Interactions of Selenate with Copper (I) Oxide Particles

The chemical mechanisms responsible for the immobilization of selenate (SeO42-) from aqueous solutions on cuprite (Cu2O) particles were determined fro...
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Langmuir 2004, 20, 6335-6343

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Interactions of Selenate with Copper(I) Oxide Particles A. Walcarius,* J. Devoy, and J. Bessie`re Laboratoire de Chimie Physique et Microbiologie pour l’Environnement, UMR 7564 CNRS Universite´ Henri Poincare´ Nancy I, 405, rue de Vandoeuvre, F-54600 Villers-le` s-Nancy, France Received February 13, 2004. In Final Form: May 10, 2004 The chemical mechanisms responsible for the immobilization of selenate (SeO42-) from aqueous solutions on cuprite (Cu2O) particles were determined from batch experiments. This was achieved by performing both solution-phase analyses and characterization of solid particles by X-ray photoelectron spectroscopy and transmission electron microscopy techniques, after equilibration of cuprite particles with selenatecontaining solutions at various pH values, solid-to-solution ratios, and ionic strengths. Two distinct mechanisms have been pointed out. In the acidic medium, where the acid-catalyzed dissolution of cuprite into CuI species occurs, the immobilization of selenate implies a redox reaction with transient CuI leading to the precipitation of copper(II) selenite, CuSeO3. In the absence of protons added in the medium, Cu2O is chemically stable and immobilization of SeO42- is essentially due to adsorption in the form of an outersphere surface complex. The uptake level of selenate by Cu2O is markedly lower than that observed for selenite species in the same conditions.

1. Introduction Mechanistic understanding of trace element cycling in the environment requires the knowledge of mass transport processes occurring across boundaries between surface waters and underlying soils or sediments. To this end, it appears necessary to investigate the (bio)chemical mechanisms that are responsible for the immobilizationleaching processes at such solid/liquid interfaces. Many efforts have been directed to characterize the biogeochemical cycle of selenium in the environment1-4 (see, e.g., investigation of sorption processes of selenite and selenate in soils),5-11 as well as to find solid scavengers likely to limit their migration in natural waters.12-17 Selenate is one of the most mobile forms (with selenite) of selenium in natural waters.4 Because of its oligo-element * To whom correspondence should be addressed. Fax: (+33) 3 83 27 54 44. E-mail: [email protected]. (1) Selenium in Agriculture and the Environment; Jacobs, L. W., Ed.; SSSA Special Publication No. 23; American Society of Agronomy: Madison, WI, 1989. (2) Masscheleyn, P. H.; Patrick, W. H., Jr. Environ. Toxicol. Chem. 1993, 12, 2235. (3) Selenium in the Environment; Frankenberger, W. T., Jr., Benson, S. M., Eds.; Marcel Dekker: New York, 1994. (4) Conde, J. E.; Sanz Alaejos, M. Chem. Rev. 1997, 97, 1979. (5) Alemi, M. H.; Goldhamer, D. A.; Nielsen, D. R. J. Environ. Qual. 1988, 17, 608. (6) Fio, J. L.; Fujii, R.; Deverel, S. J. Soil Sci. Soc. Am. J. 1991, 55, 1313. (7) Paya-Perez, A. B.; Goetz, L.; Springer, A.; Nielsen, A. J. Trace Microprobe Tech. 1993, 11, 143. (8) Johnsson, L. Ecol. Bull. 1995, 44, 123. (9) Tao, Z.; Du, J.; Dong, W.; Zheng, L. J. Radioanal. Nucl. Chem. 1996, 214, 245. (10) Pezzarossa, B.; Piccotino, D.; Petruzzelli, G. Commun. Soil Sci. Plant Anal. 1999, 30, 2669. (11) Sharmasarkar, S.; Vance, G. F. Adv. Environ. Res. 2002, 7, 87. (12) Koren, D. W.; Gould, W. D.; Lortie, L. Waste Process. Recycl. Min. Metall. Ind., Proc. Int. Symp. 1992, 171-182. (13) Kuan, W.-H.; Lo, S.-L.; Wang, M. K.; Lin, C.-F. Water Res. 1998, 32, 915. (14) Li, C.; Viraraghavan, T. Water Resources and the Urban Environment- 98, Proceedings of the National Conference on Environmental Engineering; Chicago, June 7-10, 1998; American Society of Civil Engineers: Reston, Va: 1998; pp 356-361. (15) Qiu, S. R.; Lai, H.-F.; Roberson, M. J.; Hunt, M. L.; Amrhein, C.; Giancarlo, L. C.; Flynn, G. W.; Yarmoff, J. A. Langmuir 2000, 16, 2230. (16) Qian, S.; Huang, G.; Jiang, J.; He, F.; Wang, Y. J. Appl. Polym. Sci. 2000, 77, 3216.

nature, selenium is both an essential nutrient and a potentially toxic trace element for which the concentration window between toxic and deficient levels is quite narrow.18,19 As a result of the usually low interaction of selenate with most soil components (especially clay minerals),20-22 their efficient immobilization would require the use of sorbent additives to reduce their concentration in polluted waters to nontoxic levels. Also, the 79Se isotope is a long-lived (1.1 × 106 years)23 fission product which is chemically and radiologically toxic. Under experimental conditions typical of oxidative alteration of spent nuclear fuel, selenite, SeO32- or HSeO3-, or selenate, SeO42-, are the dominant aqueous species of selenium.24 Because migration of these mobile long-lived radionuclides is not retarded by silicate minerals, it is extremely important to find solids able to trap them in the hypothesis of radioactive waste disposal in deep geological sites (e.g., in the design of barriers surrounding the waste containers).25-27 Sorption of selenate on metal oxides is well-documented in the cases of iron oxides and oxi-hydroxides27-38 and, to (17) You, Y.; Vance, G. F.; Zhao, H. Appl. Clay Sci. 2001, 20, 13. (18) Levender, O. A. Fed. Proc. 1985, 44, 2579. (19) Lockitch, G. Crit. Rev. Clin. Lab. Sci. 1989, 27, 483. (20) Ahlrichs, J. S.; Hossner, L. R. J. Environ. Qual. 1987, 16, 95. (21) Bar-Yosef, B.; Meek, D. Soil Sci. 1987, 144, 11. (22) Neal, R. H.; Sposito, G. S. Soil Sci. Soc. Am. J. 1989, 53, 70. (23) Jiang, S.; Duo, J.; Jiang, Sh.; Li, C.; Cui, A.; He, M.; Wu, S.; Li, S. Nucl. Instrum. Methods 1997, B123, 405. (24) Chen, F.; Burns, P. C.; Ewing, R. C. J. Nucl. Mater. 1999, 275, 81. (25) Fujikawa, Y.; Fukui, M. Radiochim. Acta 1997, 76, 163. (26) Wang, P.; Anderko, A.; Turner, D. R. Ind. Eng. Chem. Res. 2001, 40, 4444. (27) Duc, M.; Lefe`vre, G.; Fedoroff, M.; Jeanjean, J.; Rouchaud, J. C.; Monteil-Rivera, F.; Dumonceau, J.; Milonjic, S. J. Environ. Radioact. 2003, 70, 61. (28) Hayes, K. F.; Roe, L.; Brown, G. E., Jr.; Hodgson, K. O.; Leckie, J. O.; Parks, G. A. Science 1987, 238, 783. (29) Balistrieri, L. S.; Chao, T. T. Soil Sci. Soc. Am. J. 1987, 51, 1145. (30) Hayes, K. F.; Papelis, C.; Leckie, J. O. J. Colloid Interface Sci. 1988, 125, 717. (31) Balistrieri, L. S.; Chao, T. T. Geochim. Cosmochim. Acta 1990, 54, 739. (32) Zhang, P.; Sparks, D. L. Environ. Sci. Technol. 1990, 24, 1848. (33) Manceau, A.; Charlet, L. J. Colloid Interface Sci. 1994, 68, 87. (34) Wijnja, H.; Schulthess, C. P. J. Colloid Interface Sci. 2000, 229, 286. (35) Su, C.; Suarez, D. L. Soil Sci. Soc. Am. J. 2000, 64, 101.

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a lesser extent, of aluminum oxides.34,39-42 Interactions between selenate and these metal oxide surfaces involve most often adsorption via the formation of an outer-sphere complex (a mixture of outer- and inner-sphere surface complexes on ferric oxides was also recently reported).34,36,37 Selenate adsorbs much more weakly than selenite,29,31,32 which always forms an inner-sphere surface complex on metal oxides,13,28-33,39,43-45 and the selenate adsorption process is markedly reduced by increasing the ionic strength of the external solution28-30,39 or by the presence of competing anions.31,38-41 Sorption processes are strongly pH-dependent, increasing by shifting from a neutral to an acidic medium.27-42 To the best of our knowledge, no information is available on the interactions between selenate and copper(I) oxide, cuprite, Cu2O. The only work dealing with selenate and copper oxide-based minerals relates on the preconcentration of this analyte, prior to its detection by atomic absorption spectrometry (AAS),46 on copper(II) oxide, CuO. In a previous work, we have investigated the uptake of selenite species (SeIV) from an aqueous medium by Cu2O particles.45 Three immobilization mechanisms were pointed out as a function of pH and reaction time: bulk precipitation of crystalline CuSeO3‚2H2O in an acidic medium and fast chemisorption of SeO32- on Cu2O at a pH above 7, which was followed by phase transformation of the cuprite surface at longer times concomitantly to the formation of amorphous CuSeO3. In the present paper, we have studied the chemical mechanisms that are responsible for the immobilization of selenate (SeVI species) on cuprite (Cu2O) from an aqueous medium, with the objective to point out the type of interactions occurring at the solid/liquid interface. Investigations were carried out from batch equilibration experiments performed in aqueous suspensions of Cu2O (at various solid-to-solution ratios) adjusted to appropriate pH values, which contain selenate at selected concentrations. Experimental data were obtained from solutionphase analyses and solid characterization using various bulk solid and surface analysis methods. The effect of the ionic strength was also investigated. 2. Experimental Section 2.1. Sorbents, Chemicals, and Solutions. Commercially available cuprous oxide, Cu2O, under the form of dry powder (97%, Fluka) was mainly used in this study. The visibly hydrophobic particles were pretreated in a slightly acidic medium (∼10-4 M HClO4) during 12 h to remove the oxidizing layer that was present on their surface.47 The solid particles were then filtered, copiously washed with pure water, and dried in the absence of oxygen. After such treatment, the Cu2O particles turned hydrophilic and were easily dispersed in the aqueous medium. Granulometry analysis of this solid revealed an average particle size of 3.9 µm (d20 ) 2.5 µm, d80 ) 5.1 µm). Some experiments involving transmission electron microscopy (TEM) (36) Rietra, R. P. J. J.; Hiemstra, T.; van Riemsdijk, W. H. J. Colloid Interface Sci. 2001, 240, 384. (37) Peak, D.; Sparks, D. L. Environ. Sci. Technol. 2002, 36, 1460. (38) Wijnja, H.; Schulthess, C. P. Soil Sci. Soc. Am. J. 2002, 66, 1190. (39) Ghosh, M. M.; Cox, C. D.; Yuan-Pan, J. R. Environ. Prog. 1994, 13, 79. (40) Wu, C.-H.; Lin, C.-F.; Lo, S.-L. Colloids Surf., A 2000, 166, 251. (41) Wijnja, H.; Schulthess, C. P. Soil Sci. Soc. Am. J. 2000, 64, 898. (42) Schulthess, C. P.; Hu, Z. Soil Sci. Soc. Am. J. 2001, 65, 710. (43) Papelis, C.; Brown, G. E., Jr.; Parks, G. A.; Leckie, J. O. Langmuir 1995, 11, 2041. (44) Parida, K. M.; Gorai, B.; Das, N. N.; Rao, S. B. J. Colloid Interface Sci. 1997, 185, 355. (45) Devoy, J.; Walcarius, A.; Bessie`re, J. Langmuir 2002, 18, 8472. (46) Reddy, K. J.; Zhang, Z.; Blaylock, M. J.; Vance, G. F. Environ. Sci. Technol. 1995, 29, 1754. (47) Lefe`vre, G.; Walcarius, A.; Ehrhardt, J.-J.; Bessie`re, J. Langmuir 2000, 16, 4519.

Walcarius et al. analysis have required the use of Cu2O particles displaying smaller size (85% completion within the first 24 h, depending on the mechanism involved, see the following), which is much faster than for selenite species on the same Cu2O sorbent for which 2 weeks were necessary to get stationary values.45 The corresponding data were gathered and presented by plotting the equilibrium values of immobilized SeVI (calculated by the difference between the starting and equilibrium solution-phase concentrations) as a function of final pH. These curves are displayed in Figure 1 for the three starting selenate concentrations, 2 × 10-3 M (A), 4 × 10-3 M (B), and 8 × 10-3 M (C), and a solid-to-solution ratio of 30 g L-1. The general trend that can be observed at first glance is a rather low uptake efficiency at high pH values and much higher (even total) immobilization in the presence of protons, with a sharp breakthrough in the 7-8.5 pH region (Figure 1). This resembles at first glance the typical adsorption behavior generally reported for anions on oxidebased minerals (because positively charged surfaces are expected to interact strongly with negatively charged solutes),25,51,52 but the amount of immobilized SeVI is much greater than that corresponding to a monolayer adsorption (especially at pH below 8), suggesting the participation of (at least) one other mechanism to explain the whole uptake process, as previously reported for SeIV on Cu2O.45 These curves exhibited some other features, suggesting the participation of distinct sorption mechanisms. First, the pH window corresponding to total SeVI consumption from the solution became more and more narrow when raising the starting SeVI concentration. For example, at pH 7, the fraction of SeVI consumed decreased from about 90 to 80 to 70% for selenate concentrations varying from 2 to 4 to 8 mM, respectively; note that this corresponds to a net increase in the absolute quantity of immobilized species (e.g., 90% of 2 mM ) 1.8 mM and 70% of 8 mM ) 5.6 mM). This agrees to the fact that a high concentration (50) Wagner, C. D.; Bickham, D. M. NIST Standard Reference Database 2.0, NIST XPS Database 1.0; U.S. Department of Commerce, National Institute of Standards and Technology: Gaithersburg, MD, 1989. (51) Goldberg, S.; Glaubig, R. A. Soil Sci. Soc. Am. J. 1988, 52, 954. (52) Szczepaniak, W.; Koscielna, H. Anal. Chim. Acta 2002, 470, 263.

Figure 1. Variation of immobilized SeVI species from 25 mL of solutions containing initially 0.75 g of Cu2O and (A) 2 × 10-3, (B) 4 × 10-3, and (C) 8 × 10-3 M SeO42- (adjusted at various starting pHs, with added H+ concentrations ranging from 0 to 0.2 M or OH- in the range 0-0.01 M), as a function of the final pH measured at equilibrium.

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Table 1. Comparison between Starting and Equilibrium pH Values Measured in Cu2O Suspensions Contacting Three SeO42- Concentrations in an Acidic Medium, Given for Some Representative Experimental Data Points of the Curves Depicted in Figure 1 [SeO42-] ) 2 × 10-3 M

[SeO42-] ) 4 × 10-3 M

[SeO42-] ) 8 × 10-3 M

initial pH final pH initial pH final pH initial pH final pH 3.3 4.5 5.0 6.0

4.8 5.9 6.7 7.6

3.5 5.3 5.5 5.9

5.6 7.3 7.6 7.9

3.2 3.9 5.3 5.5

4.5 5.0 7.5 8.2

in the equilibrium solution will always support a higher uptake amount at a given pH, independently of the mechanism. It is also noteworthy that the equilibrium pH of the suspensions was significantly different from its initial value measured directly after addition of strong acid or base in the medium, and this was especially dramatic when attempting to get acidic media (Table 1). As discussed earlier45,47 and in the following, the dissolution of cuprite in the presence of protons is the main responsible pathway explaining this phenomenon. Finally, the addition of too large amounts of protons resulted in some restriction of the immobilization extent, and this effect was much more marked with higher SeVI concentrations (Figure 1C). These two last features contribute to explain the rather large dispersion of data points, which expresses a non-negligible dependence of the uptake process(es) on small variations in the experimental conditions (e.g., small differences in the amounts of added acid or base). Working with unbuffered solutions and uncontrolled ionic strength was preferred here to avoid any side effect of their components on the immobilization processes. In the alkaline medium, the SeVI uptake efficiency was much less and the process was found to be dependent on the selenate concentration and solid-tosolution ratio, but the rather high dispersion of data points prevents careful analysis at this stage (as a result of the intrinsic lack of accuracy in quantifying low SeVI consumption, the measured equilibrium concentrations being very close to the starting ones). All these preliminary observations suggest the intervention of more than one immobilization mechanism to describe the uptake of SeVI by Cu2O. 3.2. Sorption in an Alkaline Medium. Consistent with anion adsorption on metal oxides or oxyhydroxides,27-42 selenate is thought to adsorb on cuprite, at least in the pH range where this copper oxide does not undergo significant dissolution (>7). Indeed, the cuprite mineral surface contains hydroxyl groups characterized by acidbase properties leading to a point of zero change in the pH range between 7 and 9.553 (no accurate data on pKa are available probably because of the well-known dissolution of cuprite in the presence of protons, this process preventing to a great extent the protonation of the surface hydroxyl groups). The adsorption behavior is also expected from the global envelopes depicted in Figure 1 for which a nearly constant amount of SeVI species seems to be immobilized, that is, at pH 10 (this can be evaluated by multiplying the fraction of selenate consumed by the starting SeVI concentration: ≈8% of 2 mM (Figure 1A) ) ≈4% of 4 mM (Figure 1B) ) ≈2% of 8 mM (Figure 1C). The adsorption mechanism is sustained by the rather fast kinetics associated to the uptake process, indicating that 50% of the maximal immobilization level was achieved during the first 5 min, 90% within 24 h, and the equilibrium steady state after 4 days (Figure 2a). This behavior agrees well with adsorption of selenite on the

Figure 2. Evolution of the amount of (a) selenate and (b) selenite species bound to cuprite as a function of reaction time. The suspensions were prepared to contain 2.0 × 10-3 M SeVI or SeIV and 30 g L-1 Cu2O, without any added pH-adjustment reagent.

same solid, which occurred in the early stage of reaction between SeIV and Cu2O,45 but suggests the only existence of this adsorption process because no further change in equilibrium concentrations were observed after 4 days for the next 2 weeks. This is a significant difference compared to the SeIV-Cu2O system for which a second, much slower uptake process leading to the immobilization of larger amounts of selenium (Figure 2b, data from our previous work45 for comparison purposes) was reported, involving phase transformation and surface precipitation in addition to surface complexation.45 Fast adsorption of selenate of other surface hydroxyl-bearing solids was also reported elsewhere.54 Further evidence to support the adsorption mechanism was provided by measuring the variation of selenate uptake as a function of Cu2O content in the suspension, indicating a direct proportionality between the amount of immobilized SeVI and the cuprite density in the medium independently of the equilibration time (Figure 3). This linear variation was obtained both after 5 min of reaction (corresponding to half of the maximum loading capacity) and at the equilibrium, over the whole Cu2O content range between 10 and 80 g L-1, as illustrated by parts a and b in Figure 3. This is consistent with the fact that the amount of adsorbed species is directly related to the total surface area of Cu2O available for SeVI binding. For each reaction time considered here, the adsorption isotherms (inset of Figure 3) were nearly flat and confirmed that the experiments were performed with the added SeVI in excess of the binding capacity of the adsorbent. The calculated distribution coefficients (KD) were constant for all the solid-to-solution ratios considered here, at values around 6 L kg-1, which is very close to what was reported for SeIV adsorption on the same solid in similar experimental conditions.45 Note that much higher KD values can be reached when working with Cu2O contents in excess to the amount of SeVI in solution (i.e., in conditions more relevant to remediation processes): for example, a KD value of 40 L kg-1 was calculated after the treatment of 25 mL of solution of 10-5 M SeVI in the presence of 0.25 g of Cu2O, in agreement with the generally observed increase of KD values by decreasing the analyteto-sorbent ratio.29,44 (53) Raju, G. B.; Forsling, W. Bull. Electrochem. 1992, 8, 402. (54) Scott, M. J.; Morgan, J. J. Environ. Sci. Technol. 1996, 30, 1990.

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Figure 3. Variation of the amount of selenate bound to cuprite as a function of adsorbent concentration. Data points were recorded after (a) 5 min of reaction and (b) at equilibrium (4 days). The suspensions were prepared to contain 2.0 × 10-3 M SeVI and various amounts of Cu2O in 25 mL of solutions (final pH values were around 9).

Figure 4. Adsorption of SeVI species on Cu2O obtained (a) in the absence of added electrolyte, (b) in the presence of 0.1 M NaNO3, and (c) with 1 M NaNO3, as a function of the final pH measured after 24 h of equilibration. Experimental conditions: 1.0 × 10-5 M SeVI, 10 g L-1 Cu2O, 25 mL of solution, pH adjusted with diluted HClO4 or NaOH.

The maximal amount of adsorbed selenate on cuprite was equal to 12 ( 3 µmol g-1 (as estimated from the isotherm in the inset of Figure 3). Taking into account the specific surface area of the Cu2O sample (1.6 ( 0.5 m2 g-1, as measured on 4 aliquots of cuprite after acid washing; this agrees with values previously published)54,55 and knowing the ionic radius of SeO42- (0.256 nm),56 it is possible to calculate a surface coverage equal to 7.5 ( 2 µmol m-2, corresponding to 4.5 ( 1.5 molecules/nm2, which is close to, yet lower than, the monolayer coverage (ca. 4.8 molecules/nm2, as estimated on the basis of nonhydrated SeO42- hard spheres). This value is comparable but usually higher than those reported for selenate adsorption on aluminum oxide or iron oxyhydroxides.37,40,57 All the above adsorption experiments (Figures 2 and 3) have been performed in excess SeVI over the binding capacity of the adsorbents. For remediation purposes, however, the solid-to-solution ratio is often much higher, to reduce the solution-phase selenium concentration to insignificant levels, so that we have studied the adsorption of SeVI species from more diluted solutions. Adsorption was found to be strongly affected by pH: it resulted in low efficiency in highly a alkaline medium and increasing adsorption yields by decreasing pH (in agreement with the increasingly growing positive charge on the cuprite surface), the corresponding isotherm displaying a sigmoidal shape characterized by a half-maximum capacity at a pH value of about 10.5 (Figure 4a). This value is significantly higher than that reported for selenate adsorption on copper(II) oxide, CuO,46 for which only 50% SeVI binding was obtained at pH 6 and only 20% at pH 8 in the same conditions as those applied here to get the experimental data depicted in Figure 4. Indeed, more than 50% SeVI immobilization was achieved with Cu2O at a pH as high as 8 and yet 40% at pH 10, from a solution containing initially 10-5 M SeO42- (1.43 ppm). Both CuO and Cu2O sorbents are more efficient than iron oxyhydroxides (e.g., 12) of suspensions containing either CuSeO4 or CuSeO3 results in the formation of Cu(OH)2 with concomitant leaching of either SeO42- or SeO32- in solution. These species can be readily distinguished from each other and in the mixture by CE.72,73 The following experiments have been performed: 0.75 g of Cu2O was suspended into 25 mL of solution containing initially 0.1 M SeVI and 0.5 M H+. After equilibration for 24 h, the remaining selenium species in solution were in the form of SeO42- at a measured concentration of 0.051 M, meaning that 49% of the starting SeVI had been immobilized. The solid phases were then collected and treated in 0.1 M NaOH (0.1 g of solid in 20 mL of solution). The liquid phase was recovered and CE analysis was performed after appropriate dilution, indicating the presence of SeO32- to a great extent (1.65 mmol g-1) with some trace SeO42- species (0.008 mmol g-1). The presence of the first one (SeO32-) demonstrates that selenate species were effectively reduced when contacting Cu2O in an acidic medium to form the sparingly soluble CuSeO3 compound, while the second one is relative to adsorbed SeO42- on Cu2O surfaces, which is known to desorb in an alkaline medium (Figure 4). This detection of SeO32- species confirms, thus, the hypothesis of CuSeO3 formation when contacting Cu2O particles to selenate solutions in acidic media, which was somewhat suggested previously on the basis of XPS results. The overall process fully agrees with thermodynamic calculations that can be made in the Cu-H2O and Se-H2O systems.61,74 In the presence of protons, Cu2O produces transient Cu+ species (eq 2) that are subjected to disproportionation in water in the absence of any other reactant, leading to the formation of equivalent amounts of Cu2+ and Cu0 (eq 3). Both Cu+ and Cu0 are reducing agents likely to transform SeVI into SeIV (eq 7a,b).

2Cu+ + SeO42- + 3H+ f 2Cu2+ + HSeO3- + H2O (log K ) 24.5) (7a) Cu0 + SeO42- + 3H+ f Cu2+ + HSeO3- + H2O (log K ) 18.4) (7b) The selenite species generated according to such redox reactions are in the form of HSeO3- in the pH range observed at equilibrium (4.5-5.5, see Figure 1). These species are in equilibrium with unprotonated SeO32(eq 8) that is known to precipitate in the presence of Cu2+ (eq 5). (72) Gilon, N.; Potin-Gautier, M. J. Chromatogr., A 1996, 732, 369. (73) Walker, E. B.; Walker, J. C.; Zaugg, S. E.; Davidson, R. J. Chromatogr., A 1996, 745, 111. (74) Neal, R. H.; Sposito, G.; Holtzclaw, K. M.; Traina, S. J. Soil Sci. Soc. Am. J. 1987, 51, 1161.

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HSeO3- f SeO32- + H+ (log K ) -8.0)

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(8)

Both processes (involving selenate reduction by either Cu+ or Cu0) lead to the same overall reaction described by the following eq 9.

Cu2O(s) + SeO42- + 4H+ f CuSeO3(s) + Cu2+ + 2H2O (log K ) 23.1) (9) This reaction is significantly shifted to the right but also strongly pH-dependent (e.g., the apparent equilibrium constant at pH 5 is equal to 3.1). This explains why the efficiency of selenate immobilization in the presence of cuprite according to this redox-driven precipitation mechanism decreased dramatically for equilibrium pH values higher than 5-6 (Figure 1). It should be noted, however, that working at lower starting SeVI and H+ concentrations (while keeping the selenate-to-proton ratio equal to 1:4, in agreement with eq 9) resulted in shifting the equilibrium (eq 9) to the left, preventing, therefore, the formation of CuSeO3. For example, a 30 g L-1 Cu2O suspension containing initially 0.005 M SeVI and 0.02 M H+ resulted in the uptake of 90% SeVI by Cu2O and an equilibrium pH of 5.6. In such conditions, no CuSeO3 was detected in the solid phases (after treatment in NaOH and analysis by CE), because eq 9 was shifted to the left (under these conditions, less than 1% of SeO42- species are expected to be transformed according to eq 9 on a thermodynamical basis). In this case, the immobilized selenium species were in the form of selenate, most probably as a CuSeO4 surface precipitate on Cu2O because the amount of fixed SeVI was much higher (0.15 mmol g-1) than that corresponding to an adsorbed monolayer (about 15 times higher). On a thermodynamical basis, both Cu+ and Cu0 are likely to act as the reducing agent toward SeO42- species. However, it appears that, in the presence of a large excess of SeVI and huge amounts of H+ introduced in the medium (it is preferable here to speak about quantities of H+ rather than pH values because the latter are transient and evolve very rapidly as a consequence of proton consumption concomitantly to cuprite dissolution, see Table 1), the quantitative removal of SeO42- species from the solution according to reaction 9 was no longer observed. This was the case for starting SeO42- concentrations higher than 10-2 M, and this can even be noticed in Figure 1C for

which less than 100% removal was observed from a 8 × 10-3 M SeVI solution after addition of more than 0.1 M H+ (it was more evident at a higher concentration: for example, 49% removal from a 0.1 M SeVI solution, see previous). This behavior is not fully understood but might indicate that Cu+ species were mainly involved in the selenate reduction process in a homogeneous pathway (opposite to the heterogeneous Cu0-SeVI reaction). Indeed, the production of a too high amount of Cu+ ions with respect to the SeVI species would kinetically favor the disproportionation reaction (eq 3) in comparison to the CuSeO3 formation by reduction of (less concentrated) solution-phase SeVI (eq 9). This limitation could be also due to the formation of too large amounts of metallic copper (arising from the disproportionation reaction 3), which has been reported to act as a barrier to the further dissolution of cuprite.63 4. Conclusions Selenate in aqueous solutions can be immobilized on cuprite particles, but the extent of uptake is strongly dependent on pH. In a slightly alkaline medium, adsorption of SeO42- onto the hydroxylated Cu2O surface is responsible for the retention process. This event is significantly affected by ionic strength, suggesting the formation of an outer-sphere surface complex. In the presence of protons, cuprite is first dissolved into transient Cu+ species that are then oxidized by the solution-phase selenate to produce the sparingly soluble CuSeO3 compound, this reaction requiring, however, rather high SeVI and H+ starting concentrations. As for other metal oxides, the retention capabilities of cuprite for selenate is less than for selenite species (less quantitatively in an acidic medium and less strongly adsorbed in an alkaline medium). This makes, thus, cuprite and the adsorbent not satisfactory enough to be used as an efficient scavenger for SeVI in the context of radioactive waste management. Acknowledgment. This work has been supported by the French “Agence Nationale pour la Gestion des De´chets Radioactifs” (ANDRA), via a Ph.D. fellowship (J.D.). The authors are also grateful to J. Lambert (LCPME - Nancy) for recording the XPS spectra, to J. Ghanbaja for TEM, and to J. Cortot (LCPME - Nancy) for help in the experiments. LA0496136