Chemical Mechanisms Responsible for the Immobilization of Selenite

The ability of cuprite (Cu2O) to reduce the mobility of selenite species (SeO3. 2- or HSeO3. -), by means of chemical immobilization from an aqueous m...
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Chemical Mechanisms Responsible for the Immobilization of Selenite Species from an Aqueous Medium in the Presence of Copper(I) Oxide Particles Je´roˆme Devoy, Alain Walcarius,* and Jacques Bessiere 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 March 26, 2002. In Final Form: August 11, 2002 The ability of cuprite (Cu2O) to reduce the mobility of selenite species (SeO32- or HSeO3-), by means of chemical immobilization from an aqueous medium, was evaluated from batch experiments. The effect of pH was critically discussed as it was found to have a major influence on the nature of the mechanism involved in the uptake process of selenite by cuprite. Three pH-dependent mechanisms were evidenced by combining quantitative analysis of selenite solutions after equilibration with cuprite, solid-phase characterization, and surface analyses. In the presence of protons, the immobilization of selenite occurs via the formation of a copper selenite solid (i.e., chalcomenite, CuSeO3‚2H2O). The driving force is the acid-catalyzed dissolution of cuprite into CuI species, followed by their fast disproportionation into Cu0 and CuII, this latter reacting with selenite to give the precipitation product. In an alkaline medium (pH > 7.5), selenite was mainly accumulated on the cuprite surface according to two distinct mechanisms depending on the time scale of the experiment. A fast adsorption process was predominant at short equilibration times (typically a few minutes), involving the substitution of the surface hydroxyl groups of Cu2O by the selenium oxyanion. Longer times (several days) led to the progressive coverage of the cuprite surface by transient Cu(OH)2, which was found to react with HSeO3- species to form an insoluble layer on the cuprite particles. Furthermore, in relation to the possible use of cuprite as a scavenger for radioactive selenium species in polluted environments, the competitive influence of chloride anions on the immobilization processes was investigated and discussed with respect to the kind of uptake mechanisms.

1. Introduction Selenium has been recognized for a long time as an essential trace element to many living organisms (humans, animals, and some plants). This micronutrient, however, becomes potentially toxic at high concentrations, and the range between toxic and deficient levels in the diets of humans and animals is quite narrow.1-5 This dual character has led to considerable efforts to understand the biogeochemical cycle of selenium in the environment, to determine its speciation, and to characterize its mobility in natural waters, sediments, and soils, as well as its partitioning within various environmental compartments, to evaluate potential risks arising from deficiency or toxicity.6-13 In parallel, various methods were proposed * To whom correspondence should be addressed. Fax: (+33) 3 83 27 54 44. E-mail: [email protected]. (1) Rosenfeld, I.; Beath, O. A. Selenium Geobotany, Biochemistry, Toxicity, and Nutrition; Academic Press: New York, 1964. (2) Lakin, H. W. In Geochemical Environment in Relation to Health and Disease; Cannon, H. L., Hopps, H. C., Eds.; Special Paper, 140; Geological Society of America; Boulder, CO, 1972; p 27. (3) Food and Nutrition Board Selenium and Human Health Nutrition Reviews; National Academy of Science: Washington, DC, 1976; vol. 34, p 247. (4) Levender, O. A. Fed. Proc. 1985, 44, 2579. (5) Lockitch, G. Crit. Rev. Clin. Lab. Sci. 1989, 27, 483. (6) Lakin, H. W. Geol. Soc. Am. Bull. 1972, 83, 181. (7) Lakin, H. W. Adv. Chem. Ser. 1973, 123, 96. (8) Selenium in Agriculture and the Environment; Jacobs, L. W., Ed.; SSSA Special Publication No. 23; American Society of Agronomy: Madison, WI, 1989. (9) Masscheleyn, P. H.; Patrick, W. H., Jr. Environ. Toxicol. Chem. 1993, 12, 2235. (10) Sager, M. Stud. Environ. Sci. 1993, 55, 459. (11) Selenium in the Environment; Frankenberger, W. T., Jr., Benson, S. M., Eds.; Marcel Dekker: New York, 1994. (12) Bowie, G. L.; Sanders, J. G.; Riedel, G. F.; Gilmour, C. C.; Breitburg, D. L.; Cutter, G. A.; Porcella, D. B. Water, Air, Soil Pollut. 1996, 90, 93. (13) Conde, J. E.; Sanz Alaejos, M. Chem. Rev. 1997, 97, 1979.

to remove selenium species from wastewater, including chemical precipitation, adsorption on activated carbon or on metal oxide minerals, ion exchange, and reverse osmosis (i.e., refs 14-23). Selenium is (biogeo)chemically close to sulfur.24 The most common selenium species include the soluble selenite and selenate, insoluble elemental selenium, and various metal and organic selenides.25,26 The most mobile selenocompounds (i.e., most abundant in natural waters) are selenite and selenate, and to a less extent organo-selenium moieties.13 Selenium is also present in soils, where its mobility is governed by transport across and/or transformations at the interfaces between soil particles and (14) Koren, D. W.; Gould, W. D.; Lortie, L. Waste Process. Recycl. Min. Metall. Ind., Proc. Int. Symp.; Can. Inst. Min. Metall. Pet.: Montreal, Canada, 1992; p 171. (15) Saeki, K.; Matsumoto, S. Commun. Soil Sci. Plant Anal. 1994, 25, 2147. (16) Ghosh-Dastidar, A.; Mahuli, S.; Agnihotri, R.; Fan, L.-S. Environ. Sci. Technol. 1996, 30, 447. (17) Tao, Z.; Du, J.; Dong, W.; Zheng, L. J. Radioanal. Nucl. Chem., Lett. 1996, 214, 245. (18) Vagliasindi, F. G. A.; Benjamin, M. M. Proc.-WEFTEC’96, Annu. Conf. Expo., 69th; Water Environment Federation: Alexandria, VA, 1996; Vol. 3, p 429. (19) Li, C.; Viraraghavan, T. Water Resources and the Urban Environments98, Proceedings of the 1998 National Conference on Environmental Engineering; ASCE Press: Reston, VA, 1998; p 356. (20) Kuan, W.-H.; Lo, S.-L.; Wang, M. K.; Cheng, F. Water Res. 1998, 32, 915. (21) Suzuki, T. M.; Pacheco Tanaka, D. A.; Llosa Tanco, M. A.; Kanesato, M.; Yokoyama, T. J. Environ. Monit. 2000, 2, 550. (22) 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. (23) Tsuji, M.; Ikeda, Y.; Sazarashi, M.; Yamaguchi, M.; Matsunami, J.; Tamaura, Y. Mater. Res. Bull. 2001, 35, 2109. (24) Shrift, A. Fed. Proc. 1961, 20, 695. (25) Elrashidi, M. A.; Adriano, D. C.; Workman, S. M.; Lindsay, W. Soil Sci. 1987, 144, 141. (26) Masscheleyn, P. H.; Delaune, R. D.; Patrick, W. H., Jr. Environ. Sci. Technol. 1990, 24, 91.

10.1021/la025780z CCC: $22.00 © 2002 American Chemical Society Published on Web 09/14/2002

Mechanisms for Immobilization of Selenite Species

the surrounding aqueous medium. Selenite is usually adsorbed onto surfaces of soil minerals and organic matter, more strongly than selenate, and iron oxides are thought to play an major role in this respect.17,27-29 The sorption process is greatly influenced by the pH of the medium, acidic media leading to stronger immobilization. On the other hand, 79Se is a long-lived fission product with a half-life of 1.1 × 106 years,30 which is radiologically toxic. In the context of performance assessment of a disposal concept for nuclear waste in geological sites, it appears extremely important to ensure the most efficient retardation process for limiting the expansion of the radioactive dose in the geosphere, which would arise from transport of the long-lived fission products (mainly 79Se, 99 Tc, and 129I).31 Two important geochemical retardation mechanisms are solubility-controlled precipitation and adsorption by minerals or engineered materials.32 In the hypothesis of repository of nuclear waste in deep geological sites, it is thus necessary to find solids able to trap longlived radionuclides whose migration is not retarded by silicate minerals, which could be used afterward in the design of barriers surrounding the waste containers. Performance assessment of the proposed process would require knowledge of the chemical mechanisms involved in the immobilization of the radioactive species on such scavengers. In a previous paper, we have reported the ability of cuprite (copper(I) oxide, Cu2O) to immobilize iodide species from an aqueous medium, according to two different pHdependent mechanisms (precipitation and adsorption).33 With the idea to restrict the number of additives in the retardation barrier, it is interesting to evaluate the ability of this mineral for trapping selenium species. To our knowledge, no data are available on the ability, or not, for Cu2O to interact with solution-phase selenium species. It was only mentioned recently that extraction of selenite and selenate from groundwater samples can be achieved by using copper(II) oxide particles, CuO, and this behavior was exploited to concentrate them prior to detection by atomic absorption spectrometry.34 On the other hand, numerous studies were dealing with the uptake of selenite and selenate oxyanions by other metal oxides (metal ) Fe, Al, Mn),15,19,20,35-47 for which physical or chemical adsorption was most often recognized as the process (27) Goldberg, S.; Glaubig, R. A. Soil Sci. Soc. Am. J. 1988, 52, 954. (28) Fio, J. L.; Fujii, R.; Deverel, S. J. Soil Sci. Soc. Am. J. 1991, 55, 1313. (29) Xiangke, W.; Dong, W.; Jie, Y.; Tao, Z. J. Radioanal. Nucl. Chem. 1999, 242, 815. (30) Jiang, S.; Duo, J.; Jiang, S.; Li, C.; Cui, A.; He, M.; Wu, S.; Li, S. Nucl. Instrum. Methods 1997, B123, 405. (31) Melnyk, T. W.; McMurry, J.; Sargent, F. P. High-Level Radioactive Waste Management: Proceedings of the Fifth Annual International Conference, Las Vegas, Nevada, May 22-26, 1994; American Nuclear Society: LaGrange Park, IL, 1994; Vol. 3, p 1222. (32) Chen, F.; Burns, P. C.; Ewing, R. C. J. Nucl. Mater. 1999, 275, 81. (33) Lefe`vre, G.; Walcarius, A.; Ehrhardt, J.-J.; Bessie`re, J. Langmuir 2000, 16, 4519. (34) Reddy, K. J.; Zhang, Z.; Blaylock, M. J.; Vance, G. F. Environ. Sci. Technol. 1995, 29, 1754. (35) Hingston, F. J.; Posner, A. M.; Quirk, J. P. Adv. Chem. Ser. 1968, 79, 82. (36) Hayes, K. F.; Roe, L.; Brown, G. E., Jr.; Hodgson, K. O.; Leckie, J. O.; Parks, G. A. Science 1987, 238, 783. (37) Balistrieri, L. S.; Chao, T. T. Soil Sci. Soc. Am. J. 1987, 51, 1145. (38) Hayes, K. F.; Papelis, C.; Leckie, J. O. J. Colloid Interface Sci. 1988, 125, 717. (39) Zhang, P.; Sparks, D. L. Environ. Sci. Technol. 1990, 24, 1848. (40) Ghosh, M. M.; Cox, C. D. Environ. Prog. 1994, 13, 79. (41) Manceau, A.; Charlet, L. J. Colloid Interface Sci. 1994, 68, 87. (42) Papelis, C.; Brown, G. E., Jr.; Parks, G. A.; Leckie, J. O. Langmuir 1995, 11, 2041. (43) Saeki, K.; Matsumoto, S.; Tatsukawa, R. Soil Sci. 1995, 160, 265.

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responsible for mass transfer reactions, as common for the accumulation of anionic species at oxide/water interfaces.48 Adsorption of these selenium species was pH dependent, and their removal from diluted solutions increased with decreasing pH and increasing particle concentration. Selenite was usually strongly adsorbed as an inner-sphere surface complex, while selenate displayed a relatively lower affinity for the metal oxide surfaces, as evidenced by significant variation of the uptake capacity with ionic strength of the solution.20,36-38,40 The objectives of the present work are to address the question of the ability for cuprite (Cu2O) to accumulate selenite species from an aqueous medium and to determine the nature of the interactions occurring at the solid/liquid interface. This will be achieved from batch equilibration experiments involving the dispersion of Cu2O particles into selenite solutions (adjusted to appropriate pH values), by means of both solution-phase analyses and solid characterization using various bulk solid and surface analysis methods. The competitive effect of chloride species on the sorption processes was also evaluated and discussed as a function of the uptake mechanism. 2. Experimental Section 2.1. Chemicals and Solutions. The cuprous oxide, cuprite, was purchased from Fluka as a dry powder and pretreated in a slightly acidic medium (∼10-4 M HClO4) during 12 h to remove the oxidizing layer that was present on the particle surface.33 The solid particles were then filtered, copiously washed with pure water, and dried in the absence of oxygen. These steps contribute to enhance the hydrophilic properties of the material, allowing therefore an optimal dispersion of suspended Cu2O in an aqueous medium. All other chemicals were analytical grade: Na2SeO3 (Fluka), NaCl (OSI), Cu(NO3)2, HClO4, and NaOH (Prolabo). Nitrogen gas was bubbled into batch reactors to remove residual oxygen from the solutions. All solutions were prepared with ultrapure water (18 mΩ cm-1) from a Millipore Milli-Q water purification system. 2.2. Batch Experiments and Solution Analyses. Batch equilibrations were performed at 20 °C in stoppered polyethylene tubes (Nalgene) containing typically 25 mL of the selenite solution at a selected initial concentration, into which Cu2O particles were suspended at a level ranging from 10 to 80 g L-1. pH was adjusted by addition of strong acid (HClO4) or base (NaOH). Supernatant pH values were measured before and after equilibration, using a model 605 pH-meter (Metrohm). Samples were then centrifuged at 10 000 rpm for 30 min using a Jouan GR20.22 apparatus and passed through Durapore membrane filters (0.45 µm pore size, Millipore). Chemical analysis of the filtered supernatant was carried out either by capillary electrophoresis (CE) or inductively coupled plasma atomic emission spectrometry (ICP-AES). Determination of selenite by CE was performed at 25 °C with a capillary ion analyzer equipped with a negative high-voltage power supply (Waters, CIA) and a bare fused silica capillary (length, 60 cm; diameter, 75 µm). Detection was achieved on-line by indirect spectrophotometry at a wavelength of 254 nm. The electrolyte composition was 4.6 M Na2CrO4‚4H2O (OSI) + 2.5 M osmotic flow modifier (OFM-OH from Waters) + 6.4 M H3BO3 (Prolabo). Samples were injected in hydrostatic mode at the cathodic side of the capillary subject to an applied voltage of 20 kV. In these conditions, the observed migration time for selenite was 3.8 min, with a total run time of the experiment around 5 min. The ICP-AES apparatus (Plasma 2000, Perkin(44) Parida, K. M.; Gorai, B.; Das, N. N.; Rao, S. B. J. Colloid Interface Sci. 1997, 185, 355. (45) Parida, K. M.; Gorai, B.; Das, N. N. J. Colloid Interface Sci. 1997, 187, 375. (46) Davis, S. A.; Misra, M. J. Colloid Interface Sci. 1997, 188, 340. (47) Suarez, D. L.; Goldberg, S.; Su, C. ACS Symp. Ser. 1998, 715, 136. (48) Behra, P.; Douch, J.; Binde, F. In Effect of Mineral-OrganicMicroorganism Interactions on Soil and Freshwater Environments; Berthelin, J., Huang, P. M., Bollag, J.-M., Andreux, F., Eds.; Kluwer Academic/Plenum Publishers: New York, 1999; p 1.

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Elmer) was used to determine the soluble fraction of copper and to control selenium concentrations measured by CE. The detector was adjusted to a wavelength of 204.0 or 324.8 nm for Se or Cu, respectively. Measurements were performed by the standard addition method. 2.3. Apparatus for Solid-Phase Analyses. Electron microscopy was used to identify Se-rich phases that would arise from the reaction of Cu2O particles with selenite species. Samples were first randomly surveyed using a Philips XL30 scanning electron microscope (SEM) equipped with a thin-window energydispersive X-ray (EDX) system (KEVEX Sigma) for compositional analysis. This technique was applied to the visualization of copper selenite crystals based on backscattered electron contrast. For transmission electron microscopy (TEM), small fractions of cuprite samples after exposition to selenite solutions including the control samples without selenite were quickly dispersed in air onto amorphous carbon-coated grids and loaded into the analysis holder of a Philips CM20 microscope operating at 200 kV with an unsaturated LaB6 cathode. This apparatus was coupled with an EDX spectrometer equipped with an ultrathin window X-ray detector. Bulk analysis of powders was also performed by X-ray diffraction (XRD) using a classical powder diffractometer with transmission geometry, equipped with a Mo tube (quartz monochromator, KR1 radiation, λ ) 0.070930 nm) and a scintillation counter. The surface of cuprite samples, before and after reaction with selenite, was analyzed by X-ray photoelectron spectroscopy (XPS) after thorough washing with ultrapure water. XPS spectra were obtained at a residual pressure lower than 10-9 mbar, by using a VSW HA150 MCD electron energy analyzer operating with an Al KR nonmonochromatic source (1486.6 eV). The energy scale was calibrated using the aliphatic adventitious hydrocarbon C1s peak located at a binding energy of 284.6 eV. Narrow scanned spectra were used to obtain the chemical state information for copper, selenium, and oxygen.

3. Results and Discussion 3.1. Preliminary Experiments: Batch Equilibration as a Function of pH. At first, the ability of cuprite to remove selenite species from aqueous solutions was evaluated by adding the same SeIV concentration to several cuprite suspensions with various amounts of protons or eventually hydroxide species (starting suspensions of Cu2O-SeO32- mixtures were intrinsically basic, with pH values ranging from 9 to 10). Consumption of SeIV was measured as a function of time. Stationary concentrations of remaining SeIV in solution were reached after about 2 weeks, and corresponding equilibrium values for immobilized SeIV have been plotted as a function of final pH (Figure 1A). At first glance, it appears that variation in the uptake efficiency follows the general trend previously observed for selenite adsorption on iron oxides and clays,35,37,49 where the extent of SeIV binding was high at low pH values and decreased rapidly with rising pH. The case of cuprite sorbent is however much more complicated, and more in-depth analysis of results displayed in Figure 1 reveals some intriguing features. First, nonzero uptake was observed at high pH values (>9-10), contrary to what has been always reported for selenite adsorption on iron oxyhydroxides or aluminum and manganese oxides.37,42,44,45 Here, about 30% selenite binding was still achieved at pH values above 11, so that another process in addition to surface complexation is expected to occur in the Cu2O-SeIV system. Second, significant pH variations have been observed during the reaction itself, depending on the amount of added acid or base in the starting solution (some illustrative examples are given in Figure 1B). These are clearly due to the pHsensitive reactants, as cuprite can dissolve in the presence

of protons50 and SeIV does exist under three pH-dependent forms (H2SeO3, HSeO3-, and SeO32-). As a consequence, pH adjustment with strong acid or base was difficult and subject to great consumption of reactants; this point will be carefully taken into account when characterizing matter balance before and after reaction. Working with unbuffered solutions and uncontrolled ionic strength was however preferred here to avoid any side effect on the uptake processes (e.g., Cl- species can react with CuI to form sparingly soluble CuCl, buffer components could adsorb on the cuprite surface, competitive adsorption of anions could occur as previously observed for selenium adsorption by goethite,37 etc.). Finally, the rather large dispersion of the isotherm data seems to indicate a nonnegligible dependence of the uptake processes on small variations in the experimental conditions over prolonged reaction times (various cuprite samples, small differences in the amounts of added acid or base, efficiency of the

(49) Neal, R. H.; Sposito, G.; Holtzclaw, K. M.; Traina, S. J. Soil Sci. Soc. Am. J. 1987, 51, 1161.

(50) Beverskog, B.; Puigdomenech, I. J. Electrochem. Soc. 1997, 144, 3476.

Figure 1. (A) Variation of immobilized SeIV species from 25 mL solutions containing initially 0.75 g of Cu2O and 4 × 10-3 M SeO32- (adjusted at various starting pHs), as a function of final pH after 450 h equilibration. (B) Corresponding evolutions of pH for some data points of part A: (a) [H+]i ) 2.3 × 10-1 M; (b) [H+]i ) 1.4 × 10-2 M; (c) [H+]i ) 8.0 × 10-3 M; (d) [H+]i ) 4.0 × 10-3 M; (e) [H+]i ) 2.0 × 10-3 M; (f) no pH adjustment; (g) [OH-]i ) 4.0 × 10-3 M.

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solution outgassing, performance of airtightness to limit residual oxygen, etc.). Such variations are also due to the intervention of several immobilization mechanisms (see below). 3.2. Bulk Precipitation in an Acidic Medium. According to the Pourbaix diagram relative to the copperoxygen system, Cu2O undergoes disproportionation in an acidic medium.50 The chemical pathway involves at first the production of Cu+ species (eq 1), which immediately give equivalent amounts of Cu2+ and Cu0 in the absence of any stabilizing ligands (eq 2). The overall reaction is given by eq 3.

Cu2O(s) + 2H+ f 2Cu+ + H2O 2Cu+ f Cu0(s) + Cu2+

(log K ) -1.1) (1) (log K ) 6.1)

(2)

Cu2O(s) + 2H+ f H2O + Cu0(s) + Cu2+ (log K ) 5) (3) On the other hand, selenite is known to form sparingly soluble compounds when reacting with copper(II) species: amorphous cupric selenite (eq 4a) or crystalline chalcomenite (eq 4b).25

Cu2+ + SeO32- f CuSeO3(s)

(log K ) 7.8)

(4a)

Cu2+ + SeO32- + 2 H2O f CuSeO3‚2H2O(s) (log K ) 7.78) (4b) These reactions are pH dependent because SeIV can exist under two protonated forms (H2SeO3 and HSeO3-) in an acidic medium and CuII precipitates as cupric hydroxide in an alkaline medium. Therefore, the formation of cupric selenite (or chalcomenite) occurs at intermediate pH (e.g., between 5 and 10 for total concentrations of CuII and SeIV equal to 10-2 M). No redox reaction between transient Cu+ species and SeIV is expected on a thermodynamic basis. From the above thermodynamic considerations, it appears that contacting a cuprite suspension with SeIV in the presence of protons would lead to precipitation of cupric selenite, at least if added SeIV and produced CuII are in concentrations high enough to fulfill the solubility product (pKs ≈ 7.8). This is illustrated in Figure 2 where the backdiffusion scanning electron micrographs obtained on solid phases recovered after treating 0.75 g of Cu2O in 25 mL of solution containing initially 0.1 M H+ and 0.1 M SeO32are depicted. They show a mixture of two solids: one made of small particles of irregular shape (typical of Cu2O powder) and the other in the form of crystals with welldefined shape and particle size of about 15 × 30 µm. While EDX analysis of the small grains did not reveal a significant amount of selenium (only Cu2O), data obtained when focusing on a large crystal indicate the presence of selenium in the particle. A Se/Cu atomic ratio of 1.2 was evaluated, which is close to unity and argues for the formation of the expected cupric selenite compound. Chalcomenite was indeed detected by XRD (see Supporting Information). Despite the large amount of protons initially introduced in the medium, a final pH above 5.5 was measured after the reaction, confirming proton consumption during cuprite dissolution (eq 3). The overall reaction was fast (95% completed in less than 1 min), and equilibrium was reached within a few hours. Decreasing both proton and selenite concentrations (e.g., typically in the 10-3 M range) led to smaller cupric selenite crystals, which presented star and needle shapes (see Supporting Information). Below 10-4 M, no crystals were evidenced,

Figure 2. (a,b) Scanning electron micrographs of the solid phases obtained after reaction of 0.75 g of Cu2O with 0.1 M SeIV in the presence of 0.1 M HClO4; (c) EDX analysis of the small particles; (d) EDX spectra obtained when focusing on a large crystal.

in agreement with the corresponding SeIV and CuII concentrations being too low to give the precipitated product. This enables the overall process to be stated as follows. Cuprite is first “dissolved” in the presence of H+ and then reacts with selenite species (i.e., HSeO3- in such a pH range) to give chalcomenite (eq 5). This reaction is pH dependent and directly a function of both H+ and SeIV concentrations. It is quite obvious that a SeIV concentration that was too low did not give significant amounts of chalcomenite. On the other hand, its formation in the presence of a large excess of SeIV is totally controlled by the amount of H+ introduced in the medium (strictly speaking, one has to consider quantities of H+ rather than pH values as the latter are transient and evolve very rapidly as a consequence of proton consumption concomitant to cuprite dissolution).

Cu2O(s) + HSeO3- + H+ + 2H2O f CuSeO3‚2H2O(s) + Cu0(s)

(log K ) 5.5) (5)

Reaction of Cu2O with selenite species in the presence of various amounts of protons added in the medium (initial concentrations varying from 0.1 to 10-4 M) was monitored by XPS. Some typical spectra relative to copper are given in Figure 3. The binding energy of the Cu2p3/2 line for Cu2O is pointed out at 932.5 eV, in agreement with previously reported data for the same compound,33,51 while pure chalcomenite containing CuII centers is characterized by a Cu2p3/2 line located at 934.8 eV together with a large satellite observed around 940-945 eV (this includes two (51) Chawla, S. K.; Sankarraman, N.; Payer, J. H. J. Electron Spectrosc. Relat. Phenom. 1992, 61, 1.

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Figure 4. Evolution of the amount of selenite species bound to cuprite as a function of reaction time (inset: magnification of the short-time part). The suspensions were prepared to contain 2.0 × 10-3 M SeIV and 30 g L-1 Cu2O, at pH 11.

Figure 3. Montage of the Al KR XPS spectra of the Cu2p3/2 line recorded on the solid phases obtained after reaction of 0.75 g of Cu2O with 0.1 M SeIV in the presence of (a) 2.0 × 10-3 M HClO4, (b) 1.4 × 10-2 M HClO4, (c) 2.3 × 10-1 M HClO4, and (d) no added protons, as compared to pure cuprite (Cu2O) and pure chalcomenite (CuSeO3).

bands, as generally obtained for other CuII-bearing compounds).52,53 Increasing progressively the amount of protons in Cu2O suspensions containing 0.1 M SeIV resulted in progressive growing of the cupric selenite contribution (Figure 3, lines a-c). A concomitant increase in the Se3d line was also observed at 58.7 eV (data not shown). In the absence of SeIV in the medium, no contribution of CuII has been evidenced in the XPS spectrum because dissolution of cuprite produces soluble copper in the form of Cu2+ ions and contribution of Cu0 to the Cu2p line occurs at the same binding energy as that of CuI.33 3.3. Sorption in a Neutral or Alkaline Medium. 3.3.1. Initial Observations.In the pH range where cuprite is thermodynamically stable (>7), the mechanism responsible for selenite uptake would involve only adsorption processes, in agreement with previous studies with other metal oxides or oxyhydroxides (FeOOH, Fe2O3, Al2O3, MnO2, CuO) at the surface of which inner-sphere complexes between hydroxyl groups and selenite moieties have been reported.15,19,20,34-47 In the case of Cu2O, however, several experimental features claim for a more complex behavior. First, the XPS spectrum of the Cu2p3/2 line recorded from a Cu2O-SeIV mixture without any added protons reveals a rather large satellite area around 940-945 eV (typical of CuII) although CuSeO3 should not form under these conditions, as expected from the negligible dissolution of cuprite at a pH as high as 10 and as otherwise evidenced by very low signal of the Se3d line sampled with the same solid. Second, adsorption is known to be a rather fast process, at least when involving nonporous solids for which no diffusional limitations occur54 (e.g., selenite sorption on iron oxyhydroxide reaches (52) Wagner, C. D.; Bickham, D. M. NIST Standard Reference Database 2.0, NIST XPS Database 1.0, 1989. (53) Lefe`vre, G.; Alnot, M.; Ehrhardt, J.-J.; Bessie`re, J. Environ. Sci. Technol. 1999, 33, 1732.

equilibrium in 2 h and no further changes in equilibrium concentrations are observed later on).44 As depicted in Figure 4, the uptake of SeIV by Cu2O is much slower and equilibrium is reached after about 2 weeks. Moreover, the whole process seems to be separated into two parts: a fast consumption step at the beginning and a slow continuing uptake at extended equilibration periods. This behavior is not common, but it was already observed for the reaction of orthophosphate at the alumina/water interface where both surface complexation and phase transformation occurred according to fast and slow processes, respectively.55,56 Third, comparing the variation of selenite uptake as a function of cuprite density in the medium reveals a significant difference between short and long equilibration periods (Figure 5). Results obtained at short times indicate a direct relation between sorbed SeIV and the amount of Cu2O in suspension, agreeing well with adsorption phenomena, while longer times led to a nearly constant additional contribution that was independent of the cuprite density between 10 and 80 g L-1. All these features suggest the intervention of (at least) two mechanisms. 3.3.2. Surface Adsorption at a Short Equilibration Time. On the basis of several observations made above (Figures 4 and 5), it is likely that adsorption happens during the first stage of the reaction between solutionphase selenite and the cuprite solid. Some short-time experiments were then performed to assess the adsorption process. The first pseudo-equilibrium stage observed on the kinetic curve relative to selenite removal from the aqueous phase corresponds to an accumulated amount of about 8 µmol per g of Cu2O, in conditions of a large excess of selenite in solution (Figure 4, inset). Taking into account the ionic radius of SeO32- (0.239 nm)57 and the specific surface area of the Cu2O sample (1.6 m2 g-1), one can estimate a surface coverage close to a monolayer (about 80%) and equal to about 5 µmol m-2. This is more than, yet comparable to, (54) Sposito, G. The surface chemistry of soils; Oxford University Press: New York, 1984. (55) Laiti, E.; Persson, P.; Ohman, L.-O. Langmuir 1996, 12, 2969. (56) Laiti, E.; Persson, P.; Ohman, L.-O. Langmuir 1998, 14, 825. (57) Bernard, M.; Burnst, F. Usuel de Chimie Ge´ ne´ rale et Mine´ rale; Dunod: Paris, 1996; 96/03.

Mechanisms for Immobilization of Selenite Species

Figure 5. Variation of the amount of selenite species bound to cuprite as a function of adsorbent concentration. Data points were recorded after (a) 5 min reaction and (b) at equilibrium (450 h). The suspensions were prepared to contain 4.0 × 10-3 M SeIV and various amounts of Cu2O in 25 mL solutions (final pH values were around 10).

Figure 6. Variation of the distribution coefficient relative to sorption amount of selenite on cuprite as a function of adsorbent concentration. Data points were recorded after (a) 5 min reaction and (b) at equilibrium (450 h). The suspensions were prepared as in Figure 5.

values for selenite adsorption on aluminum oxides or iron oxyhydroxides for which values of typically 1-3 µmol m-2 have been reported.42,44 Though rather high, this coverage level was however not sufficient to give a measurable XPS signal of the Se3d line at this early stage of the reaction. To further support the hypothesis of adsorption, uptake experiments have been performed with various amounts of Cu2O in the presence of a large excess of SeIV in solution. The quantity of accumulated species was found to be directly proportional to the Cu2O content in suspension (Figure 5, curve a) and, therefore, to the total surface area of Cu2O available for adsorption. The Langmuir equation was respected, and the calculated distribution coefficients (KD) were constant at about 5 L kg-1 (Figure 6, curve a) over the whole Cu2O content range considered here (1080 g L-1), in agreement with similar experiments performed for selenite adsorption on ferric oxyhydroxides.44

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Decreasing the SeIV-to-Cu2O ratio led to significant increase of KD values (e.g., about 50 L kg-1 for 10-5 M SeIV and 20 g L-1 Cu2O), as often observed in adsorption experiments.37,44 One can therefore conclude that the main mechanism responsible for the immobilization of selenite species on cuprite at the early stage of the reaction is adsorption. No attempt was made here to ascribe the exact nature of the cuprite-selenite interaction, but a satisfactory assumption would be the formation of an inner-sphere surface complex by binding of selenite species to the terminal hydroxyl groups on cuprite, as always observed for adsorption of selenite on other metal oxides or oxyhydroxides.15,19,20,34-47 This is supported by decreased adsorption yield when increasing pH and by the fact that variable ionic strength in the medium (up to 0.1 M) did not result in significant changes in the adsorption isotherm. 3.3.3. Surface Precipitation at a Long Equilibration Time. The sorption behavior after reaching equilibrium cannot be explained only by a simple adsorption process. First, it takes a long time (Figure 4) and the amount of sorbed SeIV species at equilibrium is much higher than that corresponding to one monolayer. Second, the variation of sorbed quantities as a function of cuprite content does not pass through the origin and displays a rather large intercept on the y-axis (Figure 5, curve b). Third, KD values are not constant when varying the Cu2O content in suspension (Figure 6, curve b). All these observations claim for the existence of another phenomenon in addition to adsorption to explain the uptake processes. As no redox reaction between SeIV and Cu2O is expected, as complexation was already involved in the adsorption path, and according to the aforementioned results, the additional phenomenon would take into account phase transformation eventually associated with acid-base equilibrium (as the uptake processes are pH dependent over the whole period of time before reaching equilibrium; see curves f and g in Figure 1B). Figure 7 illustrates the effect of aging a Cu2O sample in an aqueous medium on its surface properties, in the absence of selenite, through recording XPS spectra of Cu2p3/2 and O1s lines after 5 days at two different pH values (10 and 12). As shown, the cuprite surface is oxidized at pH 10 (the natural pH observed for Cu2O in water without adding any reactant), and even more at a higher pH (Figure 7, part A, curves b and c). The Cu2p3/2 line typical of Cu2O (main peak at 932.5 eV) shows progressive growth of an additional signal at 935.0 eV (with a large satellite including two bands at ca. 940-945 eV), indicating the presence of CuII centers52,53 in the form of an oxidation layer around the cuprite particles. After comparison to the spectra of pure Cu(OH)2 and CuO (Figure 7, curves d and e), it seems that the oxidation layer formed onto the cuprite surface is most probably Cu(OH)2, but localization of its Cu2p3/2 main peak (935.0 eV) is too close to that of CuO (933.9 eV) to conclude unambiguously. Confirmation of the hypothesis was however given by comparison of the O1s lines, as the signal for Cu(OH)2 is shifted to the left (531.2 eV) with respect to that of Cu2O (530.1 eV) while that of CuO is shifted to the right (529.6 eV), in agreement with previous XPS standardization of copper oxide and hydroxide compounds.51 Decomposition spectra for aged cuprite at pH 10 and 12 clearly show an increased contribution of the O1s line for Cu(OH)2, via the progressive decrease of the signal located at 531.2 eV and the concomitant growth of that situated at 531.2 eV (Figure 7, part B, curves b and c); the small peak at 532.8 eV is due to residual water.

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Figure 7. Montages of the Al KR XPS spectra of the Cu2p3/2 line (A) and O1s line (B) recorded on (a) pure Cu2O, (b) Cu2O after 5 days in water (pH 10), (c) Cu2O after 5 days in 10 mM NaOH (pH 12), (d) freshly prepared Cu(OH)2, and (e) pure CuO.

This confirms the presence of Cu(OH)2 on the cuprite surface upon aging in an alkaline medium. The origin of this oxidation layer is not fully understood actually but could be due either to chemical oxidation by residual oxygen in the medium (that might be non-negligible at such long equilibration times) or to some slow dissolution of cuprite (solubility of Cu2O up to 10-6 M at pH 10 has been reported)50 followed by disproportionation and precipitation of Cu(OH)2. Anyway, the Cu(OH)2 layer does exist, and one must explain how it can intervene in the immobilization of SeIV species. In the presence of SeIV in the medium, long equilibration times gave rise to a regular thin layer at the outermost surface of Cu2O particles, as evidenced by transmission electron microscopy (Figure 8). This appears clearer than the bulk portion of Cu2O and energydispersive analysis of X-rays and reveals the presence of a significant part of selenium in addition to copper and

Devoy et al.

oxygen (the oxygen signal is also more intense than in the bulk particle). It should be recalled that no measurable selenium signal was observed by applying this technique at short reaction times corresponding to adsorption of selenite on Cu2O. The thickness of the layer was roughly estimated to that of a pile of 4-8 monolayers. The above results demonstrate the presence of selenite in this layer that was initially constituted of Cu(OH)2 as otherwise pointed out by XPS measurements. The remaining questions are how to explain these results and what is the exact nature of the SeIV-CuII compound that has been formed. On the basis of thermodynamic considerations, SeO32species are not liable to react with Cu(OH)2 (i.e., to form CuSeO3), as otherwise confirmed by a simple chemical test (Table 1). Also, treating a CuSeO3 sample with a strong base such as NaOH resulted in the destruction of CuSeO3 and formation of Cu(OH)2 with, at the same time, solubilization of SeO32- species (Table 1). CuSeO3 is not stable in a strong alkaline medium. In a slightly alkaline medium where both Cu2O and Cu(OH)2 solids are stable, however, the protonated form of selenite (HSeO3-) can exist and is liable to react with Cu(OH)2 (Table 1). The overall transformation involves an acid-base reaction followed by precipitation of CuSeO3, in which two HSeO3moieties are necessary to liberate Cu2+ from Cu(OH)2: one results from selenite precipitation as CuSeO3, while the second remains in solution (eq 6). This reaction is pointed out by some simple chemical tests (Table 1), which brought together synthetic Cu(OH)2 and selenite at variable protonation ratios, to mimic the situation liable to occur at the cuprite surface after aging. The loss of selenite from solution can be explained only by precipitation as CuSeO3, which was higher when increasing the amount of protonated selenite, in a stoichiometry agreeing with the equilibrium reaction of eq 6 (e HSeO3- consumed by CuSeO3 precipitation).

Cu(OH)2(s) + 2HSeO3- f CuSeO3(s) + SeO32- + 2H2O (6) Coming back to the immobilization of SeIV on cuprite after long equilibration times, Cu(OH)2 arises from cuprite aging, and the presence of HSeO3- is controlled by pH. It could also result from deprotonation of the surface hydroxyl groups on cuprite by SeO32-, as otherwise observed with other bases. For example, titration of the

Figure 8. Transmission electron micrograph (+ enlargement of the bottom) of a Cu2O crystal after 450 h reaction in 0.01 M SeIV; (a) EDX analysis of the particle center and (b) EDX spectra obtained when focusing on the 3-4 nm thick layer situated at the outermost boundary of the crystal.

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Table 1. Some Chemical Tests to Illustrate the Chemical Reactions between Cu(OH)2 and SeIV Species (H2SeO3, HSeO3-, and SeO32-) in Unbuffered Solutions starting reactants

final pH

equilibrium SeIV concentration (M)

1 × 10-2 M Cu(NO3)2 + 2 × 10-2 M NaOH + 1 × 10-2 M Na2SeO3 1 × 10-2 M Cu(NO3)2 + 2 × 10-2 M NaOH + 0.7 × 10-2 M Na2SeO3 + 0.3 × 10-2 M NaHSeO3 1 × 10-2 M Cu(NO3)2 + 2 × 10-2 M NaOH + 0.4 × 10-2 M Na2SeO3 + 0.6 × 10-2 M NaHSeO3 1 × 10-2 M Cu(NO3)2 + 2 × 10-2 M NaOH + 1 × 10-2 M NaHSeO3 1 × 10-2 M Cu(NO3)2 + 2 × 10-2 M NaOH + 0.7 × 10-2 M NaHSeO3 + 0.3 × 10-2 M H2SeO3 CuSeO3 in pure water (40 g L-1) CuSeO3 in 10-2 M NaOH (40 g L-1)

9.8 9.3

9.8 × 10-3 8.7 × 10-3

≈10-5 2 × 10-5

8.7

7.3 × 10-3

4 × 10-5

7.9 7.7

5.5 × 10-3 3.6 × 10-3

7 × 10-5 9 × 10-5

6.1 9.5

8.6 × 10-5 4.6 × 10-3

1.1 × 10-4