(VI) concentration in the solution from reactor IV was ascribed to the spontaneous precipitation of schoepite rather than to the redox reaction. Similarly, the natural pyrite samples of VII and VIII were only fit with a component at 381.6 eV, which indicated the precipitation of schoepite (Figure 2). However, the precipitation of schoepite may partly account for the inhibition of U(VI) reduction. For sample VI, a minor peak at 379.6 eV that contributes to ∼7.6% of the total U was fit in addition to the main contributor at 380.9 eV. This peak indicated that the redox reaction occurred and formed a mixed valent oxide of U(IV) and U(VI) after 110 days of reaction. Compared with the fresh FeS2, the Fe(2p) spectra for the U(VI)-reacted synthetic pyrite had additional peaks (only contributing to ∼0.2% of the total Fe) at 711.3 and 711.3 eV 10719



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Figure 2. Fitted Fe(2p), S(2p), and U(4f) XPS spectra for the U(VI)-reacted natural pyrite. The reaction times were 110 days for sample VI and 34 days for samples VII and VIII.

Figure 3. SEM images of samples III (A) and VI (B). The lower panels show the corresponding EDS spectra of Spots 1 and 2 in A and Selected Area 1 in B (the Si signal was from the sample holder, and the K signal might have been from the residual KCl during pH measurement or introduced during solid sample preparation).

samples of VI, VII, and VIII, the Fe(2p) spectra did not show a new peak after contact with U(VI) (the Fe(2p) and S(2p) spectra for the fresh natural pyrite are shown in SI Figure S5). This result is understandable because U(VI) reduction

for samples III and IV, respectively. These peaks were consistent with an Fe(III)−O environment, which suggested the possible formation of small quantities of ferric oxyhydroxide under neutral to alkaline conditions. In the natural pyrite 10720



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Figure 4. HRTEM micrographs of fresh synthetic pyrite (A) and aggregated U on pyrite (B) (inset: EDS spectrum of the sample with backscattered peaks of S, Fe, U, Na, and Cl). The image of (C) shows distinctive lattice fringes, and (D) is the selected area electron diffraction (SAED) pattern from the area that is shown in 4B.

SEM and TEM Analyses for U(VI)-Reacted Pyrite. Representative SEM images for the U(VI)-reacted synthetic and natural pyrite are shown in Figure 3. The SEM images indicate that synthetic pyrite has a size of ∼2 μm and is stacked up with ∼100 nm cubic particles (SI Figure S6). The U-rich spots show layered or fibrous uranium deposited on synthetic pyrite particles for sample III (Figure 3A). Similar uranium precipitation was observed on the corroded steel coupons in a previous study.38 Uranium precipitation located on the grain edge was also confirmed in sample VI for natural pyrite (Figure 3B). The transmission electron microscopic results for sample III are given in Figure 4. The morphology and size of the fresh synthetic pyrite is illustrated in Figure 4A. Figure 4B depicts Urich particles with irregularly aggregated crystals are densely distributed on the external surface of the pyrite. As shown in Figure 4C, a single crystal is found growing with a d-spacing of 0.316 nm, and it occurs as a unique phase that is different from the pyrite. The selected area electron diffraction (SAED) pattern (Figure 4D) also indicates a strong d-spacing of ∼3.16

occurred under acidic conditions in reactor VI and no redox reactions occurred in reactors VII and VIII at higher pH values. The absence of an Fe(III)−O peak in the samples of VII and VIII also indicated that the inhibition of U(VI) reduction cannot be attributed to the formation of an Fe(III)−bearing passivation coating. To fit the S(2p) spectrum, doublet peaks with identical fwhm, a spin−orbit splitting of 1.18 eV, and intensity ratio of 2:1 (2p3/2:2p1/2) were used for each S species. All of the samples (including the fresh pyrite) were fit with three contributors that were located at 161.5 ± 0.2, 162.3 ± 0.2, and 164.3 ± 0.3 eV, which corresponded to surface S2−, bulk S22− dimers, and sulfur atoms of intermediate oxidation state (e.g., polysulfides and thiosulfate),21,36 respectively. The surface S2− resulted from the fracture of S−S bonds. In this case, monomeric sulfur species were produced at the surface (nominally S− species) and subsequently reduced to S2− during relaxation through the oxidation of surface Fe2+ ions to Fe3+.37 No peak was observed at 168−169 eV, which is the binding energy of SO42−. 10721



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this study contained ∼0.44 wt % Co, although the other elements (As, Ni, Cu, and Zn) occurred at concentrations that were less than the ICP-OES detection limit. Minor elements contained within the pyrite structure may result in different pyrite reactivities, which may explain why U(VI) reduction was observed at pH 4.52−4.83 (reactor VI) despite no discernible redox reactions occurring with synthetic pyrite (reactor II). At higher pH, UO2(OH)+, (UO2)3(OH)5+, and UO2(OH)3− became the dominant solution species. The formation of ferric hydroxide was kinetically favorable upon pyrite oxidation and could be represented by reactions 4, 5, and 6:

Å along with other spacings of ∼1.67 and ∼1.98 Å that match the values reported for U4O9 (ICCD PDF-2 card No. [01-0750944]), UO2 (No. [01-041-1422]), U3O8 (No. [00-047-1493]), and U3O7 (No. [00-042-1215]). These observations confirm the occurrence of U(VI) reduction, although we cannot determine the final product. Nevertheless, the formation of UO2 is unlikely because considerable amounts of unreduced U(VI) have been detected by XPS analysis. Indeed, UO2 is unstable and can be oxidized to a UO2+x phase by the addition of U(VI) or through the U(VI) that is released by schoepite (discussed below). Possible Reaction Inhibition Mechanisms. The results from the solution analyses and XPS surface characterization indicated that the reduction of U(VI) by pyrite did not occur under most conditions. The aqueous species of U(VI) were significantly affected by pH. SI Figure S7 shows the calculation for 10−6 M U(VI) and indicates that UO22+ was the predominant species at pH < 5.4, UO 2 (OH) + and (UO 2 ) 3 (OH) 5 + were dominant at pH 5.4−7.5, and UO2(OH)3− and uncharged UO2(OH)2 became dominant at pH 7.5−9.5. Under acidic conditions, the reaction between U(VI) and pyrite can be represented by the following equation:

UO2 (OH)+(aq) + 2Fe 2 +(aq) + 5H 2O(l) = UO2(cr) + 2Fe(OH)3(s) + 5H+(aq)

(UO2 )3 (OH)5+(aq) + 6Fe2 +(aq) + 13H 2O(l) = 3UO2(cr) + 6Fe(OH)3(s) + 13H+(aq)

= UO2(cr) + 2Fe(OH)3(s) + 3H+(aq)

UO2 2 +(aq) + 2Fe 2 +(aq) = UO2(cr) + (2)

UO2 (OH)+(aq) + 2Fe 2 +(aq) = UO2(cr) + 2Fe3 +(aq) + OH−(aq) Δr G° = 119.17kJ/mol

(6)

According to the calculated ΔrG (SI Table S2), the above reactions were thermodynamically feasible at pH 6.21−8.52 (reactor III) and were observed. Because pyrite is negatively charged at pH > 2.021,45 and the negatively charged UO2(OH)3− becomes the dominant species at pH > 8.0, electrostatic repulsion could prevent the adsorption of UO2(OH)3−. Thus, this electrostatic repulsion potentially explains why no redox reaction was observed at pH > 8.5 (reactors IV and IX). Overall, this is the first study to report the inhibition of U(VI) reduction on pyrite. However, it is difficult to provide conclusive explanations for the observations. Further investigations of the differences in the physicochemical characteristics between natural and synthetic pyrite and the inflection point for the reaction pH are needed along with a longer reaction with much longer time to unravel the inhibition mechanisms. Thermodynamic Considerations. Using the latest OECD/NEA (2003) thermodynamic data for U, 3 9 PHREEQC46 calculations indicated that UO2 was the most stable phase in the pyrite-containing systems for all of the experimental conditions (see also SI Figure S8). However, similar to previous research,20−25 incomplete reduction products (e.g., U3O8/U4O9/U3O7) were inferred from the TEM and XPS results on the synthetic and natural pyrite. These observations made us question these conflicting results. As indicated by Cui and Spahiu (2002),47 UO2(s) is strongly reactive toward U(VI) and reductive deposition of a UO2+x phase on the UO2 surface was observed in an aqueous U(VI)/ NaHCO3/pH 8.6 system under anaerobic conditions. Thus, in the presence of unreduced U(VI), UO2+x can be formed on pyrite rather than UO2(s) because the first nano-UO2(s) precipitation should be extremely reactive toward aqueous U(VI). Furthermore, in a slow reaction system (e.g., for natural pyrite), the colloidal UO2+x formed at the early stage was expected to aggregate and transform to bulk/compact UO2+x that has a low solubility. In these conditions, further transformation of the insoluble UO2+x phase to UO2(s) after depletion of the aqueous U(VI) was kinetically limited.

(1)

Based on the latest thermodynamic data for U39 and iron,40 the Eh-pH predominance diagram (SI Figure S8) indicates that UO22+ can be reduced to UO2(cr) by pyrite. Nevertheless, pyrite oxidation involves circular oxidation and reduction of the aqueous Fe2+ sorbed on the pyrite surfaces.41 Under acidic conditions, the first step for U(VI) reduction by pyrite can be represented by the following reactions:

2Fe3 +(aq) Δr G° = 69.22kJ/mol

(5)

UO2 (OH)3−(aq) + 2Fe 2 +(aq) + 3H 2O(l)

FeS2(s) + 7UO2 2 +(aq) + 8H 2O(l) = 7UO2(cr) + Fe2 +(aq) + 2SO4 2 −(aq) + 16H+(aq)

(4)

(3)

The Fe2+(aq) in reactions 2 and 3 refers to aqueous Fe2+ at the pyrite/water interface. Because of the positive ΔrG°, reactions 2 and 3 are thermodynamically feasible only in the presence of high concentrations of Fe 2+ , U(VI) species, and low concentration of Fe3+. Previous studies have indicated that the pHzpc (zero point of charge) of synthetic pyrite was pH 2.0, which is similar to that of the crushed acid-washed natural pyrite.42 Under acidic conditions, the sorption of Fe2+ and UO22+ cations at the pyrite/water interface due to electrostatic attraction was not favored. Thus, eqs 2 and 3 were not favored. In addition, the oxidation of Fe2+ by UO22+ or UO2(OH)+ and the reduction of Fe3+ by pyrite were expected to have slow reaction kinetics, which would limit the extent of the reaction in reactors I, II, and V at pH 3.02−4.55. In contrast, U(VI) reduction by natural pyrite was observed at pH 4.52−4.83 in reactor VI. The occurrence of U(VI) reduction on natural pyrite has been reported under acidic conditions (pH 3.6−6.1) over a short reaction time (e.g., 4 or 48 h).20,21,23 Natural pyrite generally holds minor elements (e.g., As, Co, Ni, etc.) that enhance its reactivity, particularly for As- and Co-doped pyrite.43,44 The natural pyrite sample used in 10722



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(4) Craft, E. S.; Abu-Qare, A. W.; Flaherty, M. M.; Garofolo, M. C.; Rincavage, H. L.; Abou-Donia, M. B. Depleted and natural uranium: Chemistry and toxicological effects. J. Toxicol. Environ. Health, Part B 2004, 7 (4), 297−317. (5) Bruno, J.; Ewing, R. C. Spent nuclear fuel. Elements 2006, 2 (6), 343−349. (6) Bargar, J. R.; Reitmeyer, R.; Davis, J. A. Spectroscopic confirmation of uranium(VI)-carbonato adsorption complexes on hematite. Environ. Sci. Technol. 1999, 33 (14), 2481−2484. (7) Behrends, T.; Van Cappellen, P. Competition between enzymatic and abiotic reduction of uranium(VI) under iron reducing conditions. Chem. Geol. 2005, 220 (3−4), 315−327. (8) 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 nanoparticies. Environ. Sci. Technol. 2003, 37 (4), 721−727. (9) Liger, E.; Charlet, L.; Van Cappellen, P. Surface catalysis of uranium(VI) reduction by iron(II). Geochim. Cosmochim. Acta 1999, 63 (19−20), 2939−2955. (10) Chakraborty, S.; Favre, F.; Banerjee, D.; Scheinost, A. C.; Mullet, M.; Ehrhardt, J. J.; Brendle, J.; Vidal, L.; Charlet, L. U(VI) sorption and reduction by Fe(II) sorbed on montmorillonite. Environ. Sci. Technol. 2010, 44 (10), 3779−3785. (11) Missana, T.; Maffiotte, U.; Garcia-Gutierrez, M. Surface reactions kinetics between nanocrystalline magnetite and uranyl. J. Colloid Interface Sci. 2003, 261 (1), 154−160. (12) Jeon, B. H.; Dempsey, B. A.; Burgos, W. D.; Barnett, M. O.; Roden, E. E. Chemical reduction of U(VI) by Fe(II) at the solid-water interface using natural and synthetic Fe(III) oxides. Environ. Sci. Technol. 2005, 39 (15), 5642−5649. (13) Veeramani, H.; Scheinost, A. C.; Monsegue, N.; Qafoku, N. P.; Kukkadapu, R.; Newville, M.; Lanzirotti, A.; Pruden, A.; Murayama, M.; Hochella, M. F. Abiotic reductive immobilization of U(VI) by biogenic mackinawite. Environ. Sci. Technol. 2013, 47 (5), 2361−2369. (14) Hyun, S. P.; Davis, J. A.; Sun, K.; Hayes, K. F. Uranium(VI) reduction by iron(II) monosulfide mackinawite. Environ. Sci. Technol. 2012, 46 (6), 3369−3376. (15) 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 absorption spectroscopy study. Environ. Sci. Technol. 2000, 34 (6), 1062−1068. (16) Vaughan, D. J.; Lennie, A. R. The iron sulfide minerals - their chemistry and role in nature. Sci. Prog. 1991, 75 (298), 371−388. (17) Baeyens, B.; et al. In situ physico-chemical characterisation of Boom Clay. In Radioactive Waste Management and the Nuclear Fuel Cycle; Harwood Academic Publishers: New York, 1985; Vol. 6, pp 391−408. (18) Gaucher, E.; Robelin, C.; Matray, J. M.; Negral, G.; Gros, Y.; Heitz, J. F.; Vinsot, A.; Rebours, H.; Cassagnabere, A.; Bouchet, A. ANDRA underground research laboratory: Interpretation of the mineralogical and geochemical data acquired in the CallovianOxfordian formation by investigative drilling. Phys. Chem. Earth. 2004, 29 (1), 55−77. (19) Beaucaire, C.; Pitsch, H.; Toulhoat, P.; Motellier, S.; Louvat, D. Regional fluid characterisation and modelling of water-rock equilibria in the Boom Clay formation and in the Rupelian aquifer at Mol, Belgium. Appl. Geochem. 2000, 15 (5), 667−686. (20) Aubriet, H.; Humbert, B.; Perdicakis, M. Interaction of U(VI) with pyrite, galena and their mixtures: A theoretical and multitechnique approach. Radiochim. Acta 2006, 94 (9−11), 657−663. (21) Descostes, M.; Schlegel, M. L.; Eglizaud, N.; Descamps, F.; Miserque, F.; Simoni, E. Uptake of uranium and trace elements in pyrite (FeS2) suspensions. Geochim. Cosmochim. Acta 2010, 74 (5), 1551−1562. (22) Bruggeman, C.; Maes, N. Uptake of uranium(VI) by pyrite under Boom Clay conditions: Influence of dissolved organic carbon. Environ. Sci. Technol. 2010, 44 (11), 4210−4216.

Additionally, the formation of schoepite is kinetically favorable under alkaline conditions, and because UO2 could be oxidized by the U(VI) released by schoepite, the formation of UO2 was inhibited as long as schoepite was present. These limitations may account for the incomplete reduction products that were observed in this study and in most previous experiments regarding U(VI) reduction by Fe(II)-bearing minerals.10,20−25,48 A similar mechanism was proposed for Se(IV) reduction by Fe(II)-bearing minerals.49 The findings in this study are directly linked to key geochemical processes that govern the mobility of toxic U in the environment. In a natural system that contains pyrite, whether UO2 controls U solubility is dependent on the absence of aqueous U(VI) or U(VI)associated phases.

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ASSOCIATED CONTENT

S Supporting Information *

Details regarding the XRD analyses of the synthetic FeS2, natural FeS2, and UO2, U3O8, and UO3 references are provided in the Supporting Information. In addition, details regarding the XPS spectra for the UO2(NO3)2, UO3, U3O8, and UO2 references, the fresh natural pyrite, the preliminary experiments for U(VI) reduction by synthetic pyrite, and the Eh-pH predominance diagram for U are available. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Authors

*(M.L.K.) Phone: +86 756 3668392; fax: +86 756 3668596; email: [email protected] or [email protected]. *(C.L.L.) Phone: +86 10 62765905; fax: +86 10 62765905; email: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS Funding for this research was provided by the Special Foundation for High-level radioactive Waste Disposal (2007840, 2012-851), the National Natural Science Foundation of China (NSFC, No. 11075006, 91026010), the China Postdoctoral Science Foundation project (Grant No. 2013M530013), and a collaborative project from the Key Laboratory of Mineralogy and Metallogeny at the Guangzhou Institute of Geochemistry in the Chinese Academy of Sciences (No. KLMM20120203). In addition, the authors are very grateful to Dr. Daqing Cui (Department of Materials and environmental Chem, Stockholm University, S-106 91, Sweden.) for his helpful comments and suggestions.

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REFERENCES

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Inhibition of U(VI) Reduction by Synthetic and Natural Pyrite

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