Environ. Sci. Technol. 2000, 34, 1494-1499
Influence of Cadmium Sorption on FeS2 Oxidation B E N J A M I N C . B O S T I C K , * ,†,‡ S C O T T F E N D O R F , †,‡ BRYAN T. BOWIE,§ AND PETER R. GRIFFITHS§ Soil Science Division, University of Idaho, Moscow, Idaho 83844-2339, and Department of Chemistry, University of Idaho, Moscow, Idaho 83844
Pyrite oxidation leads to the formation of acid-mine drainage and the release of associated trace metals. A better understanding of the processes that influence pyrite oxidation will help to determine the rate of acid and cation evolution from pyritic mine tailings. The oxidation process is surface-controlled and may be influenced by cadmium sorption, which forms surface precipitates and complexes that may limit pyrite oxidation. The purpose of this research is to investigate the effect of cadmium sorption on the rate of FeS2 oxidation by molecular oxygen. Raman spectroscopy was used to track the evolution of oxidized sulfur products, and X-ray absorption spectroscopy was used to quantify iron oxidation rates. Cadmium concentrations as low as 50 µM depressed pyrite oxidation rates. Oxygenation of amorphous FeS2 (1 g/L) is described by a pseudo-first-order reaction with a rate constant of 6.85 × 10-5 s-1; FeS2 sorbed with 500 µmol of Cd g-1 prior to oxygenation exhibited an oxidation rate 5.15-fold lower with a rate constant of 1.33 × 10-5 s-1. Raman and XANES data indicate that cadmium sorption influences FeS2 oxidation. In natural systems, suppression of FeS2 oxidation by Cd and other soft Lewis acids may retard the release of acidity and trace metals to the environment.
Introduction Pyrite (FeS2) is found in mining areas, aquifers, and lake and estuarine sediments. Pyrite oxidation acidifies surface waters, potentially releasing trace elements such as arsenic, lead, zinc, and cadmium that are associated with sulfidic ores, soils, and sediments (1-3). Thus, pyrite oxidation is a pervasive environmental concern. The rate of pyrite oxidation largely determines the amount of oxidation and ultimately the amount of acidification and trace metal released into the environment. Consequently, in pyrite-containing systems, processes that control the release of trace elements as well as conditions that influence these processes are of interest. The initial step in pyrite oxidation involves the transition of sulfur to thiosulfate, releasing Fe2+ into solution where it may be oxidized by molecular oxygen (4). Ferric iron can precipitate (reaction 1) to form a hydrous iron oxide at high pH (4) or further oxidize (reaction 2) the pyrite at neutral or lower pH (2, 5, 6): * Corresponding author phone: (650)723-4152; e-mail: bbostick@ stanford.edu; fax: (650)725-0979. † Soil Science Division. ‡ Present address: Department of Geological and Environmental Sciences, Stanford University, Stanford, CA 94305-2115. § Department of Chemistry. 1494
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FeS2 +
15 7 O + H O S Fe(OH)3 + 2SO42- + 4H+ (1) 4 2 2 2
FeS2 + 14Fe3+ + 8H2O S 15Fe2+ + 2SO42- + 16H+ (2) Often pyrite oxidation occurs nonstoichiometrically or incompletely with many intermediate oxidation products including elemental sulfur, thiosulfate, and polysulfides (1, 7). Ferrous iron may sorb on the pyrite surface, blocking the active oxidation sites from either dissolved oxygen or ferric iron (8, 9). If Fe2+ can block reactive surface sites, other sorbed ions may also influence the oxidative behavior of pyrite. However, ionic species of other elements, including V, Co, Ni, Mn, Na, and Ce, do not appear to influence pyrite oxidation (1, 10, 11). Oxidants that cause iron (hydr)oxide precipitation may also passivate pyrite surfaces (12-14). Elemental sulfur production during oxidation may similarly suppress pyrite oxidation (17). Cadmium reacts with the surface of FeS2 through a complex disproportionation reaction, which leads to the formation of cadmium sulfide, oxidized iron products, and elemental sulfur (16, 17). Cadmium sorption on pyrite should thus lead to the formation of surface species that inhibit pyrite oxidation. To test this premise, we studied the oxidation rate of FeS2 reacted with cadmium.
Materials and Methods FeS2 Preparation. Fresh FeS2 was synthesized from iron hydroxide and dissolved sulfide before each experiment (18). The FeS2 suspension, and all subsequent reactions with Cd, was kept under nitrogen during synthesis and thereafter to prevent oxidation. Iron(III) was initially precipitated at pH 4 from 1 L of a 0.034 M FeCl3‚6H2O solution. Then 0.068 mol of NaHS‚H2O was slowly added. The pH was decreased to 4 with 0.25 M HCl, and the suspension was allowed to react for approximately 2 h, after which time NaCl was added to flocculate the solids. The suspension was then centrifuged and washed in distilled water until sulfide was not detected using an ion-selective electrode. The resulting suspension had a suspension density of 3.5 g/L. The stoichiometry of the product was verified to be FeS2 by quantifying the residual Fe and S in solution as well as by acid digestion of the solids. Batch Reactions. Reaction Conditions. Batch experiments were used to characterize the sorption behavior of the amorphous FeS2. Samples were prepared by mixing the stock suspension of FeS2 (3.5 g/L) with distilled water to create a 1 g/L suspension density (about 4.17 × 10-3 M FeS2). The suspension was then adjusted to the desired pH, and Cd was added from a 50 mM stock solution made with CdBr2‚2H2O. This mixture was then readjusted to the proper pH and kept in a N2(g)-filled glovebox for approximately 1 h to allow for complete sorption. Cadmium concentrations of 0, 10, 50, and 5000 µM (0, 10, 49, and 970 µmol/g, respectively) at pH values of 5, 7, and 9 were used. Raman Spectroscopy. The solids were collected by filtration and kept in an anaerobic chamber until immediately prior to oxidation. A thin film of the solids was deposited on a gold mirror and placed in air (0.2 atm of O2) at ambient humidity for Raman spectroscopic analysis during the ensuing oxidation. No control of water activity was made during this experiment, although water activity can influence oxidation (e.g., ref 19). Spectra were collected with a Renishaw Ramascope Raman microscope, equipped with a 782-nm diode laser operating at 5-mW average power, using the 50× objective. 10.1021/es990742w CCC: $19.00
2000 American Chemical Society Published on Web 03/09/2000
FIGURE 1. Experimental reactor used for XANES data collection (adapted from ref 23). Spectra were collected between 140 and 1000 cm-1 Raman shift at 6 cm-1 resolution. Three spectra were co-added, each with 50 s/point integration times. Spectra were collected at regular intervals from the onset of oxidation; each spectrum was collected in a different area to prevent sample photodegradation. Standard spectra were collected for a range of potential reaction products including elemental sulfur, ferrihydrite (Fe5HO8‚5H2O), goethite (R-FeOOH), hematite (R-Fe2O3), lepidocrosite (γ-FeOOH), greenockite (R-CdS), cadmium hydroxide (γ-Cd(OH)2), various iron sulfates, and sodium sulfite. Bands were also compared to published spectra (20, 21). For the analysis of polysulfides and iron (hydr)oxides, cyclohexane was used to dissolve elemental sulfur prior to spectral collection. Kinetic Experiments. Reaction Conditions. A suspension density of 1 g/L synthetic FeS2 in 0.01 M 1,4-piperazinediethanesulfonic acid (PIPES) buffer was placed into a sealed 150-mL reaction vessel. The reaction vessel (Figure 1) contained a gas inlet, outlet, and sampling ports (22, 23). Suspension mixing was accomplished by stirring with a magnetic stir bar. The pH was adjusted to 6.5 with 0.1 M NaOH. The system pH was monitored before and after the reaction; it changed less than 0.05 pH unit. Samples containing sorbed Cd were created immediately before analysis in order to prevent oxidation. Cadmium, as a 50 mM CdBr2‚2H2O stock solution, was added to the suspension to obtain a final concentration of 500 µM total Cd. No soluble Cd was detectable in the residual solution. The system was kept purged with nitrogen to prevent oxidation prior to oxygenation measurements. Oxidation of FeS2 was carried out by continuous gas injection through a glass frit at a flow rate of 85 mL/min. The partial pressure of O2 was kept constant at 0.2 atm. Reacted suspension was then pumped into a flow cell consisting of polyethylene tubing in the X-ray beam path. Fresh sample from the primary reaction vessel entered the flow cell for each analysis. XANES Spectroscopy. XANES spectra were collected on beamline 4-3 at the Stanford Synchrotron Radiation Laboratory (SSRL). A Si(111) monochromator was used to scan the incident X-ray beam through the K-edge of ironsfrom 7050 to 7320 eV. Incident and transmitted intensities were measured with N2-filled ionization chambers; fluorescence intensity was measured using a 13-element Ge detector. Internal energy calibration for each sample was accomplished using an iron foil positioned between the second and third in-line ionization chambers, with the first inflection point set to 7112 eV. Because of the dilute nature of the sample suspension, fluorescence data were used for analysis. Spectra were collected periodically during the reaction until oxidation was visibly apparent by the presence of a bright orange suspension. Each scan took approximately 3 min to collect,
and the midpoint of the scan was assumed to be the average time for a measurement. Data analysis was carried out using Peak-Fit 4.0 (Jandel Scientific). The background was subtracted, the jump height was normalized to unity, and the spectrum smoothed (5%) using a first derivative Savitzky-Golay method. The resulting first derivative XANES spectrum was then fit using Voigt amplitude functions. These fitted peaks were used to quantify the fraction of Fe(II) and Fe(III) in the samples by comparison with a standard curve generated from mixtures of FeOOH and FeS (22, 23). Independent data analyses of samples by the curve-fitting procedure resulted in iron oxidation states within 4%. The fitting is based on the shift in binding energy of the Fe K-shell (1s orbital) as a function of oxidation state. The higher oxidation state, Fe(III), has a larger effective nuclear charge and thus has a slightly larger binding energy. While the pre-edge region has also been used to determine Fe oxidation states (e.g., ref 24), these features are much weaker than main-edge components and thus result in less favorable signal-to-noise ratios for monitoring changes in Fe(II):Fe(III) ratios.
Results and Discussion FeS2 Characterization. The FeS2 suspension was predominantly X-ray amorphous and had a framboidal morphology as noted by scanning electron microscopy. No oxidized sulfur products were detected using Raman spectroscopy. The synthetic FeS2 had a surface area of approximately 125 m2/g based on ethylene glycol monoethyl ether adsorption (25). Raman spectra of the synthetic FeS2 have very broad peaks, shifted approximately 20 cm-1 to lower energies relative to crystalline pyrite, indicating that the FeS2 is disordered. Raman Spectroscopy. The extent of sorption was determined by measuring residual Cd in solution following reaction with FeS2. Greater than 95% of the Cd initially added sorbed for all samples except for those containing a 5000 µM initial Cd concentration. For this sample, the surface loading was 970 µmol of Cd/g (approximately 20% of the Cd added partitioned to the solid phase). The amorphous FeS2 sorbed Cd more strongly and at higher surface concentrations than ground mineral pyrite (16, 17), possibly due to the presence of surface defects in the precipitated iron disulfide. Without the addition of Cd to the FeS2 samples, bands attributable to elemental sulfur are not visible in the Raman spectrum (Figure 2A). As Cd is added, the magnitude of the elemental sulfur scattering peaks increase, indicating that this phase formed as a product of Cd sorption. Sulfur-rich surface layers must be considered in analyzing data from subsequent intervals in order to separate the effects of oxidation induced by molecular oxygen versus those induced by Cd sorption. Oxidation of FeS2 is rapid, with elemental sulfur present within minutes of oxygenation (Figure 2B). Significant oxidation occurred for all samples except for the 970 µmol g-1 sample. Increasing elemental sulfur is mirrored by the evolution of iron (hydr)oxides resulting from the concurrent oxidation of iron. Lepidocrosite and elemental sulfur are produced after 255 min of oxygenation, with ferrihydrite being detected in small amounts in a few samples. There are large spatial variations in ferrihydrite and lepidocrosite, with the latter being more prevalent and well distributed. It is interesting to note that the broad iron oxide Raman bands at 300-400 cm-1, attributed to ferrihydrite, present after 40 min change into sharper bands characteristic of lepidocrosite when oxidized for 255 min or longer. The development of sharper bands can be explained by the development of more stable, crystalline species over time. The transformation of ferrihydrite to lepidocrosite is unlikely because ferrihydrite must dissolve to form lepidocrosite (26). Therefore, the two species probably form independently. The VOL. 34, NO. 8, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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spectral intensities. However, the general trend of decreasing sulfur spectral intensities with increased Cd loading was reproducible, indicating that Cd sorption did suppress the formation of elemental sulfur. After 24 h, oxidation of FeS2 without Cd is essentially complete in oxygenated suspensions. Elemental sulfur is lacking due to its conversion to sulfate, as detected by Raman and confirmed by the presence of hydrous iron(III) sulfate, detected by infrared spectroscopy; formation of sulfate is necessary for stoichiometric oxidation (1). Cd-sorbed FeS2 samples contain large concentrations of elemental sulfur following oxygenation for 24 h, indicating that oxidation is suppressed. This observation is in contrast with the shrinking core kinetic model for pyrite oxidation (28), in which oxidation products do not interfere with the reactive surface.
FIGURE 2. Raman spectra of FeS2 with various cadmium loadings at pH 7. (A) Raman spectra of Cd-loaded pyrite prior to oxidation indicate the formation of elemental sulfur due to Cd sorption (peaks at 222, 473, 245, and 438 cm-1). (B) Normalized Raman spectra of FeS2 during oxidation show the development of iron oxides and elemental sulfur. At 40 min, ferrihydrite is present, as indicated by the broad peaks at 300-400 cm-1. After 4.25 h, lepidocrosite is present, indicated by peaks at 249 and 378 cm-1. (C) Raman spectra of FeS2 samples oxidized for 5 h at pH 7. (D) Raman spectra of hexane-cleaned samples following 24 h of oxidation. formation of ferrihydrite would be most significant initially, while lepidocrosite predominates at longer times. This behavior is expected because lepidocrosite is more crystalline and thus has a higher energy of nucleation, causing it to form more slowly (26). The surface of FeS2 may catalyze the precipitation of ferrihydrite relative to lepidocrosite. Anion sorption, including sulfate formed during pyrite oxidation, may also influence which iron (hydr)oxide precipitates (e.g., refs 22 and 27). Further research is needed to determine which of these processes are operative. As the sample dries, the oxidation rate slows, consistent with other studies (e.g., refs 1 and 19). However, the trend of increasing elemental sulfur with time generally continues for at least 8 h. A band attributable to polysulfide at 454 cm-1 is noted after oxidation for 255 min. However, polysulfide was generally not detected after long-term oxidation. After extended periods of oxidation (∼5 h), the effect of Cd addition on oxidation is readily apparent (Figure 2C). With one exception (the 10 µmol of Cd g-1 loading), samples with cadmium sorbed form less elemental sulfur than FeS2 alone. The 10 µmol of Cd g-1 sample shows an overall increase in elemental sulfur resulting from Cd sorption and FeS2 oxidation by molecular oxygen. At this low loading, passivation of FeS2 is not appreciable, and the net result of cadmium sorption is enhanced elemental sulfur formation. At 50 µmol of Cd g-1 loading, elemental sulfur is produced due to oxygenation, although less than for FeS2 controls. The samples with 970 µmol of Cd g-1 loading have very little net production of elemental sulfur from oxygenation. Thus, for cadmium loadings of 50 µmol g-1 and higher, oxidation from exposure to oxygenated waters is retarded. Attempts to subtract elemental sulfur produced by cadmium sorption were not possible because of the difficulty in reproducing 1496
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The system pH had a large effect on the products of oxidation for Cd-sorbed FeS2 at a surface concentration of 50 µmol of Cd g-1 (Figure 2D). Samples were washed with cyclohexane in order to remove elemental sulfur, allowing the residual iron (hydr)oxides to be studied in more detail. At pH 5 and pH 7, the dominant iron oxyhydroxide formed was lepidocrosite. Lepidocrosite also forms as the dominant oxidation product from amorphous iron sulfide (22) and commonly forms as the initial oxidized product of iron(II) oxidation (26). At pH 9, substantial ferrihydrite was present while lepidocrosite was at a much lower concentration. The presence of various iron (hydr)oxides indicates that reaction mechanisms and oxidation retardation effects of sorbates will vary with pH. Under most of the experimental conditions studied, lepidocrosite was formed as the predominate iron (hydr)oxide. This is contrary to previous research (14, 29-31) which has determined that ferrihydrite is the product of oxidation. However, Nicholson et al. (12) noted that maghemite (γFe2O3), a dehydrated lepidocrosite, formed from pyrite oxidation. In addition, iron oxyhydroxides were suggested to be an important product of pyrite (19) and arsenopyrite (32) oxidation in aerated water. Oxidation Kinetics Measured with XANES Spectroscopy. Significant variability in the excitation volume, dispersion, and other instrumental factors prevent the collection of quantitative Raman spectra. Although still valuable for qualitative to semiquantitative measurement, this difficulty is compounded by the fact that FeS2 corrosion occurs in a very heterogeneous fashion. Therefore, XANES spectroscopy was used to determine reaction rates for the oxidation of FeS2 and of Cd-sorbed FeS2. Using the pulsed flow technique described here, XANES spectroscopy provides a useful means of determining oxidation rates quantitatively. XANES spectra of FeS2 oxidized for varying lengths of time are illustrated in Figure 3. The spectra are described by a multitude of Voigt peaks; the pre-edge contains two peaks at approximately 7112 and 7115 eV that shift slightly with increased oxidation, while the main-edge portion of the spectrum is described by five additional peaks. Peaks at 7123, 7129, and 7132 eV generally maintain constant height and position throughout the reaction, indicating that these peaks are not likely associated with edge shifts induced by oxidation. Peaks at 7125 and 7127 eV, however, change in amplitude with time and can be used to quantify the contributions of Fe(II) and Fe(III). The 7125 eV peak, which decreases with time, is assigned to Fe(II) while the 7127 eV peak, which increased in intensity overtime, is identified as Fe(III). Quantifying Fe(II) and Fe(III) was based on peaks area ratios. Peak ratios have also been used to determine the fractions of Cr(VI) and Cr(III) (33) and Mn(IV) and Mn(II) (34). Generally, peak areas provided superior quantification because they can compensate for peak asymmetry, peak
FIGURE 4. XANES spectra of FeS2 (A) and FeS2 sorbed with 500 µmol of Cd/g (B) as a function of oxidation duration. The shift in the edge for FeS2 indicates a large change in oxidation state; no change is apparent for Cd-sorbed FeS2.
TABLE 1. Fe(II) and Fe(III) Fractions during the Oxidation of Pyrite and Cd-Sorbed (500 µmol of Cd g-1) Pyrite at pH 7 Determined by XANES Spectroscopya pyrite
Cd-sorbed pyrite
FIGURE 3. Normalized XANES spectra of oxidized FeS2 at pH 6.5 during oxidation and the Fe(II) and Fe(III) components used in fitting.
t (min)
Fe(II) (%)
Fe(III) (%)
t (min)
Fe(II) (%)
Fe(III) (%)
1 7 13 21.5 50.5 81 117 255
13.3 (7.3)b 11.2 16.0 16.3 21.3 27.0 40.0 69.3
1
broadening, and different Gaussian and Lorentzian contributions to the Voigt function used in fitting. The XANES spectra of FeS2 indicate a large shift in the intensity of the Fe(II) and Fe(III) peaks during oxygenation and suggest that most FeS2 was exhausted within a few hours (Figure 4). Extensive oxidation was confirmed by the color change from black to orange observed in the reaction vessel. Initially, the iron in FeS2 was 93% Fe(II) while it decreased to 31% Fe(II) after reaction with oxygenated water for 255 min (Table 1). Some of the ferric iron may be due to traces of oxidation; however, oxidized iron species may also form on FeS2 surfaces prior to exposure to O2 (35). The Cd-FeS2 samples are characterized by the presence of a large fraction (33%) of Fe(III) initially, thus Cd sorption causes the net oxidation of pyritic iron. Raman data also show the presence of oxidized sulfur and iron products formed during Cd sorption. However, the Cd-FeS2 samples oxidize more slowly than the FeS2 without Cd (Figure 4). Over 145 min, the Fe(III) fraction increases by only 9%sfrom 33 to 42% (Table 1). The oxidation of FeS2 by molecular oxygen is thought to be first-order with respect to pyrite (1, 8, 28, 36). The pseudofirst-order rate constant for native FeS2 is 6.85 × 10-5 s-1 for 1 g/L FeS2 in air-saturated water at pH 6.5 (Figure 5). This rate can be compared to the general rate equation determined for pyrite oxidation by molecular oxygen (9):
86.7 (92.7)b 88.8 84.0 83.7 78.7 73.0 60.0 30.7
66.9 (65.4)b 59.6 63.6 62.8 61.9 60.6 60.6 58.3
33.1 (34.6)b 40.4 36.4 37.2 38.1 39.4 39.4 41.7
[O2]0.5 aq r ) 10-8.19 + 0.11 [H ]
(3)
9 22.5 40.5 59 77 118 145.5
a Percentages were determined with peak areas and were reproducible within 4%. Regression determined that the 95% confidence interval for each point was e6%. Note the high fraction of Fe(III) initially in pyrite with sorbed cadmium and the inhibition of oxidation in the Cdsorbed pyrite. b Determined using the intercepts in first-order kinetic plots.
where r is the rate in mol m-2 s-1. Using eq 3, the predicted rate is 6.18 × 10-10 mol m-2 s-1. The pseudo-first-order equation can be adjusted to the units of eq 3 using the following relation:
r)
k′[FeS2] A
(4)
where k′ is the pseudo-first-order rate constant (in s-1), [FeS2] is 0.00833 M (1 g/L), and A is the surface area of synthetic FeS2 (in m2/L). Given a 125 m2/g surface area for synthetic FeS2, the experimental rate, r, is 4.56 × 10-9 mol m-2 s-1. The experimental rate is slightly higher, but within an order of magnitude, than the calculated values. Experimental error may explain some of the variation in the observed rate from VOL. 34, NO. 8, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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Iron disulfide oxidation is influenced by the sorption of cadmium. Generally, oxidation proceeds quickly at circumneutral pH and results in the formation of iron (hydr)oxides, elemental sulfur, and sulfates. Even low surface coverages of cadmium induce a noteable decrease in oxidation rates, which may be important in natural systems. Trace element concentrations in acid-mine drainage are commonly near 10-50 µM (43), the lower limit of Cd concentrations that retard pyrite oxidation in this study. Additionally, these systems often contain higher concentrations of Cd in the solids, which may increase the effect of Cd on pyrite oxidation. It is also possible that other soft Lewis acids, such as Pb2+ and Zn2+, may also react similarly with FeS2 and in combination of these elements could significantly retard pyrite oxidation. FIGURE 5. First-order kinetic plots of the oxidation of 1 g/L of FeS2 and 500 µmol/g of Cd on pyrite at pH 6.5. that of previous research, but an additional explanation is appropriate. One possible explanation for the higher rate constant is the higher reactivity of synthetic FeS2 relative to natural pyrite (4, 37, 38). Observed rates of FeS2 oxidation may also be due to the method of detection. XANES spectroscopy detects a change in oxidation state of all iron in a suspension. Most previous research has determined the extent of pyrite oxidation by measuring the stoichiometric evolution of (soluble) iron (36, 37) or sulfate (8, 39). Pyrite oxidation is often nonstoichiometric (7, 40), so iron oxides and sulfur species, including elemental sulfur, are not measured in solution-phase experiments. Pyrite oxidation does produce ferrous iron during the oxidation process (equation 2). However, at the neutral pH used in this study, the final product of pyrite oxidation is ferric oxyhydroxide (equation 1). Thus, the net reaction can be described by the formation of ferric iron. The rate reported here is the convolution of the rate of both pyrite and soluble ferrous iron oxidation. However, the rate of Fe(II) oxidation at pH 7 is very rapid and is not likely to be rate-limiting (e.g., refs 1 and 8). Therefore, in effect, the surface reaction (the net oxidation of FeS2) is being measured. This hypothesis is supported by two experimental factors: the rate of oxidation is comparable to the rate of pyrite oxidation reported by others, and the rate is affected by the sorption of Cd to the surface, implying that a surface reaction is rate-limiting. The pseudo-first-order rate constant of Cd-FeS2 (500 µmol g-1) is 1.33 × 10-5 s-1 (Figure 5), which corresponds to a ∼5-fold decrease relative to FeS2 without Cd. Assuming that the decrease in the rate constant is due to a change in reactive surface area, the consumed surface area (100.7 m2/g) has sorbed 500 µmol of Cd. This suggests a site density of 4.96 µmol of Cd/m2, similar to other reported values (41). Raman data suggest that Cd concentrations as low as 50 µM may decrease the rate of FeS2 oxidation, a concentration too low to appreciably affect the total surface area. The reactive surface area can, however, be significantly smaller than the total surface area (37, 38). Therefore, even if Cd only covers a small portion of the surface, a large decrease in the oxidation rate may be observed. Furthermore, Cd sorption involves the formation of CdS, elemental sulfur, and other species, all of which would cover more reactive surface area than Cd alone, thus enhancing passivation. Cadmium preferentially reacts with defect sites on the FeS2 surface (17), which are typically the most reactive sites toward oxidation (37, 38, 42). Once all of the reactive sites are occupied, additional sorbed Cd would have minimal effect on the oxidation rate. This mechanism of oxidation suppression is supported by the nonlinear trend in oxidation rates observed in the Raman data. 1498
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Acknowledgments Funding for this research was provided by the NSF-Idaho EPSCoR program and the National Science Foundation (Grant OSR-930539). The authors thank three anonymous reviewers for their helpful suggestions.
Supporting Information Available A table containing the Raman peaks for several iron (hydr)oxides, sulfides, and polysulfide standards (2 pages). This material is available free of charge via the Internet at http://pubs.acs.org.
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Received for review July 1, 1999. Revised manuscript received January 24, 2000. Accepted January 25, 2000. ES990742W
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