Effects of Mn(II) on UO2 Dissolution under Anoxic and Oxic Conditions

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Effects of Mn(II) on UO2 Dissolution under Anoxic and Oxic Conditions Zimeng Wang,*,† Bradley M. Tebo,‡ and Daniel E. Giammar† †

Department of Energy, Environmental and Chemical Engineering, Washington University in St. Louis, St. Louis, Missouri 63130, United States ‡ Institute of Environmental Health, Oregon Health & Science University, Portland, Oregon 97239, United States S Supporting Information *

ABSTRACT: Groundwater composition and coupled redox cycles can affect the long-term stability of U(IV) products from bioremediation. The effects of Mn(II), a redox active cation present at uranium-contaminated sites, on UO2 dissolution in both oxic and anoxic systems were investigated using batch and continuous-flow reactors. Under anoxic conditions Mn(II) inhibited UO2 dissolution, which was probably due to adsorption of Mn(II) and precipitation of MnCO3 that decreased exposure of U(IV) surface sites to oxidants. In contrast, Mn(II) promoted UO 2 dissolution under oxic conditions through Mn redox cycling. Oxidation of Mn(II) by O2 produced reactive Mn species, possibly short-lived Mn(III) in solution or at the surface, that oxidatively dissolved the UO2 more rapidly than could the O2 alone. At pH 8 the Mn cycling was such that there was no measurable accumulation of particulate Mn oxides. At pH 9 Mn oxides could be produced and accumulate, while they were continuously reduced by UO2, with Mn(II) returning to the aqueous phase. With the rapid turnover of Mn in the redox cycle, concentrations of Mn as low as 10 μM could maintain an enhanced UO2 dissolution rate. The presence of the siderophore desferrioxamine B (a strong Mn(III)-complexing ligand) effectively decoupled the redox interactions of uranium and manganese to suppress the promotional effect of Mn(II).

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the access of protons to the surface.12 A recent investigation13 found that UO2 dissolution at both oxic and anoxic conditions was inhibited by Ca(II) and Zn(II) at concentrations that are relevant to contaminated sites. X-ray absorption spectroscopy and electron microscopy suggested that a Ca-U(VI) phase and a Zn carbonate (hydrozincite, Zn5(CO3)2(OH)6), respectively, precipitated and coated the UO2 surfaces. As a redox-active divalent cation, Mn(II) may have more complex effects on the dissolution kinetics of UO2. Analogous to Ca(II) and Zn(II), adsorption and/or precipitation of Mn(II) at the UO2−solution interface may inhibit UO2 dissolution. On the other hand, UO2 oxidation by dissolved oxygen as well as Mn(III) and Mn(IV) species is energetically favorable, and Mn(II) oxidation to Mn(III) or Mn(IV) species may provide oxidants that are kinetically more reactive than molecular oxygen. 14−18 Even at low dissolved oxygen concentrations, oxidation of Mn(II) is thermodynamically favorable and can be catalyzed by Mn-oxidizing microorganisms19 and mineral surfaces.20,21 Active biological Mn oxidation can couple with the oxidation of U(IV)17 and

INTRODUCTION Uranium contamination of soil and groundwater is a legacy of past production and waste disposal activities associated with nuclear weapons development and is a continuing concern associated with the mining, milling, and processing of materials for the nuclear power industry. Microbial reduction of soluble U(VI) to U(IV) can produce biogenic UO2 as well as other low solubility U(IV) species as part of proposed bioremediation strategies.1−5 Biogeochemical and transport processes governing the longevity of reduced U(IV) species determine the effectiveness of uranium bioremediation.6−9 The chemical composition of groundwater can impact the kinetics of UO2 dissolution and reoxidation. Dissolved oxygen (DO) and dissolved inorganic carbon (DIC) promote the dissolution of UO2 by enhancing the oxidation of the UO2 surface and facilitating the detachment of labile U(VI) species from the UO2 surface.10,11 The impact of groundwater cations on UO2 dissolution has been less characterized. Non-redoxactive divalent cations inhibit UO2 dissolution through adsorption and surface precipitation reactions that block the access of oxidants to U(IV) sites. Santos et al. reported an inhibitory effect of Ca(II) on the anodic dissolution of UO2 and attributed it to the adsorption of Ca(II) on the UO2 surface that either suppressed the formation of a hydrolyzed U(VI) surface species that was the precursor to dissolution or blocked © 2014 American Chemical Society

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Cr(III)22−24 minerals to promote their oxidative dissolution. A recent study reported that abiotic heterogeneous Mn(II) oxidation on Cr(OH)3(s) surfaces resulted in substantially faster oxidative dissolution.25 Groundwater Mn(II) is at appreciable concentrations at uranium-contaminated sites;26,27 up to 20 μM soluble Mn has been reported at the Rifle former uranium mill tailings site.28 Seasonal fluctuations of the water table lead to oscillation of the redox status of the aquifer. The spring water table rise delivers dissolved oxygen to suboxic groundwater, while the dissolved oxygen concentrations decrease as the water table drops during the fall and winter. A remarkable inverse correlation of the dissolved Mn concentration with the dissolved oxygen concentration at the Rifle site indicated dynamic seasonal Mn redox cycling.8 The objective of this study was to evaluate the effects of abiotic processes involving Mn(II) on the kinetics of UO2 dissolution. This evaluation was pursued to gain insights into the mechanisms of U(IV) reactions with redox-active cations in subsurface environments. The central hypothesis was that Mn(II) would have different effects on UO2 dissolution under anoxic and oxic conditions. To test this hypothesis, dissolution experiments were performed at oxic and anoxic conditions with varied Mn(II) concentration and pH, and solids from these experiments were characterized.

Table 1. Summary of Batch and CSTR UO2 Dissolution Experiments batcha final Mnaqb (%)

UO2 dissolution rate (nmol g−1 s−1)

0.00 0.10 1.00 0.10 0.00 0.10 1.00 0.00 0.01 0.10 0.10 CSTRe

NA 96 86 95 NA 100 95 NA 34 16 97

0.0205c 0.0044c 0.0024c NA >0.09d >0.171d >0.096d >0.198d >1.016d >0.334d NA

expt

pH

1 2 3 4 5 6 7 8 9 10 11

8.0 8.0 8.0 8.0 8.0 8.0 8.0 9.0 9.0 9.0 8.0

DFOB and Mn(II) ≈ UO2 only. The promotional effect of Mn(II) on UO2 dissolution, which was identified in the absence of DFOB (Figure 1c), was not observed when DFOB was added together with Mn(II). DFOB is a strong ligand for Mn(III) that can promote the air oxidation of Mn(II) to form stable Mn(III) complexes, which keeps Mn(III) in solution (i.e., not on the surface or in solids) and prevents the further oxidation of Mn to form MnO2. The half-life of homogeneous Mn(II) oxidation in the presence of ambient oxygen and equimolar concentrations of DFOB and Mn(II) (0.5 mM) was reported to be 2.8 h for pH 8.1,30 suggesting that soluble DFOB-Mn(III) species could be produced in the time scale of the present study. Our recent study18 indicated that the Mn(III)−DFOB complex was not an effective oxidant for UO2 due to the strong hexadentate coordination that limits the ability of the complex to associate with the UO2 surface. Therefore, the role of DFOB in the present experiment was to capture and stabilize the Mn(III) 5551

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Figure 4. Effluent uranium (solid symbols, left axis) and manganese (empty symbols, right axis) concentrations from the CSTR experiments. (a−d) Results from expts 12−15. Results of duplicate experiments (triangles and circles) are shown. The lines indicate ideal CSTR behavior with constant reaction rates determined by the steady-state effluent U concentrations.

to a rate equation derived from CSTR experiments,16 10 μM chemogenic δ-MnO2 could increase the UO2 oxidation rate at pH 7.5 by only 0.02 nmol g−1 s−1, which was much less than the extent of promotion observed in the present study. On the other hand, a kinetic model predicted that 10 μM aqueous Mn(III) complexed by pyrophosphate (at a pyrophosphate:Mn(III) ratio of 6:1) could oxidize UO2 at pH 7.5 at a rate of 0.6 nmol g−1 s−1.18 Given that a lower ligandto-metal ratio is expected to increase the reactivity of aqueous Mn(III), this predicted rate may be a lower bound for an uncomplexed Mn(III) species that may be the dominant oxidant at the 10 μM initial Mn condition of the present study. While this explanation remains a hypothesis that would be difficult to test directly, the results are consistent with the speculation that short-lived Mn(III) species play a major role in Mn-mediated UO2 oxidation. Steady-State Dissolution: CSTR Experiments. The effects of Mn(II) on UO2 dissolution rates quantified in CSTR experiments under oxic conditions were consistent with those observed in the batch experiments. Under ambient oxic conditions at pH 7.7, the introduction of Mn(II) induced an immediate enhancement of UO2 dissolution (Figure 4a and b). The effluent uranium concentration responded differently after the introduction of 0.1 mM versus 1 mM Mn(II). Introduction of 0.1 mM Mn(II) led to a restabilized steady state with an effluent uranium concentration that was 40% higher than without Mn(II). In comparison, introduction of 1 mM Mn(II) induced an initial increase of uranium release, but the subsequent effluent uranium concentrations showed a consistent decease. These results were consistent with the batch experiments (Figure 1c), that increasing the Mn(II) concentration from 0.1 to 1 mM could not further enhance the dissolution rate of UO2. The subsequent downward trend in

species produced by O2, which prevented interactions of oxidized Mn species with UO2. UO2 dissolved faster in the presence of just 100 μM DFOB than in the presence of 100 μM DFOB and Mn. As a U(IV)complexing ligand, DFOB can promote UO2 dissolution.41 In the presence of an equimolar concentration of Mn(III) (from the air-oxidation catalyzed by DFOB), DFOB may be preferentially bound to aqueous Mn(III), which limits its ability to facilitate ligand-promoted dissolution of UO2. This could explain why UO2 dissolved faster in the presence of just 100 μM DFOB than in the presence of both DFOB and Mn. Possible Role of Mn(III). The preceding discussion as individual sections cannot completely resolve the question of whether the major promoter of UO2 oxidation is Mn oxides or other short-lived intermediates (e.g., Mn(III)). The DFOB experiments alone could not directly exclude the possibility that DFOB was limiting Mn(II) effects on UO2 oxidation by preventing the formation of Mn oxides and not by stabilizing Mn(III) in a form that did not react with UO2. However, on the basis of the collection of observations from multiple experiments a solid-associated or dissolved Mn(III) intermediate species appears to offer the most reasonable explanation of the present results. At pH 9, 100 μM Mn(II) with substantial production of Mn oxides led to slower UO2 oxidation than 10 μM Mn(II) (Figure 1e and f), suggesting that particulate Mn oxides were not the major promoter. Uncomplexed free Mn(III) cannot persist in solution due to its disproportionation. However, in the presence of UO2 as a reductant and a surface, short-lived Mn(III) intermediates at the surface or in solution can rapidly oxidize UO2 before disproportionation. Previous results provided estimates of the relative reactivity of preformed MnO2 and aqueous Mn(III) species. According 5552

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the effluent uranium concentration observed for 1 mM Mn(II) after the initially elevated effluent uranium concentration could be attributed to precipitation of MnCO3 (with a saturation index of 1.7 but slow kinetics), although the extent of precipitation was not detectable by SEM or XRD. Under a moderately oxic condition (PO2 = 0.01 bar), the introduction of 0.1 mM Mn(II) at pH 7.7 increased the UO2 dissolution rate by a factor of 2 (Figure 4c). A CSTR experiment conducted at 0.01 bar PO2 and pH 8.6 with 0.01 mM Mn(II) verified the ability of a low concentration of Mn(II) to sustain Mnmediated UO 2 oxidation even under low dissolved O 2 concentrations (Figure 4d). While the UO2 dissolution rates in the absence of Mn(II) under the two oxic conditions were lower in the present experiments than in previous studies with similar material,11 the differences were less than an order of magnitude. The approximately 5-fold difference of the dissolution rates between the experiments at 0.01 and 0.21 PO2 was comparable to what had been observed previously. Environmental Implications. Compared with groundwater cations that are not redox active (e.g., Ca and Zn),12,13 Mn can have more complicated effects on U(IV) stability. Temporal oscillation and spatial stratification of redox conditions are common in subsurface environments. This study highlighted the dual roles of Mn(II) under oxic and anoxic conditions. When the environment is strictly anoxic, Mn(II) could inhibit UO2 dissolution just as non-redox-active divalent cations do. Although dissolved oxygen that re-enters the previously reducing sediments serves as the ultimate electron acceptor and can directly oxidize UO2, the presence of Mn(II) can further enhance the rate of oxidative dissolution. Beyond their relevance to uranium, the present results may be applicable to other redox active metals that may be considered for reductive remediation (e.g., Cr, Np, Tc). While it had been established that Mn oxides can act as an oxidant for UO214−16 and that active microbial oxidation of Mn can significantly increase the dissolution rate of UO2,17 the present results demonstrated that Mn(II) could also abiotically promote the dissolution of UO2 as long as oxygen is available. With Mn behaving catalytically, the promotional effect can be sustained even without substantial accumulation of Mn oxides. Therefore, the role of Mn(II) is analogous to that of a catalyst; only a small amount of Mn(II) (e.g., 10 μM, which is relevant for contaminated sites8,28) was needed to sustain the redox coupling of uranium and manganese, which highlighted the previously underappreciated major influence of Mn in geochemical redox processes even with its minor fraction in terms of concentrations.42 The results also suggested that the UO2 dissolution rate as a function of Mn(II) concentration was nonlinear. Given the emerging evidence of the abundance and reactivity of Mn(III) species as intermediates of Mn redox cycling,18,43,44 geochemical models of Mn-mediated oxidation processes may not be adequate if they merely consider the oxidation by particulate Mn oxides as the final products of Mn oxidation.

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Article

AUTHOR INFORMATION

Corresponding Author

*Tel: 314-935-9437. E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors acknowledge helpful discussions with Sung-Woo Lee, Jeff Catalano, John Bargar, and Rizlan Bernier-Latmani. This research was supported by the U.S. Department of Energy, Office of Science, Subsurface Biogeochemical Research Program (DE-SC0005324). SEM was done in the Institute of Materials Science and Engineering at Washington University in St. Louis. Comments and suggestions of three anonymous reviewers and Associate Editor Stephan Hug helped us improve an earlier version of this paper.

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ASSOCIATED CONTENT

S Supporting Information *

Additional information, including one table (equilibrium constants for reactions involving dissolved Mn(II)) and two figures (adsorption isotherm of Mn(II) to UO2 and Mn(II) speciation). This material is available free of charge via the Internet at http://pubs.acs.org. 5553

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