Environ. Sci. Technol. 1998, 32, 2084-2091
Potential Role of Bicarbonate during Pyrite Oxidation V . P . E V A N G E L O U , * ,† A . K . S E T A , ‡ A N D A. HOLT† Department of Agronomy, University of Kentucky, Lexington, Kentucky 40546-0091, and Fakultas Pertanian, Universitas Bengkulu, Bengkulu 38371-A, Indonesia
According to Frontier molecular orbital (FMO) theory, the surface-exposed sulfur atom of pyrite possesses an unshared electron pair which produces a slightly negatively charged pyrite surface that can attract cations such as Fe2+. Because of surface electroneutrality and pH considerations, however, the pyrite surface Fe2+ coordinates OH. We proposed that this surface Fe2+ OH when in the presence of CO2 is converted to -FeCO3 or -FeHCO3, depending on pH. In this study, using Fourier transform infrared spectroscopy (FT-IR) we demonstrated that such complexes form on the surface of pyrite and continue to persist even after a significant fraction of the surface Fe2+ was oxidized to Fe3+. FT-IR spectra also showed the presence of two carbonyl absorption bands (1682 and 1653 cm-1) on the surface of pyrite upon exposure to CO2 which suggested that pyrite surface carbonate complexes existed in two different surface chemical environments, pointing out two potential mechanisms of pyrite surfaceCO2 interactions. One potential mechanism involved formation of a pyrite surface-Fe(II)HCO3 complex, whereas a second potential mechanism involved formation of a pyrite surface-carboxylic acid group complex [-Fe(II)SSCOOFe(II)]. We hypothesized that these pyrite surface-CO2 complexes could promote abiotic oxidation of pyrite by accelerating the abiotic oxidation of Fe2+. Iron (III) would oxidize the disulfide (-S2) by accepting its electrons. Using a miscible displacement technique, oxidation of FeS2 with H2O2 was carried out in the absence or presence of 10 or 100 mM NaHCO3. The data show that 100 mM NaHCO3 significantly increased the oxidation rate of FeS2. Furthermore, the data show that FeS2 oxidation kinetics were more dependent on HCO3- but were less dependent on H2O2 for the range of HCO3- and H2O2 concentrations tested.
Introduction The need to prevent the development of acid mine drainage (AMD) by oxidation of pyrite has triggered numerous investigations into the mechanisms of its oxidation (1-9). Singer and Stumm (1) reported that Fe3+ can oxidize pyrite at a much higher rate than O2. The role of Fe3+ in the oxidation of pyrite is demonstrated below * Corresponding author e-mail:
[email protected]. † University of Kentucky. ‡ Universitas Bengkulu. 2084
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FeS2 + 14 Fe3+ + 8H2O
15Fe2++ 2SO42– + 16H+ (1)
O2 iron-oxidizing bacteria
At low pH, Thiobacillus ferrooxidans (an acidophilic, chemoautotrophic, iron-oxidizing bacterium) catalyzes and accelerates the oxidation of Fe2+ by a factor larger than 106 (1). Thus, the processes of reduction of Fe3+ to Fe2+ by pyrite and oxidation of Fe2+ to Fe3+ by atmospheric O2, catalyzed by T. ferrooxidans, represent an effective continuous pyrite oxidation cycle. It is believed that pyrite oxidation can be controlled by adding alkaline material to pyrite, a widely used field practice. This is due to the fact that the activity of T. ferrooxidans diminishes and Fe3+ becomes insoluble at high pH (10). However, recent findings (4) revealed that nonmicrobial pyrite oxidation rates increased as pH increased. At circumneutral pH, nonmicrobial pyrite oxidation appeared to be carried out by a surface-catalyzed mechanism involving Fe2+, O2, and Fe3+ (6, 11). According to Moses and Herman (4) and Moses et al. (5), Fe3+ is an effective and direct pyrite oxidant at low pH as well as at circumneutral pH, and the role played by dissolved O2 is to sustain the reaction by regenerating Fe3+. It is wellknown that the abiotic rate of Fe2+ oxidation increases as pH increases (1, 12). The role of OH in abiotically oxidizing Fe2+ has been postulated by Luther et al. (9) to be due to the potential increase in frontier molecular orbital electron density of Fe2+ upon binding to oxygen by coordinating OH. Luther et al. (9) demonstrated that an increase in electron density also increases the potential of Fe2+ to oxidize rapidly to Fe3+ when in the form of a complex with a ligand containing oxygen as the ligating atom. Nicholson et al. (2, 3) reported that pyrite oxidation kinetics in a bicarbonate (HCO3-) buffered system initially increased, but hypotheses regarding this behavior were not given. Hood (11) also reported that the rate of pyrite oxidation by atmospheric O2 at room temperature under abiotic conditions increased when HCO3- concentration increased. The author postulated that the cause for the increase in pyrite oxidation was formation of a pyrite surface-Fe(II)CO3 complex which facilitated electron transfer from Fe(II) to O2. However, no evidence of pyrite-Fe(II)CO3 complexes was presented by Hood (11). The objectives of this study were to provide molecular evidence of pyrite surface-Fe(II)-CO3 complexes and to evaluate the influence of solution HCO3on pyrite oxidation.
Theoretical Considerations To explain the potential enhancement of pyrite oxidation by HCO3- or CO2 gas, we are postulating the following models. Formation of a ferrous-bicarbonate complex on the surface of pyrite occurs as follows
Fe-SA-SB:Fe(OH)2 + CO2 f
Fe-SA-SB:FeHCO3+ + OH- (2)
or
Fe-SA-SB:Fe(OH)2 + HCO3- f
Fe-SA-SB:FeHCO3+ + 2OH- (3)
S0013-936X(97)00829-8 CCC: $15.00
1998 American Chemical Society Published on Web 05/21/1998
Formation of the ferrous-bicarbonate complex on the surface of pyrite would increase the basicity of Fe2+ leading to its rapid oxidation (9, 12)
Fe-SA-SB:FeHCO3+ + 1/4O2 + H+ f Fe-SA-SB:Fe(OH)2+ + CO2 + 1/2H2O (4) Reaction 4 spontaneously forms a persulfido bridge after Fe(II) oxidation followed by decarboxylation and pyrite oxidation initiates as shown below Fe–SA–SB:Fe3+
• Fe–SA–SB :Fe2+
persulfido bridge
(5)
free radical
This leads to a continuous loss of electrons until thiosulfate is produced as shown below – –
– O
Fe–SA–SB:(Fe3+)5 + 3H2O
– –
Fe–SA SB O + 6H+ + 5Fe2+ (6) O – thiosulfate
In the above reactions, -S2 undergoes five continuous redox reactions to produce thiosulfate. Thiosulfate is rapidly oxidized by H2O2 or Fe3+ to SO4, yielding eight more H+ ions (details are given in ref 13 and references therein). It follows that pyrite oxidation by an oxidizer, such as H2O2, is an autocatalytic process since one of the oxidation products, Fe3+, can also oxidize pyrite (1, 4, 5)
FeS2 + 7H2O2 f Fe2+ + 2SO42- + 2H+ + 6H2O (7) Fe2+ + H+ + 0.5H2O2 f Fe3+ + H2O
(8)
and
FeS2 + 14Fe3+ + 8H2O f 15Fe2+ + 2SO42- + 16H+ (9) The rate law describing production of SO4 through oxidation of FeS2 by H2O2 (eq 7) and Fe3+ (eq 9) can be written as (13)
d[SO4]/dt ) (k1[H2O2] + k2[Fe3+])S
(10)
where SO4 is directly related to the quantity of decomposed pyrite at any time t; [H2O2] and [Fe3+] are the concentrations of H2O2 and Fe3+; k1 and k2 are the rate constants of H2O2 and Fe3+, respectively; and S is the surface area of pyrite available to react. Assuming that, during oxidation of a small portion of the available pyrite, the newly exposed pyrite surface (S) remains proportional to the number of moles of unreacted FeS2 in the system, then (14)
S ) K[FeS2]
(11)
where K is a constant. Considering also that during pyrite oxidation, as t approaches zero S, [H2O2] and Fe3+ (eq 10) remain relatively constant, the oxidation process can be expressed as zero-order reaction
SO4produced ) k′t
(12)
where k′ is an empirical rate constant related to the concentrations of H2O2 and Fe3+; and the corresponding rate constants, k1 and k2 (eq 10), overall oxidizing solution composition; and surface area of available pyrite. Since HCO3- can enhance oxidation of Fe2+, increasing HCO3concentration should increase the rate of pyrite oxidation by rapidly regenerating Fe3+ (1, 4).
Ionic strength is a second potential reason for the HCO3to affect the oxidation rate of pyrite. Its influence on the rate of reaction between any two chemical species was described by Benson (15) as follows:
log(kexp/kid) ) 1.0ZAZB(I)1/2
(13)
where kexp ≈ k′ (eq 12) denotes the experimental rate constant under a given ionic strength; kid denotes the rate constant at infinite dilution; ZA and ZB denote the charges of ions A and B, respectively; and I denotes ionic strength. A plot of log(kexp/kid) versus (I)1/2 would produce a straight line with slope ZAZB. Benson (15) pointed out that when one of the Zi values in eq 13 is 0, the ionic strength would not influence the reaction rate. On the other hand, when one of the Zi values is negative and the other is positive, the influence of ionic strength on the rate of reaction would be negative. When both Zi values are positive, the influence of ionic strength on reaction rate would be positive. In the case of Fe2+ being oxidized to Fe3+ by H2O2, the latter’s charge is 0, and for this reason the influence of ionic strength on Fe2+ oxidation would be negligible. Assuming that Fe2+ oxidation is the rate-controlling step in pyrite oxidation (2) under our experimental conditions, one could then assume that ionic strength would have negligible effect on pyrite oxidation. However, any increase in pyrite oxidation due to increasing HCO3- concentration would imply that -Fe(II)HCO3- complexes were responsible for the rapid regeneration of Fe3+.
Materials and Methods The pyrite sample used in this study was obtained from the University of Kentucky Applied Energy Research Laboratory, Lexington KY. It was separated by gravity, using a water column, from coal material obtained from seam number 9 in western Kentucky. The sample was pulverized and passed through a 37 µm sieve, then washed with 4 M hydrofluoric acid to remove silicate and Fe oxides, rinsed repeatedly with N2-purged distilled water and acetone, and stored in a desiccator under vacuum (16, 17). This pyrite sample was then characterized prior to surface spectroscopic studies. Characterization of sample crystallinity was carried out by X-ray diffraction analysis (XRD) using an 1840 Philips Co-Ka diffractometer, while surface morphologies of individual pyrite grains were examined using a scanning electron microscope (SEM) [Hitachi S-800]. Specific surface analysis was carried by multipoint BET method using a Quantachrome Autosorb 6 instrument. Pyrite oxidation experiments were carried out by placing approximately 50 mg of clean pyrite into small glass vials. The unstoppered vials were then put into a chamber, and various gas treatments were introduced. These treatments were the following: (a) a control sample with a constant exposure of pyrite to an N2 gas atmosphere, (b) a 14-day exposure to atmospheric air at 100% relative humidity [After exposure to 100% relative humidity, a portion of the sample was rinsed with 10 mL of 4 M HCl through the use of a buchner filter-funnel under a N2 atmosphere], (c) a 14-day exposure to atmospheric air at 100% relative humidity followed by heating at 50, 100, and 249 °C [heating was carried out by using 25-50 mg of pyrite in a furnace DuPont S/N-X, FTIR-TGA interface] for 4 h under a N2 gas atmosphere, followed by a 2-h evacuation in a desiccator, and (d) a 14-day exposure to atmospheric air at 100% relative humidity followed by extraction with 1 M KCl to determine total iron [Fe3+ plus Fe2+] and Fe2+ on the surface of pyrite (18) [all samples representing all treatments were evacuated for 2 h, using a desiccator, prior to scanning]. For iron analysis, 10 mg of the pyrite samples representing the control and the 14-day exposure to 100% relative humidity was placed into 50-mL centrifuge tubes to which 10 mL of VOL. 32, NO. 14, 1998 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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1.0 M KCl was added. After equilibration for 2 h using endto-end shaking, the samples were centrifuged and the supernatants were collected for analysis. Iron was quantified colorimetrically following procedures outlined by American Water Works Association (19). For total Fe determination, 150 µL aliquots of samples and standards were pipetted into a disposable microplate and treated with 60 µL of hydroxylamine/phenanthroline solution and 30 µL of an ammonium acetate pH 5 buffer solution. The samples were shaken on a vortex mixer for 10 min before they were scanned at 540 nm (Bio-Tek Instruments microplate autoreader EL311). For Fe2+ determination, 150 µL aliquots of samples and standards were pipetted into a disposable microplate to which 60 and 30 µL of the phenanthroline and ammonium acetate buffer solutions, respectively, were added. The samples were shaken on a vortex mixer for 10 min before scanning them at 450 nm. Infrared (IR) analyses were carried out employing a Nicolet 5XSC Fourier transform infrared spectrophotometer (FT-IR). The procedure for generating FT-IR spectra involved mixing 100 mg of KBr with 2.5 mg of pyrite and placing the mixture into the FT-IR sample holder (20). All FT-IR analyses were performed under a nitrogen gas atmosphere. For background correction KBr alone was used after the instrument was purged with N2 gas and 50 scans were collected. This number of scans was determined by trial and error using as criteria reproducibility of noice-clean spectra. A similar approach was used for all pyrite samples scanned. In addition to pyrite spectra, FT-IR spectra of reagent grade NaHCO3, Na2CO3, KHCO3, and K2CO3 were also obtained as references for carbonate and bicarbonate species. All treatments were duplicated, and a minimum of two subsamples from each duplicate was used to generate FT-IR spectra. The kinetics of pyrite oxidation were studied employing a bed-reactor, consisting of a nuclepore swin-lok filter holder (25-mm inside diameter) similar to that described by Jardine and Sparks (21). The segregated pyrite from the coal, as previously described, was ground and passed through 47and 75-µm sieves. The fraction 47 µm and 75 µm was used for the present study. Prior to using the pyrite for leaching studies, the samples were washed with 4 M hydrofluoric acid and repeatedly rinsed with N2-purged distilled water and acetone. For each test run, approximately 50 mg of the separated pyrite was placed in a nuclepore filter holder between two Whatman no. 42 ashless filters. The column was first rinsed with 5 mL of 1 N HCl and then 5 mL of 5 mM NaCl to remove any surface oxidation products prior to initiating leaching. The leaching solutions were passed through the pyrite nuclepore filter column at a constant flow rate of 0.50 mL min-1 at room temperature (22 °C). Aliquots of the effluent solution were collected every 2 min for the first 16 min and then at 10 min intervals for a total of 1 h using an Eldex Universal fraction collector. Five duplicate tests were run: (1) 10 mM NaHCO3 with 20 mM H2O2; (2) 100 mM NaHCO3 with 20 mM H2O2; (3) 10 mM NaHCO3 with 2 mM H2O2; (4) 100 mM NaHCO3 with 2 mM H2O2; and (5) 10 mM NaHCO3, 20 mM H2O2, and 100 mM NaNO3. The ranges of NaHCO3 and NaNO3 concentration were chosen on the basis of the results reported by Millero and Izaquirre (12). Their results showed that abiotic solution Fe2+ oxidation rates were not affected by 10 mM solutions of NaHCO3 or NaNO3 relative to the control (absence of NaHCO3 or NaNO3). Furthermore, solution Fe2+ oxidation rates remained unaffected when in the presence of 100 mM NaNO3 but increased approximately 10-fold when in the presence of 100 mM NaHCO3. Activation energies (Ea) were determined by assessing pyrite oxidation kinetics at temperatures of 10, 25, 30, and 40 °C using 20 mM H2O2 with 10 or 100 mM NaHCO3. Constant temperature during each run was maintained by 2086
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a 2095 Bath & Circulator using a pyrite sample holder which consisted of a Pharmacia XK 26-mm i.d. column equipped with thermostat jacket. The data were plotted according to the Arrhenius equation (eq 12) [log k′ vs 1/T, where k′ denotes experimental rate constant (eq 12) and T denotes temperature in degrees Kelvin]. Sulfate concentration in all leachates was measured using turbidimetry with BaCl2 (23). Regression analysis on SO4 production versus time was carried out using the duplicated data with slope denoting the conditional zero-order rate constant (eq 12).
Results and Discussion X-ray diffraction analysis of the iron sulfide sample substantiated that the sample was pyrite (24). The micrograph in Figure 1A shows the surface morphology of the pyrite sample, demonstrating a spongy-like appearance and a large specific surface. The micrograph in Figure 1B reveals the absolute size and size distribution of the pyrite particles selected for the wet chemistry oxidation study using H2O2 and HCO3-. It is important to note that the majority of the particles appear to be fairly uniform in size, although a small number of extremely small particles are also present. Specific surface analysis for the less than 75 µm particles from a fourpoint N2 BET isotherm (25) was found to be 9.75 m2 g-1. Commonly, coarse-grained, massive pyrite specimens in the range of 38 to 250 µm particles have surface areas that range from 0.071 to 0.047 m2 g-1 (5, 26). The large specific surface of the pyrite in the present study may account for the high acid mine drainage production potential for the geologic stratum, Kentucky coal seam 9. Figure 2 shows the carbonate and bicarbonate spectra of KHCO3/K2CO3 and NaHCO3/Na2CO3, respectively. The data in Figure 2I reveal a major vibrational band for Na2CO3 occurring around 1406 cm-1 and one for K2CO3 occurring at around 1450 cm-1 (Figure 2II). These bands represented the ν3 of the CO3 vibrations (27). The carbonate anion (CO32-) is highly symmetrical, and ν3 is infrared active (27). When the symmetry of carbonate is perturbed due to coordination with cations, the degenerate vibrations (ν3) split into a number of distinct vibrations which depend on metal-carbonate coordination. For example, in the case of reagent grade K2CO3 (Figure 2II), the CO3 band split into three distinct bands signifying that CO3 coordinated in a bidentate fashion (27). In the case of NaHCO3 (Figure 2I) and KHCO3 (Figure 2II), the two major vibrational bands were at 1601 and 1631 cm-1, respectively (28). There were significant differences in the vibrational spectra of reagent grade Na2CO3/NaHCO3 and K2CO3/KHCO3 spectra. These differences were most likely due to differences in the physicochemical properties between the two different carbonate-associated cations, Na+ and K+. One major difference between these two cations is their ionic potential, defined as the ratio between charge and radius (c/r2) (29). The ionic potential for Na+ is 1.11 and for K+ is 0.56. Difference in ionic potentials produces differences in the electron orbital perturbation of CO3 associated with these two metals; for this reason the vibrational spectra of Na2CO3/ NaHCO3 and K2CO3/KHCO3 differ (Figure 2). The FT-IR spectra for pyrite before and after oxidation by exposure to humidified air for 14 days at room temperature were presented in Figure 3. These spectra show that before oxidation no absortion bands were apparent. After oxidation, however, pyrite showed absorption bands at 3539 and 3406 cm-1. These bands represented OH most likely associated with both Fe on the surface of pyrite and pyrite surface adsorbed water. The data in Table 1 show the presence of iron on the surface of the pyrite sample representing the control. Furthermore, iron concentration on the pyrite surface increased approximately 4-fold during oxidation,
FIGURE 1. Scanning electron microscope photograph of (A) surface of pyrite and (B) a number of pyrite particles less than 75 µm size before any oxidation treatment. although the ratio between Fe2+ and total Fe before and after oxidation did not differ dramatically (0.38 vs 0.31, respectively) The difference between total Fe and Fe2+ was assumed to represent Fe3+. After exposing the pyrite to humidified air for 14 days, the sample was heated to various temperatures in a N2 atmosphere and then cooled to room temperature in a desiccator under vacuum. The spectra of these pyrite samples under the various heating treatments are presented in Figure 4. Note, after heating the pyrite to 50 °C, the intensity of the 3539 cm-1 band (Figure 4I) decreased, suggesting that the OH, coordinated to Fe2+ or Fe3+ on the surface of pyrite, was consumed perhaps by increasing surface acidity. Pyrite surface pH may decrease by decreasing water content (30) on the pyrite surface and/or by increasing electron transfer between the disulfide and Fe3+ at increasing temperatures. As temperature increased, the intensity of the H2O band (3406 cm-1) decreased as expected. Finally, when the pyrite was rinsed with 4 M HCl, neither OH band was apparent (spectra E in Figure 4I) signifying that most surface iron was removed from the surface of pyrite (also OH loss via the acid effect outlined above). Support for the above observations and conclusions is also given by the spectra behavior around the 1000 cm-1 frequencies (Figure 4II). The emergence of 1227, 1090, and 1009 cm-1 absorption bands indicates the presence of SO42associated with Fe3+ on the surface of pyrite (31). The sulfate ion (SO42-) is highly symmetrical, and ν3 (around 1100 cm-1) is infrared active (27). When the symmetry of sulfate is perturbed due to coordination with cations, the degenerate vibrations (ν3) split into a number of distinct vibrations, depending on strength of the complex, and ν1 is activated (27). The apparent splitting of ν3 (1227, 1090 cm-1) and the presence of ν1 (1009 cm-1) on the spectrum of oxidized pyrite (Figure 4II) suggested that the sulfate may be bonded to Fe3+ in a monodentate fashion (27). When the oxidized pyrite sample was heated to 249 °C, the 1227 cm-1 band disappeared, which suggested, as in the case of the OH bands (Figure 4I), that pyrite surface pH decreased through water
loss (30), weakening Fe3+-SO42+ interactions by forming HSO4-, or through surface acidification by increasing electron transport between the disulfide and Fe3+ at increasing temperature; SO4 was most likely coordinated to the newly produced Fe2+ on the pyrite surface in a weak electrostatic fashion. The spectra of oxidized pyrite (Figure 4III) also exhibit carbonate absorption bands at 1429 cm-1, which are assigned to the ν3 of CO3 vibrations. According to Nakamoto (27), the extent of splitting (number of distinguishable bands) of the 1429 cm-1 band, representing the free CO32- ion, could be directly related to the bond strength between CO32- and a metal cation or surface. Nonsplitting of the CO3 band in Figure 4III indicates that pyrite surface-CO3 complexes were most likely weak electrostatic complexes. One would assume that the strong acid environment on the surface of pyrite, created by the oxidation of the disulfide, is expected to prevent formation of surface carbonate complexes. However, persistence of the carbonate absorption band (1429 cm-1), as shown in Figure 4III, demonstrated that some of the acid produced on the pyrite surface was removed through reaction with pyrite as follows:
FeS2 + H2SO4 ) FeSO4 + H2S + S°
(14)
Equation 14 shows that H2SO4 could be removed from the pyrite surface through formation of H2S, a weak acid with pKa’s 7 and 12.9 (32). Therefore, after pyrite exposure to humidified air and heating, its surface was most likely covered with ferrous sulfate and S° and most active anodic sites (SB) were saturated with Fe2+ (4, 32). The spectra in Figure 4III also show a number of vibrational bands in the 1600 cm-1 range. The 1621 cm-1 band was most likely due to H2O deformation. The absorption bands at 1653 and 1670 cm-1 most likely represent carboxylate species in two different chemical environments (27, 28). When the oxidized pyrite was heated to 50 °C under a N2 gas atmosphere, the absorption band at 1653 cm-1 was eliminated perhaps due to pyrite surface acidification inducVOL. 32, NO. 14, 1998 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 1. Total Fe and Fe2+ on the Surface of Pyrite before and after Oxidation sample
total Fe (mg g-1)
Fe2+ (mg g-1)
before oxidation after oxidationa
4.72 22.48
1.82 6.99
a
14-day exposure to humidified air.
1670 cm-1 bands (spectra E, Figure 4III), suggesting that the pyrite surface carbonate was removed due to surface acidification, inducing conversion of pyrite surface HCO3to CO2 gas. The absence of the 1621 cm-1 band from spectra E (Figure 4III) could be due to production of elemental sulfur on the surface of pyrite (23), a hydrophobic substance. It follows from above that carbon dioxide may react with the surface of pyrite via two mechanisms. One mechanism may involve coordination of HCO3- by Fe2+ on the surface of pyrite as shown below: –
–O –Fe(II)–SA–SB:Fe(II)OH + C O
O –Fe(II):SA–SB:Fe(II)OC
–
–
(15) OH
This pyrite surface-bicarbonate complex is speculated to be sensitive to acid attack, releasing CO2 at about pH 4.2 [approximately 2 pH units lower than the pKa of H2CO3 (32)]. The second pyrite surface-CO2 complex may involve formation of a sulfur-carbon bond on the pyrite surface as shown below: –
–O –Fe(II)–SA–SB: + C O
–
FIGURE 2. FT-IR diffuse reflectance spectra of (IA) Na2CO3; (IB) NaHCO3; (IIA) K2CO3; and (IIB) KHCO3.
FIGURE 3. FT-IR diffuse reflectance spectra of pyrite: (A) before oxidation (after washing with 4 M hydrofluoric solution) and (B) after oxidation by exposure to air at 100% relative humidity for 14 days. ing conversion of pyrite surface HCO3- to CO2 gas (decarboxylation), while the 1670 cm-1 band remained intact. When the temperature was raised to 100 °C, the intensity of the 1621 cm-1 band decreased due most likely to water removal from the surface of pyrite; the 1670 cm-1 band was completely eliminated. This suggested that the 1670 cm-1 band could be representative of a second carbonyl which was removed at a higher temperature than the carbonyl represented by the 1653 cm-1 band. After washing the oxidized pyrite sample with HCl, its spectrum did not exhibit the 1429, 1653, and 2088
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O (16)
–Fe(II):SA–SB–C
–
O
This complex is speculated to be more resistant to acid attack, perhaps because of the covalent interaction between the carbon and sulfur atoms; its removal from the pyrite surface takes effect at higher temperature (Figure 4III). However, the exact reasons for this difference in temperature resistance between the two carbonyls are not clearly apparent from the present data. The data in Figure 5 and Table 2 show that pyrite oxidation is affected by the presence and concentration of HCO3-. At a H2O2 concentration of 20 mM, increasing the HCO3concentration from 10 to 100 mM increased oxidation rate by about 174% (calculated from the difference between slopes) (Figure 5A). Similarly, when H2O2 concentration was set at 2 mM, the oxidation rate was about 125% higher at the 100 mM HCO3- concentration than at the 10 mM HCO3concentration (Figure 5B). The concentration of H2O2 in this experiment also influenced pyrite oxidation. Its influence, however, was found to be smaller than that of HCO3within the range of H2O2 and HCO3- concentrations tested. At a HCO3- concentration of 10 mM, increasing the H2O2 concentration from 2 to 20 mM increased the oxidation rate by about 75%. Similarly, when HCO3- concentration was set at 100 mM, the oxidation rate was about 113% higher at the 20 mM H2O2 concentration than at the 2 mM H2O2 concentration. The role of H2O2 on pyrite oxidation was most prominent at the higher HCO3- concentration. The above results imply that at the higher HCO3- concentration, the oxidation rate was controlled mainly by the more rapid regeneration of Fe3+ due to HCO3-. It appears that Fe3+ is a more effective pyrite oxidant than H2O2 because of its small ionic radius (steric effect) and large potential to form innersphere complexes (transition-state intermediate) with the pyrite surface. On the other hand, H2O2 is a molecule with a relatively large radius, thus limiting its potential to form
FIGURE 5. Influence of sodium bicarbonate concentration on the oxidation of pyrite using (A) 2 mmol L-1 H2O2 and (B) 20 mmol L-1 H2O2. The two slopes within each graph are significantly different at the 95% confidence level according to the t-test.
TABLE 2. Summary of Reaction Rates of the Various Systems Tested
FIGURE 4. FT-IR diffuse reflectance spectra of pyrite after the following treatments: (A) oxidation with air for 14 days; (B) heating at 50 °C; (C) heating at 100 °C; (D) heating at 249 °C; and (E) washing with 4 M HCl (Roman numerals I, II, and III signify different wavenumber ranges). pyrite surface complexes with all reactive sites potentially available to Fe3+, and with a low potential to form pyrite
oxidizing systems
slope
r2
20 mmol L-1 H2O2 plus 10 mmol L-1 HCO320 mmol L-1 H2O2 plus 100 mmol L-1 HCO32 mmol L-1 H2O2 plus 10 mmol L-1 HCO32 mmol L-1 H2O2 plus 100 mmol L-1 HCO320 mmol L-1 H2O2 plus 100 mmol L-1 NaNO3
0.042 0.115 0.024 0.054 0.034
0.999 0.996 0.999 0.996 0.996
surface inner-sphere complexes (16). The data in Figure 6 show that the increase in pyrite oxidation when HCO3- concentration was increased from 10 to 100 mM could not be attributed to ionic strength. This can be deduced from the fact that the addition of NaNO3 (100 mM) in the 10 mM HCO3- solution plus 20 mM H2O2 did not in any way affect the oxidation of pyrite. Activation energy (Ea) data (Table 3) were also consistent with what one might expect if indeed -Fe2+HCO3 complexes controlled the rate of pyrite oxidation. In other words, increasing HCO3concentration in the oxidizing solution should decrease Ea. The data in Table 3 show that upon increasing HCO3concentration from 10 mM to 100 mM, Ea decreased from 28.96 to 25.92 kJ mol-1. The difference in Ea between the two systems was small but meaningful. Note, however, the Ea VOL. 32, NO. 14, 1998 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 3. Apparent Activation Energies of Pyrite Oxidation by 20 mmol L-1 H2O2 as Affected by NaHCO3
a
NaHCO3 treatment (mmol L-1)
Arrhenius regression equations (log k′ ) log A Ea/RT)
10 100
log k′ ) 3.692 - (1511.5)(1/T) log k′ ) 3.306 - (1352.4)(1/T)
n
apparent activation energies (kJ mol-1)
log k′ vs 1/T (r 2 )
4 4
28.96 ( 0.045a 25.92 ( 0.023a
0.896 0.977
Confidence interval at 95%.
the abiotic pyrite oxidation rate aided by HCO3- is not expected to be as rapid as the biotic oxidation rate (1).
Acknowledgments This research was supported, in part, by funds provided by the United States Department of Energy and the Pittsburgh Energy Technology Center under Grant DE-FG22-95PC5226. This support is gratefully acknowledged. We also wish to thank Ms. Joni Norris and Mr. Martin Vandiviere for their help in generating the data. This research was published with the approval of the Director of Kentucky Agricultural Experimental Station, Lexington.
Literature Cited
FIGURE 6. Influence of 100 mmol L-1 NaNO3 on the oxidation of pyrite by 20 mmol L-1 H2O2 and 10 mmol L-1 sodium bicarbonate. The two slopes are not significantly different at the 95% confidence level according to the t-test. values reported represent apparent activation energies, which include diffusion-controlled reactions on the surface of pyrite as well as adsorption and precipitation. The overall data suggest that -Fe2+HCO3 complexes on the surface of pyrite increased its oxidation rate.
Conclusions FT-IR spectroscopic evidence is presented which demonstrates that pyrite exposed to humidified CO2 plus O2 formed pyrite surface-CO2 complexes. Two potential mechanisms were proposed to account for these complexes. One mechanism involved formation of a weak pyrite surface-Fe(II)HCO3 complex, whereas a second mechanism involved formation of a pyrite surface-carboxylic acid group complex [-Fe(II)SSCOOFe2+]. It is hypothesized that both complexes could promote abiotic pyrite oxidation by accelerating regeneration of Fe3+, an effective pyrite oxidant. Our experimental data clearly showed that abiotic pyrite oxidation increased in the presence of solution HCO3-. The approach currently used to prevent pyrite oxidation in the field is mainly based on eliminating Fe3+ from pore waters (10, 13). This approach includes using limestone to precipitate Fe3+ as iron hydroxide/oxyhydroxide (13) and raising pH to diminish activity of T. ferrooxidans (1). Recent evidence (4), however, showed that Fe3+, associated with the surface of pyrite, is an effective pyrite oxidant at low pH as well as at near-neutral pH. The findings in this study suggest that in pyritic geologic waste the presence of HCO3-, added in the form of either limestone or Na2CO3, may accelerate the abiotic oxidation rate of pyrite. Therefore, application of alkaline materials to pyritic waste may neutralize AMD but may not inhibit sulfate production. Observations to this effect have been reported in the literature (13). However, 2090
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(1) Singer, P. C.; Stumm, W. Science 1970, 167, 1121. (2) Nicholson, R. V.; Gillham, R. W.; Reardon, E. J. Geochim. Cosmochim. Acta 1988, 52, 1077. (3) Nicholson, R. V.; Gillham, R. W.; Reardon, E. J. Geochim. Cosmochim. Acta 1990, 54, 395. (4) Moses, C. O.; Herman, J. S. Geochim. Cosmochim. Acta 1991, 55, 471. (5) Moses, C. O.; Nordstrom, D. K.; Herman, J. S.; Mills, A. L. Geochim. Cosmochim. Acta 1987, 51, 1561. (6) Brown, A. D.; J. J. Jurinak. J. Environ. Qual. 1989, 18, 545. (7) Luther, G. W., III Geochim. Cosmochim. Acta 1987, 51, 3193. (8) Luther, G. W., III In Aquatic Chemical Kinetics; Stumm, W., Ed.; John Wiley & Sons, Inc.: New York, 1990; p 173. (9) Luther, G. W., III; Kostka, J. E.; Church, T. M.; Sulzberger, B.; Stumm, W. Mar. Chem. 1992, 40, 81. (10) Jaynes, D. B.; Rogowski, A. S.; Pionke, H. B. Water Resour. Res. 1984, 20, 233. (11) Hood, T. A. Ph.D. dissertation, University of Miami, 1991. (12) Millero, F. J.; Izaguirre, M. J. Solution Chem. 1989, 18, 585. (13) Evangelou, V. P. Pyrite oxidation and its control; CRC Press: Boca Raton, FL, 1995. (14) Turner, R. C. Can. J. Soil Sci. 1960, 40, 219. (15) Benson, S. W. The foundations of chemical kinetics; Robert E. Krieger Publishing Co.: Malabar, 1982. (16) Zhang, Y. L.; Evangelou, V. P. Soil Sci. 1996, 161, 852. (17) Huang, X.; Evangelou, V. P. In The Environmental Geochemistry of Sulfide Oxidation; Alpers, C. N., Blowes, D. W., Eds.; American Chemical Society: Washington DC, 1994. (18) Thomas, W. G. In Methods of Soil Analysis, Part 2, Chemical and Microbiological Properties, 2nd ed.; Page, A. L., Ed.; Soil Science Society of America: Madison WI, 1982; Agronomy Series. (19) Clesceri, L. S., Greenberg, A. E., Trussel, R. R., Eds. Standard Methods for the Examination of Water and Wastewater, 17th ed.; APHA-AWWA-WPCF: Washington, DC. (20) Griffiths, P. R.; Fuller, P. M. Mid-infrared spectroscopy of powdered samples; Heyden: Chichester, U.K., 1981; Vol. 9, p 63. (21) Jardine, P. M.; Sparks, D. L. Soil Sci. Soc. Am. J. 1984, 48, 39. (22) Daniels, F.; Alberty, R. A. Physical Chemistry, 4th ed.; John Wiley & Sons: New York, 1975. (23) Massoumi, A.; Cornfield A. H. Analyst 1963, 88, 321. (24) Wei, D.; Osseo-Asare, K. Colloids Surf. 1997, 121, 27. (25) Brunauer, S.; Emmett, P. H.; Teller, E. J. Am. Chem. Soc. 1938, 60, 309. (26) Rimstidt, J. D.; Newcomb, W. D. Geochim. Cosmochim. Acta 1993, 57, 1919. (27) Nakamoto, K. Infrared and Raman spectra of inorganic and coordination compounds; John Wiley & Sons: New York, 1986. (28) Harter, R. D.; Ahlrich, J. L. Soil Sci. Soc. Am. Proc. 1967, 31, 30. (29) Keay, J.; Wild, A. Soil Sci. 1961, 92, 54. (30) Mortland, M. M.; Raman, K. V. Clays Clay Miner. 1968, 16, 393. (31) de Donato, P.; Mustin, C.; Berthelin, J.; Marion, P. C. R. Acad. Sci., Ser. II 1991, 321, p 241.
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Received for review September 16, 1997. Revised manuscript received March 24, 1998. Accepted April 21, 1998.
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