Environ. Sci. Technol. 2009, 43, 1237–1238
Response to Comment on “Reductive Immobilization of Uranium(VI) by Amorphous Iron Sulfide” We thank C. Noubactep for his interest in our recent paper (1) on the uptake and reduction of U(VI) by amorphous iron sulfide (FeS), in which we demonstrated that a rapid removal (within 1 h) of U(VI) by FeS from the aqueous phase was accompanied by a simultaneous release of Fe(II) due to ion exchange, and the reduction of U(VI) occurred at a much longer time scale than the initial removal. XPS analysis of the solid products further indicated the U(VI) reduction took place, forming U3O8/U4O9/UO2, polysulfide, and ferric iron as products. Noubactep voiced three concerns on our data interpretation (2); we thereby try to address his concerns one by one. The first comment is on the selection of experimental conditions that could potentially precipitate U(VI) as schoepite and whether U(VI) reduction is quantitative. Since high levels of uranium do exist in contaminated sites (e.g., up to 84-210 µM in the groundwater at the Field Research Center in Oak Ridge, TN (3)), we believe our use of 168.0 µM as the starting concentration is justified. We were aware of the potential of U(VI) precipitation in our experimental systems and pointed out that “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” (1). Our data, however, do not support precipitation being a major process. We have not observed U(VI) precipitation in the FeS-free control experiments, which is also consistent with other reports (4, 5). While surface nucleation could accelerate precipitation in general, attributing the removal of U(VI) in our system to surface precipitation is not appropriate because it cannot explain the simultaneous release of Fe(II) that accompanies U(VI) uptake at the constant pH. For example, near stoichiometric amounts of Fe(II) were released at pH 6.90 (Figure 1d in ref (1)). We, therefore, believe the uptake of U(VI) is through an ion exchange mechanism, at least at pH 5.99 and 6.90. In the alkaline solution, the amounts of Fe(II) released were less than the stoichiometric amounts of U(VI) removed (1); so we could not completely rule out the possibility of partial U(VI) precipitation of, as stated in the original paper (1). Nevertheless, we consider the formation of Fe(OH)2(s), which limits the solubility of Fe(II), is a more reasonable explanation for the low soluble Fe(II) because the experimental data agree well with the modeling results (Figure S3 in ref (1)). We are puzzled by Noubactep’s use of a paper by Read et al. (6) to support his claim of surface precipitation, because that paper did not report surface precipitation of U(VI) on “several” minerals. We also hope to reiterate that we have not observed or stated in the paper that U(VI) was quantitatively reduced to UO2, as implied by Noubactep (2). The XPS results showed the presence of reduced forms of uranium (UO2, U4O9, and U3O8), and the reductive immobilization occurred because these products are less mobile than the soluble U(VI) species. The second concern raised is on the use of 25 mM NaHCO3 for U extraction and speciation, although we are mystified by the logic behind this criticism (2). We measured the NaHCO3-extractable U in the presence of FeS and considered the extractable amount ([U(VI)]ex) represented the sum of U(VI) in the aqueous phase ([U(VI)]aq) and on the surfaces ([U(VI)]s). The rationales include: (i) carbonate ions form strong complexes with U(VI) (7, 8) and can completely desorb 10.1021/es803528t CCC: $40.75
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2009 American Chemical Society
U(VI) from ferrihydrite (9) and hematite (10) under alkaline conditions; (ii) no U(VI) uptake by FeS was observed in the presence of 25 mM NaHCO3 in our experiments; (iii) even though complete uptake of U(VI) by FeS occurred within 1 h (Figure 1a in ref (1)), significant amounts of U(VI) could be extracted after the initial uptake, suggesting that U(VI)(s) was extractable by the bicarbonate solution; and (iv) interference from the oxidative dissolution of U(IV) was unlikely, because carbonate did not dissolve UO2 (e.g., uraninite) even at concentrated concentrations (11), and the reoxidation of reduced U could not occur under the anoxic condition for extraction. We hope to emphasize that the conclusion of U(VI) reduction is not based on the initial uptake of U(VI) from the aqueous phase, but on the time-dependent NaHCO3 extraction results and more importantly, the XPS analysis of the reaction products. Noubactep’s assertions that amorphous iron oxides coprecipitated U(VI), which could not be dissolved by NaHCO3 so the sequestered U(VI) could not be extracted etc. (2), is not based on the facts relevant to the examined system. Naubactep’s third concern is on the mechanism of U(VI) reduction, but this might be caused by misunderstandings. In our system, we believe that reduced sulfur was the primary reductant, so when Fe(II) was released, the reduction of U(VI) associated with the surface could occur. Structural Fe(II) might also contribute to the reduction based on the XPS results (1, 7, 8, 10, 12). We do not believe that the aqueous Fe(II) was the reductant for U(VI) in our system and no Fe(III) was generated via oxidation of aqueous Fe(II), consistent with the literature reports (13, 14). Therefore, while Fe(OH)3 sequestration of U(VI) might be plausible under some other conditions, our observation does not support that mechanism in our systems. Furthermore, our reoxidation tests showed that upon exposure to oxygen that should also produce Fe(III)-(hydr)oxides, most uranium could still be extracted by carbonate solution. This suggests that when no U(VI) could be extracted by the 25 mM NaHCO3 solution under the anoxic condition, the sequestered U is likely in some reduced species formed through reduction, and Fe(III) sequestration could not be the dominant mechanism of U(VI) removal. Noubactep maintained that “...in the ideal case of dynamic equilibrium, when one atom of Fe(II) reacts with one atom of U(VI), one atom of Fe(II) is released from FeS”. Our best guess is that there must have been some typos here, because should Fe(II) be the reductant, U(VI) reduction to UO2 would require two Fe(II) ions, and no Fe(II) ions would be released. In summary, we maintain our original interpretations of the experimental data that demonstrated reductive immobilization of U(VI) in the presence of FeS under the anoxic condition. To evaluate the significance of this abiotic mechanism for U(VI) immobilization in the field, however, further research is needed to understand the impact of other geochemical parameters (e.g., CO2 and calcium).
Literature Cited (1) Hua, B.; Deng, B. Reductive Immobilization of Uranium(VI) by Amorphous Iron Sulfide. Environ. Sci. Technol. 2008, 42, 8703–8708. (2) Noubactep, C. Comment on “Reductive Immobilization of Uranium(VI) by Amorphous Iron Sulfide”. Environ. Sci. Technol. 2009, 43, 1236. (3) Wu, W.-M.; Carley, J.; Fienen, M.; Mehlhorn, T.; Lowe, K. B.; Nyman, J. A.; Luo, J.; Gentile, M. E.; Rajan, R.; Wagner, D.; Hickey, R. F.; Gu, B.; Watson, D.; Cirpka, O. A.; Kitanidis, P. K.; Jardine, P. M.; Criddle, C. S. Pilot-scale in situ bioremediation of uranium in a highly contaminated aquifer. 1. Conditioning VOL. 43, NO. 4, 2009 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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of a treatment zone. Environ. Sci. Technol. 2006, 40, 3978– 3985. 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.; Duff, M. C.; Gorby, Y. A.; Li, S. W.; Krupka, K. M. Reduction of U(VI) in goethite (R-FeOOH) suspensions by a dissimilatory metalreducing bacterium. Geochim. Cosmochim. Acta 2000, 64, 3085–3098. Read, D.; Lawless, T. A.; Sims, R. J.; Butter, K. R. Uranium migration through intact sandstone cores. J. Contam. Hydrol. 1993, 13, 277–289. 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. Hua, B.; Deng, B. Kinetics of uranium(VI) reduction by hydrogen sulfide in anoxic aqueous systems. Environ. Sci. Technol. 2006, 40, 4666–4671. 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. 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.
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(11) 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. (12) 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. (13) Liger, E.; Charlet, L.; Van Cappellen, P. Surface catalysis of uranium(VI) reduction by iron(II). Geochim. Cosochim. Acta 1999, 63, 2939–2955. (14) 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.
Bin Hua and Baolin Deng* Department of Civil and Environmental Engineering, University of Missouri, Columbia, Missouri 65211
* Corresponding author tel: (573) 882-0075; fax: (573) 882-4784; e-mail:
[email protected].
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