A vibrational spectroscopic 18O tracer study of pyrite oxidation

Gmchm~ca c, Cosmi~himica Copyright 0 1991 Pergamon

0016.7037/91/$3.00 + .OO

/Ma Vol. 55, pp. 1609-1614 Press pk. Printed in U.S.A.

A vibrational spectroscopic lsO tracer study of pyrite oxidation BRIAN J. REEDY,’ JAMESK. BEATTIE,’and RICHARDT. LOWSON’ ‘School of Chemistry, University of Sydney, NSW 2006,Australia ‘Environmental Science Program, Australian Nuclear Science and Technology Organisation, Lucas Heights Research Laboratories, Sutherland NSW 2232, Australia (Received July 11, 1990; accepted in revisedform March 4, 199 1)

Abstract-Pyrite was oxidised under 1802gas in H2160 solutions, with and without added ferric ion, and the sulfate produced was analysed by vibrational spectroscopy to determine the relative amounts of sulfate isotopomers (S’60,L80::,) formed. At 70°C and pH 1, with no added Fe3+, the majority of the sulfate formed was that which derived all four oxygen atoms from water (i.e., S”O$-), but significant amounts of two other isotopomers, S’603r802- and S’602’*O$-, which derive one or two oxygen atoms from molecular oxygen were observed. When Fe3” was added at the start under identical conditions, no S1602’80$- was observed. The major isotopomer formed was still S’60z-, with S’“03’802- present as a minor product. Experiments which were performed at initial pH 7 yielded similar results, as did others performed at 2O”C, although the amounts of the minor isotopomers formed vary with temperature. All of the results were confirmed by performing identical experiments with the source of the oxygen isotopes reversed, that is, by oxidising pyrite under air in H 2’gO solutions and obtaining the same products in isotopic reverse. INTRODUCTION

Several oxygen isotope studies have been performed on sulfate from pyrite oxidation with the aim of determining the contributions (to the sulfate) of oxygen from water and dissolved molecular oxygen (SCHWARTZ and CORTECCI, 1974; BAILEY and PETERS, 1976; TAYLOR et al., 1984a,b; TORAN, 1987; VAN EVERDINGEN and KJ~OUSE, 1988). SCHWARCZand CORTECCI (1974) oxidised pyrite at 25°C under air as a slurry in water of very low ‘*O-enrichment and concluded that approximately half of the oxygen in the sulfate produced came from water. BAILEYand PETERS ( 1976) performed ‘*O-tracer experiments on pyrite oxidation at high temperatures (85-l 1O’C) and pressures (up to 66.4 atm) which showed that water is the source of 73-100% of sulfate oxygen under the conditions studied. TAYLORet al. (1984a,b) studied the oxygen isotope composition of sulfate from pyrite oxidation both in acid mine waters and in the laboratory under a variety of conditions (with and without added Fe3+; with and without Thiobacillus ferrooxiduns, an iron and sulfur oxidising bacterium; aerobic and anaerobic), using natural isotope abundance oxygen gas and water. In their sterile (totally inorganic) laboratory experiments, they found that submersed, aerobic oxidation of pyrite at pH 2 produced sulfate which derived most of its oxygen from water. They concluded that this indicated the dominance of reaction (2) over reaction (1) (as an overall process) under those conditions. However, in alternating wet/dry experiments, where Fe3+ had less contact with the pyrite surface, this figure was reduced somewhat. In contrast, where bacteria were allowed to mediate in the process, most of the sulfate oxygen came from molecular oxygen instead. VAN EVERDINGEN and KROUSE(1988) also performed similar submerged/aerobic pyrite oxidation experiments and concluded that between 29 and 100% of sulfate oxygen was derived from water in experiments they surveyed (including their own). TORAN ( 1987) found that between 65 and 85% of the oxygen in sulfate from a carbonate aquifer came from water, These results give some information about

THE MECHANISM OF

pyrite oxidation, a process of environmental and potential economic importance, is still poorly understood, despite extensive study. Two general, non-bacterial reactions for the oxidation of pyrite in acidic solutions are postulated. One involves dissolved molecular oxygen as the oxidant: Fe& + 7/202 i- Hz0 -t Fe*’ f ZSOZ- + 2H+,

(1)

while the other involves ferric iron as the oxidant (GARRELS and THOMPSON, 1960): Fe& + 14Fe3+ + SH,O + 15Fe2+ f 2SO:- + 16H+.

(2)

Reaction (2) is rate-limited by the availability of Fe’+, which is generated by the oxidation of Fez+ with molecular oxygen (SINGER and STUMM, 1970): Fe’+ + 1/402 + H+ + Fe3’ + 1/2H20.

(3)

Both Eqns. (1) and (2) represent multi-step oxidations involving the transfer of 7 electrons per mole of sulfur. In reaction (I}, as it is written, the product sulfate derives 7;‘9 of its oxygen atoms from molecular oxygen and ‘la from water, while in Eqn. (2) the sulfate derives all of its oxygen atoms from water. However, since these are multi-step reactions, Eqns. (1) and (2) represent only the extreme possibilities for the source of the oxygen atoms in the sulfate product; as pointed out by VAN EVERDINGEN and KROUSE (1985) and TORAN and HARRIS (1989), they do not necessarily describe the actual source of those oxygen atoms. This depends upon the relative concentration and availability of the two oxidants under local environmental conditions, as well as upon the contributions of the various mechanistic pathways available to the oxidants under these conditions. In nature, these pathways can be bacterially mediated, and this can affect the source of oxygen. 1609

B. J. Reedy, J. K. Beattie, and R. T. Lowson

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the nature of the overall process but can give no insight into the actual mechanism of the reactions taking place. In this work, we present the results of new isotope tracer experiments on abiotic pyrite oxidation using high isotopic purity H2’80 and 1802. The pyrite is oxidised either in H2160 under “Oz or in Hz”0 under 1602_ Instead of using mass spectroscopy, the sulfate product from these reactions is analysed by both Fourier Transform Infrared (FTIR) and laser Raman vibrational spectroscopy. This enables the determination not only of the I80 content, but also of the relative amounts of sulfate isotopomers (S’60,‘80i_;) formed in the reaction, which might shed more light on the mechanism(s) involved as disulfide (S:-) is oxidised to sulfate. This kind of mechanistic information cannot be provided by mass spectroscopic techniques, such as those used by TAYLOR et al. ( 1984a,b) and BAILEYand PETERS(1976), which involve the destruction of the sulfate anion and yield only the ratio ‘60:‘80. Because the observed vibrational band is the totally symmetric v’(Al) stretching mode, it is insensitive to the isotope of the central sulfur atom. The natural abundance of 34S,moreover, is only 4%. The ability to reverse the source of the isotopes means that any isotopomer observed can be confirmed as a product of the reaction and not an impurity or sulfate which is the product of isotopic contamination. This is because any isotopomer formed in an ‘802/H2’60 experiment (e.g., S’603’802-) should be produced in isotopic reverse (S’60’80:-) in the analogous ‘602/Hz’80 experiment, within the limits ofany isotope fractionation effects. The use of near pure isotopic reagents also enables one to distinguish between sulfate formed on the pyrite surface prior to the experiments and that which is actually formed during the experiments. This serves as a useful test of the pyrite cleaning procedures. EXPERIMENTAL Pyrite for these experiments was obtained as pure massive specimens from Rum Jungle, N.T., Australia. These were hand-ground and seived to obtain the fraction with a grain size in the range 106180 pm. In order to ensure the removal of any oxidation products from the mineral surface, the pyrite was cleaned according to a procedure similar to that employed by MOSES et al. (1987). Samples of the pyrite were suspended in boiling 6 M HCI for at least 15 min, after which the HCl was replaced and the boiling procedure was repeated. The samples were then rinsed three times with additional boiling 6 M HCl and then three times with acetone. The pyrite was dried quickly in air and used in experiments within 15 min. The efficacy of this cleaning process will be discussed later. All of the isotopic tracer experiments were carried out in one of three specially constructed glass reaction manifolds. Each of these consisted of a small jacketed reaction vessel (about 5 mL in capacity) connected via a water condenser to a narrow bore vacuum line. A mercury manometer and a Toeplar pump were used to measure and control gas pressure within the system. “0 gas could be admitted to the manifold from a cylinder through a glass-metal seal. Before all experiments, the reaction vessel and the water condenser were rinsed thoroughly in acetone to render them free of bacteria, and dried in a nitrogen stream. Experiments with “0 2 The experiments involving ‘so2 gas and ordinary water solutions were carried out in the following way. The clean pyrite (0.1 g) and a magnetic stirrer bar were placed in the reaction vessel and covered with 2 mL of HCI or HC104 (0.1 M), or distilled water (all freshly boiled and degassed), depending on the required initial pH. (All pH

values given here are initial values measured at room temperature only.) In pH 1 experiments, if the presence of Fe(II1) was required, it was added as FeCI, - 6H20 (0.2 g). Although this is a sub-stoichiometric amount of Fe 3*, these reactions were alwavs stopped before all of the Fe3+was consumed. The entire reaction m>xturewas further degassed and then frozen quickly under Nz (using solid CO? or liquid Nz) to enable the connection (to the manifold) and subsequent evacuation of the reaction vessel. With the mixture still frozen, “02 gas (CEA-ORIS, Bureau des Isotopes Stables, France, 98% “0; or ICN Biomedicals, 98-99% “0) was introduced into the system until its pressure equalled or just exceeded the partial pressure of oxygen in air. Reactions were then run at either 70°C (for 2-3 days) or at the ambient laboratory temperature (-20°C) (for 5-6 days). The stirring rate was just sufficient to suspend some of the pyrite. In each case. this was sufficient time for enough sulfate to be produced for spectroscopic analysis. Experiments with Hz’s0 In the reactions where the pyrite was oxidised under air in H?“O solutions, only 0.05 g of pyrite and 1.0 g (0.9 mL) of Hz’s0 (ICN Biomedicals, 97-99%-‘*O) were used. For the pH 1 experiments, the latter was acidified usina concentrated HCI (12 M, 7 uL) or HCIOa (1 1.6 M, 7 PL) to give a-n H+ concentration bf 0.1 M.’when Fe(III) was required, it was added as anhydrous F&I,. The reactions were then run in a closed system as before, using carefully dried apparatus, but under ordinary air. Vibrational Spectroscopy At the completion of each experiment, the reaction mixture was quickly filtered (in air) using Millipore filters (0.22 pm pore size) to remove unreacted pyrite and any solid material (such as sulfur) produced during the experiment. For Raman spectroscopic analysis, this filtrate was concentrated by evaporation with an argon stream. Raman spectra were collected using a PC-driven Jobin Yvon Ramanor UlOOO monochromator with sample irradiation (-500 mW at 488 nm) by a Spectra Physics 2020 Ar+ laser into a quartz spinning cell. To produce barium sulfate for examination by FUR spectroscopy, some of the same solution of reaction products was acidified, if necessary, with 0.1 M HCl and then treated with barium chloride solution. The resulting precipitate was collected on a Millipore filter, washed several times with warm HCI (0.1 M), and then several times with distilled water before being dried at about 100°C. Samples were prepared by finely grinding BaS04 (3%) in spectroscopic grade KBr powder. FUR spectra were obtained at a resolution of 1 cm-’ with a Biorad/Digilab FTS 20/80 infrared spectrometer using the DRIFTS (diffuse reflectance) technique. The characteristic vi infrared frequencies for barium sulfate isotooomers are 984 cm-’ (S’60!m). 965 cm-’ (S’60~‘80z-). 954 cm-’ (Si602i80:-), 940 cm-’ (S’60i8&), and 929 cm” (S’BO:-); while the corresponding Raman frequencies for the sulfate anions in solution are 982, 965, 952, 939, and 927 cm-‘, respectively (REEDY et al., 1990, and refs. therein).

RESULTS AND DISCUSSION The presence of any of the five sulfate ‘60/‘80 isotopomers, S’60,,‘80’,~,, can be detected qualitatively by FTIR spectroscopy of the barium salt, a quick and convenient technique. The u’(A’) system in the 900-1000 cm-’ region of the spectrum of BaSOa will contain one major band for each isotopomer present, although these bands exhibit splitting due to symmetry effects which also prevent a quantitative determination of the relative isotopomer concentrations (REEDY et al., 1990). However, a quantitative determination can be made from Raman spectra of the sulfate solution. While this does not give an isotopic ratio as accurate or precise as that given by mass spectroscopy, that technique can give no information on isotopomer concentrations.

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Sulfate isotopomers in pyrite oxidation 929

Reactions without Added Fe3+ Figure 1 is the FTIR absorbance spectrum of the vi region of BaS04 from pyrite oxidised under “Oz gas in natural abundance water solution (pH 1, 7O”C, 60 h). Three bands are observed, a major band at 984 cm-’ and minor bands at 965 and 954 cm-‘, while the rest of the spectrum (not shown here) is typical of BaS04 with some “O-enrichment, with no significant bands which might be attributed to other species. This is important for the correct identification of the bands seen in the V, region since barium sulfite, for example, has weak infrared bands at 965 (the same as BaS’603’80), 946, and 935 cm-’ which could interfere with sulfate isotopomer bands. The bands at 965 and 954 cm-’ then are certainly u’(BaS’603180) and vl(BaS’602’802), respectively, while the main band at 984 cm-’ is obviously Vl(BaS’604). Therefore, while most of the sulfate formed by oxidation of the pyrite under these conditions derives all four of its 0 atoms from water, a significant amount derives either one or two 0 atoms from the I802 gas. This result is confirmed by the reverse experiment, namely the oxidation of pyrite under the same conditions, but using air (containing oxygen mainly as 1602) and Hz”0 instead. Here, it is expected that the FTIR spectrum of the product sulfate in the V’region will be the “mirror image” of that in Fig. 1; that is, it should show a large band at 929 cm-‘, the frequency for ul(BaS’804), and smaller bands at 940 (BaS’60’803) and 954 cm-’ (BaSL60zt802)-the isotopic reverse of the result in Fig. 1. Figure 2 shows the spectrum for this revem experiment, and, indeed, most of the sulfate again derives all of its 0 atoms from water (Hz”0 this time), with smaller amounts having one or two 0 atoms which come from oxygen gas ( 1602). This result is also important because no BaS1604 is observed, which demonstrates that the cleaning procedure employed did remove all significant amounts of sulfate that existed on the pyrite surface (as a natural oxidation

960 940 Wavenumbers FIG. 1. ITIR spectrum of BaS04 from oxidation of pyrite in H2160 under “02 gas (without added Fe’+).

/ I 960

I

940

Wavenumbers

I

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FIG. 2. ITIR spectrum of BaSO, from oxidation of pyrite in Hz’*0 under air (1602)(without added Fe’+).

product) before these experiments. This supports the choice made by MOSES et al. (1987) of this cleaning method for preventing sulfate “spikes” in time-course experiments on pyrite oxidation. It should be noted here that regardless of the source of the isotopes, the oxidation of Fe’+ to Fe3+ by molecular oxygen causes a “leak” of the molecular oxygen isotope into the water (which contains the other isotope) because the oxidation produces water containing oxygen from the atmosphere. However, this cannot be a significant contamination, because the relative molarities of O2 and water in the reaction could not be responsible for the production of mixed isotope isotopomers at the levels observed here. As mentioned previously, the relative concentrations of the isotopomers produced can be measured by Raman spectroscopy, but there are difficulties associated with this. Firstly, these relative concentrations do not remain constant over the course of the reaction. When two identical samples of pyrite are oxidised under identical conditions for different lengths of time, one day (24 h) and seven days, respectively, a time dependence of the relative isotopomer concentrations is observed. The concentrations of the minor isotopomers, S’603’802- and S’602’80:-, were greater relative to that of S’60:- after one day than they were after seven days. This is obvious from the FTIR spectra of the sulfate produced (as BaS04) by the two samples, but even after concentrating the reaction mixture from the first (one-day) run, the sulfate concentration was still too weak to measure the relative amounts of the three isotopomers by Raman spectroscopy. A rough estimate from the FTIR spectrum would be: 60% S’60:-, 30% S’603’802-, and 10% S’602’80:- (giving a total isotope content of about 88% 160, 12% “0); for the seven-day sample the concentrations are 8 1% S’60:-, 13% S’603’802-, and 6% S’602’80$- (giving a total of 94% I60 and 6% l8O) from Raman spectroscopy.

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There are a few possible reasons for the decrease in the concentrations of S’603’802- and S’602’80:~ relative to that of S’60:-. The most probable is simply the decrease in the amount of ‘*02 available as this is consumed without being replaced. That is, the partial pressure of oxygen in these experiments is not kept constant but allowed to decrease as the reaction proceeds. This change can be monitored on the manometer attached to the reaction manifold. Another possible reason is that S’603’802- and/or S’60z’“O:- are produced by reaction pathways which are important before enough Fe3+ is produced to make reaction (2) the dominant process. The homogeneous exchange of oxygen atoms between sulfate and water, which would lead to an increase in S’60$- at the expense of S’603’802- and Si602”O:-, does occur under conditions of high temperature and acidity (HOERING and KENNEDY, 1957; LLOYD, 1968; RADMER, 1972; CHIBA and SAKAI, 1985). However, under the conditions of the present experiments, no homogeneous exchange should occur. A similar exchange catalysed in some way by the pyrite surface was a possibility, so a blank experiment was conducted in which pyrite was heated (7O’C) and stirred in a solution of “O-enriched sulfate at pH 1 (0.1 M HC104) for seven days under an argon atmosphere. No change in the level of isotopic enrichment of the sulfate was detectable by vibrational spectroscopy, so it seems safe to assume that no significant exchange of sulfate oxygen takes place over the timespans of any of the tracer experiments. A further problem with measuring the relative amounts of the sulfate isotopomers produced by the pyrite oxidation experiments is that the concentrations are not exactly reproducible even under the same conditions and reaction times. This may be due to slight differences in experimental technique such as in the cleaning procedure or to air oxidation of intermediate sulfoxy ions (such as sulfite) after the conclusion of the ‘802/H2’60 experiments. If these problems could be overcome it would then be possible to conduct kinetic studies on the formation of the different isotopomers under a variety of conditions. The actual production of three different sulfate isotopomers in these tracer experiments points to the existence of more than one pathway for the formation of sulfate from pyrite oxidation in acid solution. However, because of the number of oxidation steps which must be involved (as S:- is oxidised to SOf), these pathways may not be independent of each other, but may involve common intermediate species or reaction steps. Sulfur species of intermediate oxidation state (between S:- and SO:-) that have been identified in pyrite reaction mixtures are elemental sulfur (SO, usually as S,) (NORDSTROM, 1982, and refs. therein), thiosulfate ion (S20:-), sulfite ion (SO:-), and polythionate ions (S,Oi-), especially tetrathionate (S,Oa-) and pentathionate (S,Oi-) (GOLDHABER, 1983; MCIQBBEN and BARNES, 1986; MOSES et al., 1987). These workers found that the dominant sulfoxy anion products from pyrite oxidation varied markedly with pH and that below pH 4 sulfate was the only significant sulfoxy species. Thiosulfate decomposes rapidly at low pH, to form elemental sulfur and sulfite. It is also oxidised very quickly by Fe3+ to give sulfate, with the extra oxygen coming from water. This is probably also the case with elemental sulfur and other intermediate species such as the polythionates

(MOSES et al., 1987). Sulfite is rapidly

oxidised by O2 to give sulfate. However, these species could still exist as short-lived intermediates in pyrite oxidation even at low pH. Both MOSES et al. (1987) and LUTHER (1987) proposed mechanisms for pyrite oxidation which involve the eventual formation of thiosulfate prior to oxidation to sulfate. In our tracer experiments, the oxidation or decomposition of intermediate sulfur species along different pathways probably leads to our observation of the three different sulfate isotopomers when the oxygen isotope in water is different to that in molecular oxygen. No mechanistic information about the formation of St60:- (the major isotopomer) is provided by these experiments, but there is no reason to assume that it is all formed by exactly the same pathway. One possibility is that it is produced by Fe3+ oxidation ofthiosulfate. The S’603’802- could be formed by ‘*02 oxidation of sulfite (possibly formed by the acid decomposition of thiosulfate). If that is the case, then it is not possible to look further back in the process, since SO:- exchanges oxygen atoms with water very rapidly and so will always reflect the isotopic composition of solvent water (BETTS and VOSS, 1970). Thiosulfate also exchanges oxygen with water, but not as quickly (PRYOR and TONELLATO, 1967). It is more difficult at this stage to suggest how the S’602’“O:- is formed, but this process is unlikely to go via sulfite. The possibility that S’603’802- and S’602’80:are formed (in the ‘s02/H2’60 experiment) by “02 oxidation of partly oxidised species that may have existed on the pyrite surface after cleaning, but before the commencement of the experiment, can be excluded, since the reverse experiment (‘602/H2’80) would not give the analogous isotopomers (S’60’80:and S’602’80:-, respectively) as it does, but S’60$- instead. (That is, any S1603 or S”j02 species already existing on the pyrite surface from air oxidation would form S’60$m if they were oxidised by 1602 in the reverse experiment. No S’60:m is detected, so such a process does not occur to any significant extent.) An isotopic tracer study of the oxidation of species such as sulfite, thiosulfate, and tetrathionate with Fe3+ and O2 as competing oxidants might help to explain the formation of different sulfate isotopomers in pyrite tracer experiments. Reactions with Added Fe3+ The FTIR spectrum (vl region) of BaS04 from an ‘*02/ H2160 pyrite oxidation tracer experiment to which Fe3+ was added as an oxidant is shown in Fig. 3. All other reaction conditions were identical to those in the experiments without Fe3+. Again, the major sulfate isotopomer formed is S’60i(984 cm-‘), but this time only S’603’802m (965 cm-‘) is observed as a minor product. This reaction was repeated several times over approximately the same length of time (in the range 64-69 hours) with the same qualitative result, but the concentration of S’603’802- relative to that of S’60:m was not always reproducible. For one oxidation, the isotopomer concentrations were measured to be 9 1.5% Si60:- and 8.5% S’603’802- after 64 h by Raman spectroscopy, but greater amounts of S’603’802- were observed in other runs. The reverse (‘602/02’80) of this experiment was only performed once, and the reaction had to be stopped after only 22 h for

Sulfate isotopomers in pyrite oxidation

1000

980

950

940

Wavenumbers

920

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FIG. 3. FTIR spectrum of BaSO, from oxidation of pyrite in H2160 under “02, with added Fe’+.

practical reasons (a breakage). The FTIR (v’) of BaSO, produced from the reaction mixture is shown in Fig. 4. As for the experiments without added Fe3+, reversing the sources of the two oxygen isotopes has reversed the pattern of the isotopomer u’ bands, so that now the major product is S’80:- (929 cm-‘), with S’60’80:- (940 cm-‘) as the minor product. However, a very small band at 955 cm-’ seems to indicate the presence of a small amount of S’602’80:-. The apparent lack of S’602’80:- after 64 h of reaction time in these experiments and its very low concentration (relative to the other isotopomers) after 22 hours could be because the added Fe3+ causes the production of the other two isotopomers to be much quicker than in the reactions without added Fe3+, and so the infrared band due to S’602’80:- is swamped relative to the others. The persistence, at a similar relative concentration (after the same length of time), of the S’603’R02- band in the ‘802/H2’60 experiment (or S’60’80:- in the ‘602/H2’80 experiment) when Fe3+ is added, indicates that its formation does involve Fe3+; if it was totally the product of oxidation by a competing oxidant (i.e., molecular oxygen), one might expect that increasing the availability of Fe3+ would reduce the amount of S’603’802produced relative to S’60$-. Again, it is important to remember that however it may be formed, any -SO3 intermediate produced is likely to have exchanged oxygen with water and to have assumed the oxygen isotopic composition of that water, unless it is oxidised to sulfate at a rate faster than its exchange rate.

20”C/pH 7. These pH values are the initial values only, since the systems studied were not buffered and the pH was allowed to drop during the course of each reaction as acid was produced. No Fe3+ was added to any of these reactions. The reactions conducted at 70°C and pH 7 yielded sulfate containing the same isotopomers (i.e., S%-, S’603’802-, and S’602’80~- in the ‘s02/H2’60 experiment) as observed for the 70”C/pH 1 experiment already discussed above. The relative isotopomer concentrations are also approximately the same, although it may not be meaningful to compare concentrations from experiments run for different lengths of time. This is an interesting result since at the higher initial pH it is expected that the dominant intermediate sulfoxy species would be different from and longer lived than those formed at low pH. This result suggests that the same overall pathway(s) are responsible for the production of sulfate at these higher pH values as at pH 1. This is despite the fact that Fe3+ solubility is very low above about pH 3. Further experiments using buffered solutions would be required in order to draw firm conclusions. The reactions conducted at 20°C (whether at pH 1 or pH 7 initially) produced sulfate with generally lower concentrations of the two minor isotopomers S’60~‘802- and S’602’80:m compared with the 70°C experiments, but it must be noted that the 20°C reactions were allowed to run for greater lengths of time. CONCLUSIONS In this work, pyrite oxidation in acid media has been studied through “0 tracer experiments performed using near isotopically pure O2 and H20 for the first time. The sulfate produced was analysed not by mass spectroscopy, but by vibrational spectroscopy. This enabled the identification for the first time of sulfate isotopomers formed during the oxidation 929

Experiments at Different pH Values and Temperatures

To ascertain the effects of varying pH and temperature in the experiments without added Fe3+, further ‘802/H2’60 pyrite oxidation runs (and in most cases ‘602/H2’80 reverse runs) were conducted in the same way with the following temperature/pH combinations: 70”C/pH 7; 20”C/pH 1;

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Wavenumbers FIG. 4. FTIR spectrum of BaS04 from oxidation of pyrite in Hz’*0 under 1602, with added Fe’+.

B. J. Reedy, J. K. Beattie, and R. T. Lowson

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of pyrite when the oxygen isotope in the molecular oxygen was different from that in the water of the reaction solutions. The results show that at 70°C and pH 1 the majority (up to about 90%) of the sulfate formed derives all four oxygen atoms from water. This is in basic agreement with the results of other workers, notably TAYLOR et al. (1984b), and points to mechanisms which involve Fe3+ as the sole or dominant oxidant. The results do not exclude the possibility, however, that O2 is the initial oxidant of pyrite, but that in subsequent rapid exchange reactions its isotopic signature is lost. Regardless of whether Fe3+ or O2 is the initial oxidant, the detection at significant levels of two sulfate isotopomers, with one and two oxygens respectively derived from molecular oxygen, indicates that mechanisms involving different intermediate sulfoxy species are also involved. This is even the case when Fe3+ is added at high levels to the reactions, with the isotopomer having one oxygen atom derived from the atmosphere still being formed at similar relative concentrations. The formation of the minor isotopomers seems to be depressed at lower temperatures (20°C). The major conclusions from this work have been verified by use of pure isotopes and pairs of experiments in which the oxygen isotopes are reversed. However, in a few replicate experiments, quantitative measurements of the relative isotopomer concentrations were not exactly reproducible. If this problem can be overcome, then kinetic studies of the timedependent isotopomer concentrations would be possible and might yield valuable mechanistic information. Isotope tracer experiments involving sulfate isotopomer determination have important implications for the study of the oxidation of sulfide minerals such as pyrite, which can undergo oxidation along a variety of different pathways, both inorganic and bacterial. In the latter case, oxygen isotope tracer experiments on pyrite oxidation might establish characteristic isotopomer signatures for bacterially mediated oxidation and even for specific strains of bacteria. Certainly, these experiments can aid in the interpretation of the isotope fractionation data collected by other workers and shed more light on the intermediate oxidation.

processes involved in sulfide mineral

Acknowledgments-B.J.R. gratefully acknowledges the receipt of a postgraduate studentship from the Australian Institute of Nuclear Science and Engineering. Editorial handling: G. Faure REFERENCES BAILEYL. K. and PETERSE. (1976) Decomposition of pyrite in acids by pressure leaching and anodization: the case for an electrochemical mechanism. Canadian Met. Quart. 15, 333-344.

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