Environ. Sci. Technol. 2001, 35, 2252-2257
Aqueous Geochemical and Surface Science Investigation of the Effect of Phosphate on Pyrite Oxidation ALICIA R. ELSETINOW,† MARTIN A. A. SCHOONEN,‡ AND D A N I E L R . S T R O N G I N * ,† Department of Chemistry, Temple University, Philadelphia, Pennsylvania 19122, and Department of Geosciences, SUNY-Stony Brook, Stony Brook, New York 11794-2100
Aqueous geochemical and surface science techniques were used to investigate the effect of phosphate on the oxidation of crushed and {100} pyrite. Studies showed that the presence of phosphate in solution significantly impeded the oxidation rate of crushed pyrite at pH values equal to or greater than 4. At pH 3, the presence of phosphate had almost no experimentally discernible effect on pyrite oxidation. X-ray photoelectron spectroscopy (XPS) studies showed that phosphate, at pH values greater than 4, became irreversibly bound to an Fe3+-bearing product on the pyrite surface during the oxidation process. Once bound to this region on the pyrite surface, the adsorbed phosphate inhibited further oxidation of the pyrite (based on XPS determinations of sulfur and iron oxide product concentrations) under our experimental conditions. These results suggested that the rate of pyrite oxidation in the absence of phosphate was facilitated at or near Fe3+bearing oxidation phases on the surface. Phosphate bound on iron(III) oxide product either prevents O2 adsorption on this phase or electronically modifies these surface regions, but in either case it inhibits electron transfer from pyrite-Fe2+ sites to molecular O2.
Introduction Oxidation of pyrite and other metal sulfides present in mine waste leads to acidification of waters (acid mine drainage) and the mobilization of toxic metals, such as arsenic and selenium. With miles of streams in the United States alone impacted by acid mine drainage, there is a need to understand the oxidation process and to develop strategies to abate acid mine drainage formation. The development of abatement strategies requires that the fundamental steps of the oxidation process are understood. On the basis of various types of modern surface studies (1, 2), aqueous geochemical studies (3, 4), and theoretical studies (5, 6), pyrite oxidation involves a sequence of elementary steps that in part take place at the mineral surface. Most abatement strategies are designed to inhibit the interactions of the oxidants with the pyrite surface. For example, various types of organic and inorganic compounds have been added to pyrite with the objective to either create a protective layer on the pyrite surface or specifically bind to reactive surface sites (7-10). Recently, surface science studies have started to unravel those sites on the pyrite surface that control its chemistry * Corresponding author telephone: (215)204-7119; fax: (215)2041532; e-mail:
[email protected]. † Temple University. ‡ SUNY-Stony Brook. 2252
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(1, 11). In the present study, we build upon this prior work to understand the effect of phosphate (a proposed abatement chemical) on the oxidation chemistry of pyrite. Prior work concerning the effect of phosphate on the rate of pyrite oxidation has been studied under conditions in which a protective coating of ferric phosphate formed on the pyrite (7). The work detailed in this paper investigated the interaction of phosphate with pyrite under experimental conditions where encapsulation of the pyrite could not occur. However, it is shown that phosphate was still able to exhibit an inhibitory effect on the oxidation of pyrite by binding to active sites on the pyrite surface. Importantly, the use of phosphate in this study as an adsorbate turned out to be a useful probe molecule for shedding light on the mechanism of pyrite oxidation. The strategy employed in this study was to combine macroscopic observations with molecular level observations. The macroscopic observations consisted of reaction rate data obtained in batch oxidation experiments using pyrite slurries in the pH range of 3-6 with and without phosphate. The molecular level observations were derived from surface science studies utilizing X-ray photoelectron spectroscopy (XPS). These studies helped to understand the microscopic aspects of the mode of oxidation inhibition exhibited by phosphate adsorbed on the mineral surface. The integration of both types of techniques yielded a detailed picture of how phosphate affects the kinetics and surface reactivity of pyrite.
Experimental Section Bath Oxidation Experiments. The methods and procedures for these experiments have been described in detail in an earlier publication (12). In brief, 3 g of crushed, acid-washed pyrite was added to 750 mL of N2-purged DI contained in a stirred, water-cooled Pyrex vessel. The vessel was mounted on an optical bench equipped with a 1000-W xenon lamp. After addition of the pyrite, the suspension was kept under a nitrogen atmosphere for 6-24 h. The pH was kept constant using a pH stat (Titroline Alpha). During this initial period, the pyrite particles underwent some reaction that led to the release of some iron and sulfate. After this period, pure oxygen was bubbled through the suspension, and the release of protons (measured as uptake of sodium hydroxide titrant) was continuously recorded. These measurements were complemented by sulfate and iron determinations on filtered solution aliquots. Aliquots were filtered using membrane filters with a nominal pore size of 0.2 µm. Once a steadystate oxidation rate was attained, phosphate was added in increments, so that the effect of phosphate concentration on reaction rate could be determined. Phosphate was added in the form of a 5 mM sodium phosphate solution doped with a known amount of KBr. After the phosphate solution was added, aliquots (about 5 mL) were periodically drawn to determine the uptake of phosphate onto the pyrite particles, the sulfate concentration, and the bromide concentration. It was assumed that bromide was not taken up by the presumably negatively charged pyrite surface (12). On the basis of this assumption, bromide concentrations were used to calculate the total phosphate concentration in the suspension. (Note that the volume in the batch experiments changes as aliquots are withdrawn and titrant is added.) Bromide, sulfate, and phosphate concentrations were determined in triplicate using ion chromatography (Dionex DX-500 equipped with an AS4A-SC column). Total dissolved iron was analyzed spectrophotometrically (HACH Method 8008). In a few experiments, the effect of illumination was studied on phosphate-containing suspensions by exposing 10.1021/es0016809 CCC: $20.00
2001 American Chemical Society Published on Web 04/27/2001
FIGURE 1. Rate of pyrite oxidation (in terms of acid production rate) at pH 6 vs time (experiment P19, see Table 1). The curve consists of segments that represent the rate of reaction after the addition of PO4. The inset shows the partitioning of phosphate between solution and pyrite for each segment indicated by the letters. The oxidation rate showed a significant decrease with phosphate addition. the vessel to a water-cooled beam of visible light from the 1000-W Xenon lamp. Ultra-High-Vacuum (UHV) Experiments. Experiments presented in this contribution were performed using a UHV chamber with an integrated reaction cell that was operated at environmentally relevant pressures. The facility was described in an earlier publication (13) and allowed samples to be transferred from 10-9 Torr to environmentally relevant conditions (1 bar) without exposure to the laboratory atmosphere. Samples used in the UHV experiments were natural pyrite {100} surfaces originating from Logrono, Spain. These surfaces were approximately 1 cm2 × 2 mm and were asgrown surfaces. Once introduced into the UHV environment, the surfaces were ion bombarded with 1000 eV He+ to remove carbon and oxygen impurity. Following bombardment, the samples were acid-rinsed in the reaction cell with 0.5 M HCl for 90 s. This treatment was found to generate {100} planes that were similar to those produced by mechanical fracture (14). Phosphate experiments were carried out by placing a droplet of 50 mM NaH2PO4 (maintained at pH 5) on one face of pyrite {100} for 1 h in the reaction cell, after which the sample was rinsed with 10 mL of N2-purged DI water. This water washing removed any possibility that a phosphate salt precipitate adhered to the surface. Furthermore, any phosphate left on the pyrite surface after this treatment could be presumed to be strongly adsorbing to the mineral surface. At the pH used in these studies, the predominate aqueous phosphate species was H2PO4-. For the sake of brevity, this species is referred to as phosphate in this paper. XPS data were obtained using Mg KR (1253.6 eV) unmonochromatized radiation and a double-pass cylindrical mirror analyzer (CMA). Fe 2p, S 2p, and P 2p data were analyzed with a CMA pass energy of 25 eV.
Results Aqueous Geochemical Results. Phosphate was found to be an effective retardant in near-neutral suspensions, but at pH 3 it had little or no effect. Furthermore, illumination of phosphate-containing suspensions negated any retardation. An example of the effect phosphate had on a suspension kept at pH 6 is shown in Figure 1. The pH dependence of the phosphate retardation efficiency is summarized in Figure 2a. In Figure 2, all rate data are normalized to the oxidation
FIGURE 2. (A) Relative rate of pyrite oxidation (i.e., relative to pyrite in the absence of phosphate) vs amount of phosphate sorbed on the pyrite surface. (B) Effect of illumination on oxidation rate of pyrite slurry containing phosphate. After the last addition of phosphate in experiment P7 (pH 4), the slurry was illuminated, subsequently returned to the dark, and illuminated once more. The sequence of events at the conclusion of experiment P7 is indicated with the arrows. rate in the absence of phosphate. The rate is shown as a function of the amount of phosphate sorbed, calculated by assuming that phosphatesorbed ) phosphatetotal - phosphatedissolved. It should be noted that the relative uptake of phosphate onto the pyrite surface was higher in acid suspensions than in near-neutral suspensions. In other words, the distribution of phosphate between the surface and solution was pH dependent. As seen in Figure 2a, the addition of phosphate at pH 3 had little or no effect on the oxidation rate. In experiments at higher pH there was a significant effect. While at pH 4 the retardation effect appeared to depend linearly on phosphatesorbed, the experiments at pH 5 and pH 6 showed a nonlinear dependence on phosphatesorbed. Figure 2b shows the effect of illumination of a suspension that contained phosphate and that exhibited a significant degree of retardation in the dark. In this particular experiment, the suspension was illuminated, then returned to the dark, illuminated once again, and returned to the dark for a final rate determination. The results showed that illumination increased the rate by a factor of 4-5 (close to the initial rate in the absence of phosphate). After the lamp was shut down, the rate returned to a level close to the value obtained before illumination. The subsequent cycle reaffirmed this effect. UHV Results. (a) Exposure of {100} Pyrite to Deoxygenated and Oxygenated Phosphate Solution. Clean pyrite surfaces were exposed to phosphate solutions with and without dissolved oxygen. In either circumstance, the phosphate solution was removed from the pyrite surface by VOL. 35, NO. 11, 2001 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 3. Fe 2p3/2 and P 2p data for {100} FeS2 and after this surface was exposed to a deoxygenated phosphate and oxygenated phosphate solution in separate experiments. Only in the oxygenated circumstance was a significant amount of Fe3+ (∼711 eV) bearing oxidation product and adsorbed phosphate experimentally observed. The dominate solution phosphate species at pH 5 is H2PO4-. flushing the surface with deoxygenated water. XPS of the Fe 2p core level showed that exposure of {100} pyrite to the deoxygenated solution containing phosphate had no experimentally discernible effect on the oxidation state of iron in the near-surface region of pyrite (Figure 3). Exposure of {100} pyrite to the oxygenated solution resulted in the evolution of spectral weight near 711 eV, which was assigned to Fe3+ in an iron oxide product. Prior research (11) has shown a similar effect by dissolved oxygen (although the prior work did not include phosphate). XPS of the P 2p level (Figure 3) showed that a greater amount of phosphate was present on the {100} pyrite surface after it was exposed to the oxygenated phosphate solution as compared to the deoxygenated solution (Figure 3). It can be inferred from these P 2p data that phosphate did not bind significantly on the pyrite surface under the anoxic experimental conditions. The P 2p peak maximum for the adsorbed phosphate species, after exposure of pyrite to the oxygenated solution, is ∼133.5 eV, which is consistent with P 2p values for various phosphate reference compounds published previously (15). (b) Adsorption of Phosphate on Defective {100} Pyrite. The results presented immediately above suggest that phosphate adsorption on {100} pyrite is associated with the presence of Fe3+ or Fe3+-bearing oxide phases on the mineral surface. This contention is better supported by experiments presented in this section that investigated the adsorption of phosphate on pyrite that was pretreated to increase the surface concentration of Fe3+-bearing oxide phases during exposure to oxygenated phosphate solution. These experiments were carried out by ion bombarding {100} pyrite samples with 1000 eV He+ for various times (10 or 40 min) prior to being exposed to the phosphate solution. Prior research suggested that this treatment (without subsequent annealing or acid-washing) increased the population of S2and Fe3+ surface species (11). The increased concentrations of these species led to a more significant oxidation of the 2254
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FIGURE 4. Fe 2p3/2 and P 2p data for {100} FeS2 samples with differing amounts of initial surface defects (induced by He+ bombardment), after they were exposed to a phosphate solution. After being exposed to the phosphate solution, the most defective surface showed the most iron(III) oxide and adsorbed phosphate. The clean (no ion bombardment) showed the least iron(III) oxide formation and adsorbed phosphate. iron component of the pyrite as compared to surfaces that did not receive this pretreatment. We suspect, based on arguments by Nesbitt et al. (16) for the presence of Fe3+ and S2- on mechanically cleaved pyrite, that the sulfur removal from the surface during ion bombardment leads to transient S1- species (resulting from cleavage of the disulfide bond, S22-) that oxidize neighboring Fe2+ sites to form minority Fe3+ and S2- species. Fe 2p3/2 and P 2p XPS results (Figure 4) showed that the surface concentration of phosphate that remained as an insoluble surface product was a function of the amount of Fe3+-bearing oxidation product. In particular, as the amount of Fe3+-bearing oxide product was increased on pyrite {100} (by increasing the ion bombardment pretreatment), there was a concomitant increase in the amount of surface-bound phosphate. In separate experiments, ion-bombarded pretreated surfaces were exposed to oxygenated H2O but without phosphate. Fe 2p of these surfaces after exposure to the oxygenated H2O showed iron(III) oxide surface concentrations that were similar to the corresponding pretreated pyrite {100} surfaces in the presence of phosphate (shown in Figure 4). These results suggest that the presence of phosphate does not affect the formation of the Fe3+-bearing oxide product but that the presence of the oxide phase is essential for its adsorption. (c) Oxidation of {100} Pyrite in Gaseous H2O/O2 with and without Surface Phosphate. Clean {100} FeS2 and FeS2 with adsorbed phosphate were individually exposed to a gaseous mixture of H2O (0.02 bar) and O2 (1 bar) for 20 h at 298 K. The gaseous O2/H2O environment was used,since the sulfur oxidation products (e.g., SO42-) could be maintained on the surface for investigation. In the aqueous environment, SO42- is rapidly removed from the surface. The phosphate/
FIGURE 5. Fe 2p data for FeS2{100} and after it was exposed to gaseous O2 and H2O (top), and FeS2{100} with adsorbed phosphate before and after exposure to gaseous O2 and H2O (bottom). The presence of phosphate significantly reduced the amount of Fe3+ oxidation product as compared to the phosphate-free surface. FeS2 surface was prepared by exposing a clean pyrite surface to an oxygenated solution of phosphate and then rinsing the surface with deoxygenated H2O (with subsequent drying in a pure N2 stream). Fe 2p XPS data (Figure 5) show that the amount of Fe3+-bearing oxidation product after exposure to the H2O/O2 mixture was significantly reduced when phosphate was preadsorbed as compared to pyrite that had no adsorbed phosphate. The Fe 2p data showed that the amount of Fe3+ oxidation product on the phosphate-bearing surface after exposure to the gaseous O2/H2O(g) environment was very similar to the initial concentration of oxide product (prior to exposure). These results suggested that the adsorbed phosphate was extremely effective in inhibiting oxidation of the iron component of pyrite in the O2/H2O(g) environment. Complimentary S 2p data (Figure 6) showed that the amount of sulfur oxidation was also significantly reduced if phosphate was adsorbed prior to exposure to O2/H2O(g). Specifically, the concentration of S6+-bearing product (presumably SO42-) was a factor of 3 greater on the pyrite surface that had no adsorbed phosphate.
Discussion The transfer of electrons from Fe2+ sites of pyrite to molecular oxygen is generally accepted to be a key step in pyrite oxidation in oxic aqueous or gaseous environments. The specifics of this step, however, are quite complicated. Moses and co-workers (4) have investigated pyrite oxidation in the presence of aqueous Fe3+ and proposed that adsorbed Fe3+ acts as an electron conduit for the transfer of electrons from Fe2+ to molecular oxygen (i.e., formation of the superoxide, O2-). Eggleston et al. (17) have suggested that a similar scenario occurs during the oxidation of pyrite in gaseous O2 and H2O. In the gaseous circumstance, Fe3+ bound as oxide product on the pyrite surface facilitates the electron-transfer
FIGURE 6. S 2p data for clean FeS2, after it was exposed to gaseous O2 and H2O, and after FeS2 with adsorbed phosphate was exposed to gaseous O2 and H2O. Less S6+-bearing product (presumably SO42-) was produced when phosphate was present on the pyrite surface. step from Fe2+ to O2. It might be suspected that pyrite oxidation in the aqueous environment, with or without sorbed aqueous Fe3+, also is facilitated by electron transfer at Fe3+bearing oxide sites. Our results for the effect of phosphate on pyrite oxidation lend support to the above model of oxidation where Fe3+bearing oxide facilitates the oxidation process. XPS results indicate that the phosphate anion binds preferentially and irreversibly to the Fe3+-bearing oxidation phases at a pH near 5 on the pyrite surface. These experimental observations are in agreement with prior geochemical studies that show that phosphate binds strongly to iron(III) hydroxides and oxides (18) but interacts weakly with bare pyrite surface (19). While the XPS experiments show that no phosphate sorbs on the bare pyrite surface, the aqueous geochemistry studies show that phosphate sorbs onto the pyrite surface at pH 3. This difference is probably due to the presence of dissolved iron in solution in the batch experiments. Earlier work has shown that the presence of iron renders the pyrite surface a positive charge (12), which facilitates the uptake of anions, including phosphate (19). Over the pH range studied here, the dominant species in solution is H2PO4-. However, with the techniques used in this study, it is not possible to determine the identity and/or geometry of the adsorbed species. Nevertheless, phosphate preferentially bound to the Fe3+-bearing oxide significantly inhibits pyrite oxidation under our experimental conditions. This experimental observation alone strongly supports the notion that iron(III) oxide phases are intimately involved with the elementary steps of pyrite oxidation at pH values near 5. At least two mechanisms appear plausible to explain phosphate inhibition. Adsorbed phosphate may act as a site blocker for molecular oxygen adsorption on the iron(III) oxide phases, limiting electron transfer (Figure 7) to the oxidizing agent, or adsorbed phosphate may modify the surface VOL. 35, NO. 11, 2001 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 1. Summary of Oxidation Ratesa expt
pH
rate (no PO4)
[PO4]addedb
[PO4]sorbedc
rate (PO4, max)d
P17 P7 P14 P9 P19
3 4 5 6 6
0.22 3.97 5.87 2.32 2.29
304 600 254 800 213
262 490 109 127 59
0.21 1.21 2.18 0.39 0.81
a Rates in 10-9 mol m-2 s-1; PO concentrations in µmol. b Maximum 4 amount of PO4 added to suspension. c Sorption achieved at maximum d amount of PO4 added. Oxidation rate at the maximum concentration of PO4.
FIGURE 7. Proposed models for the effect of phosphate on the oxidation of pyrite. Without adsorbed phosphate, Fe3+-bearing oxide phases serve as conduits for the transfer of electrons from Fe2+ in pyrite to adsorbed molecular O2 . Phosphate may physically inhibit the adsorption of molecular O2 or alter the electronic structure of the iron(III) oxide phase, making electron transfer less energetically favorable. electronic structure. Eggleston et al. (17) proposed that the band structure of the iron oxide (assumed to be hematite, Fe2O3) allowed it to serve as an electron conduit; hence, phosphate bound on iron(III) oxide phases may alter the band structure enough to increase the barrier for electron transport. While it is not possible to determine if phosphate acts as a site blocker or as an electronic modifier on the basis of our experimental results, the illumination experimental results tend to support the latter possibility. Illumination of the phosphate /pyrite system at pH 5, for example, results in a return of the pyrite oxidation rate to a level consistent with the phosphate-free system. Prior studies show that illumination of pyrite in the absence of phosphate only lead to minor changes in reaction rate (20). Results in this study show that the illuminated phosphate/pyrite and dark pyrite oxidation rate is similar, suggesting that the illumination activates regions of the surface that were originally nonactive in the dark with phosphate. It may be that the promotion of an electron from the pyrite valence band to the conduction band allows the previously unreactive surface to participate in the oxidation process. This photoinduced process may compensate for the modified electronic structure of the Fe3+ oxidation product by the adsorbed phosphate. This study shows that the inhibition effect of phosphate is significant at pH >3.5, but below this value, this species does not affect the oxidation rate. The overall oxidation rate of pyrite, however, decreases with a reduction in pH (Table 1 summarizes the absolute rates as a function of pH), an effect that is observed in prior experiments without phosphate (20). Below pH 3.5 the Fe3+-bearing oxidation product is soluble, thus eliminating the binding of phosphate and its inhibition of the oxidation rate. While the changes in pyrite oxidation rate with pH might be expected to be a function of many factors, it is possible that the decrease in oxidation rate below pH 3.5 is due to the absence of surface-bound Fe3+ phases that can facilitate the oxidation of pyrite-Fe2+ sites. Also, prior research (21) has shown that the oxidation rate of aqueous Fe2+ by O2 decreases with pH. Whether this result can be directly related to pyrite oxidation is not known. 2256
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It does suggest that changes in pyrite oxidation may also be due to changes in the Fe2+ to Fe3+ couple as the pH is altered. Finally, we mention that although the effect of phosphate on the oxidation kinetics of pyrite oxidation was significant, in no experiment was a complete inhibition observed. There are at least two possible reasons for this experimental observation. First, the phosphate concentration might have been too low to cover all the iron(III) oxide regions that initially appear on the surface at the onset of oxidation. This possibility cannot be excluded, but the data in Figure 2a for experiments at pH 5 and pH 6 suggest that with more phosphate there would have been little or no further retardation. An alternative and favored explanation is that the oxidation of pyrite occurs through at least two pathways; one pathway involving electron transfer via the iron(III) oxide regions and a second pathway involving reactive sites on the bare surface (i.e., oxide-free). Since the phosphate interacts only weakly with unoxidized pyrite surfaces, the process can continue, albeit slower on these sites. On the basis of this contention, the overall oxidation rate determined at pH 3 or below will be presumably dominated by reactions on primarily oxide-free (no adsorbed phosphate) pyrite. At higher pH, in the absence of phosphate, the contribution of the oxide patches will dominate the overall rate. In the presence of phosphate, however, the reactivities of these sites are significantly suppressed, and the oxidation of pyrite will again be dictated by the electron-transfer rate via the bare surface.
Summary and Implications Results of this study are best explained by a model in which phosphate interacts exclusively with iron(III) oxide regions that are present on pyrite surfaces in solutions with pH >3.5. Sorption onto these regions hinders the electron transfer and decreases the oxidation rate. In acidic solutions, the surface oxide regions are no longer stable, and phosphate has little or no effect on the oxidation rate. The implication is that below pH 3, addition of retardants that work by hindering electron transfer via the iron(III) oxide regions will be ineffective, unless the retardant also binds to reactive sites on the bare surface or forms a protective coating. The formation of an encapsulating phosphate coating requires far higher concentration than used in this study. Given that in many mine-waste environments the pH is well below 3 and sometimes negative (22), an Fe3+-bound retardant may not work. Results presented in this paper show that the inhibitory effect of sorbed phosphate on pyrite oxidation is largely removed when the surface was exposed to light. With a significant amount of mine waste spread out along stream beds and holding ponds, it is important to further study the effect of light on the efficiency of possible retardants.
Acknowledgments D.R.S. and M.A.A.S. greatly appreciate support from the Department of Energy, Basic Energy Sciences, from Grants DEFG0296ER14644 and DEFG029ER14633, respectively. Three
anonymous reviewers are thanked for their valuable comments and suggestions.
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(13) Guevremont, J. M.; Elsetinow, A. R.; Strongin, D. R.; Bebie´, J.; Schoonen, M. A. A. Am. Mineral. 1998, 83, 1353-1356. (14) Elsetinow, A. R.; Guevremont, J. M.; Strongin, D. R.; Schoonen, M. A. A. Am. Mineral. 2000, 85, 623-626. (15) Briggs, D.; Seah, M. P. Practical Surface Analysis by Auger and X-ray Photoelectron Spectroscopy; John Wiley & Sons: New York, 1983. (16) Nesbitt, H. W.; Bancroft, G. M.; Pratt, A. R.; Scaini, M. J. Am. Mineral. 1998, 83, 1067-1076. (17) Eggleston, C. M.; Ehrhardt, J.; Stumm, W. A. M. Am. Mineral. 1996, 81, 1036-1056. (18) Davies, J. A.; Kent, D. B. Surface Complexation Modeling in Aqueous Geochemistry; Mineralogical Society of America: Madison, WI, 1990; Vol. 23. (19) Bebie, J.; Schoonen, M. A. A. Earth Planet. Sci. Lett. 1999, 171, 1-5. (20) Schoonen, M.; Elsetinow, A.; Borda, M.; Strongin, D. Geochem. Trans. 2000, 4. (21) Stumm, W.; Morgan, J. J. Aquatic Chemistry; Wiley-Interscience: New York, 1981. (22) Nordstrom, D. K.; Alpers, C. N.; Ptacek, C. J.; Blowes, D. W. Environ. Sci. Technol. 2000, 34, 254-258.
Received for review September 19, 2000. Revised manuscript received March 12, 2001. Accepted March 14, 2001. ES0016809
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