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Journal of Environmental Quality 32:773-780 (2003)
© 2003 American Society of Agronomy, Crop Science Society of America, and Soil Science Society of America

TECHNICAL REPORTS
Ecosystem Restoration

Iron Sulfide Oxidation as Influenced by Calcium Carbonate Application

L. R. Hossner*,a and J. J. Doolittleb

a Dep. of Soil and Crop Sciences, Texas A&M Univ., College Station, TX 77843
b Plant Science Dep., South Dakota State Univ., Brookings, SD 57007

* Corresponding author (l-hossner{at}tamu.edu)

Received for publication March 5, 2002.

   ABSTRACT
TOP
 ABSTRACT
INTRODUCTION
 MATERIALS AND METHODS
 RESULTS AND DISCUSSION
 CONCLUSIONS
 REFERENCES
 
Two overburden materials, with different FeS2 contents (1.9 and 4.1%) and low acid neutralization potential, were limed with CaCO3 at rates of 0, 25, 50, 75, 100, and 125% based on the amount of CaCO3 needed to provide an acid–base account deficit (A/Ba) of zero (A/Ba = neutralization potential - potential acidity - exchangeable acidity). The limed overburden materials were inoculated with Thiobacillus ferrooxidans and leached weekly with deionized water. Residual FeS2 and CaCO3 were determined in samples over a 378-d period. Oxidation followed zero-order kinetics with respect to FeS2 concentration at pH values greater than 4 and first-order kinetics at pH values less than 4. Zero-order oxidation rates ranged from 0.0l to 0.46 µmol g-1 d-1 in the overburden with 1.9% FeS2 and from 0.01 to 0.22 µmol g-1 d-1 in the overburden with 4.1% FeS2. Oxidation following the first-order rate law had a first-order rate constant of 0.03 d-1 in the 1.9% FeS2 overburden and 0.01 d-1 in the 4.1% FeS2 overburden. The calculated half-life was 23 d for the 1.9% FeS2 overburden and 69 d for the 4.1% FeS2 overburden. Additions of CaCO3 affected FeS2 oxidation by controlling the pH of the system. Liming to greater than 50% of the acid–base account deficit did not significantly affect the zero-order oxidation rate. Dissolution of the applied CaCO3 was found to be faster than the oxidation of FeS2 at pH values greater than 4. It was projected that at lime rates up to 125%, the CaCO3 would dissolve and leach out of the system before all the FeS2 oxidized, leaving the potential for acid minesoil formation.


   INTRODUCTION
TOP
 ABSTRACT
 INTRODUCTION
MATERIALS AND METHODS
 RESULTS AND DISCUSSION
 CONCLUSIONS
 REFERENCES
 
SUCCESSFUL RECLAMATION of surface-mined lands can be accomplished using mixed overburden as a topsoil substitute. The overburden in this reclamation method is removed as a whole with a dragline or other mining equipment and piled on a previously mined area. The resultant mixed overburden spoil piles are leveled and revegetated. Mixed overburden can form minesoils that have better chemical and physical properties than pre-mine soils (Dixon et al., 1982; Bearden, 1984; Hons et al., 1978). Some overburden materials have the potential to become highly acidic after exposure to the atmosphere and oxidation of exposed iron sulfides (FeS2). Pyrite (cubic FeS2) was the major form of FeS2 in the overburden used in this study. Marcasite (orthorhombic FeS2) was also identified (Arora et al., 1978).

Sulfuric acid is a product of FeS2 oxidation and is largely responsible for decreased minesoil pH values. Acid minesoils can create problems for revegetation and discharge water quality. Research has shown that the rate of pyrite oxidation is influenced by parameters such as FeS2 morphology and surface area, microorganisms, and pH (Pugh et al., 1984; Smith and Shumate, 1970). The neutralization of acidity by lime is limited by the solubility of CaCO3 and by slow reaction kinetics (Geidel, 1979). Marcasite has been found to be more reactive than pyrite when all other parameters are held constant (Pugh et al., 1984; Temple and Delchamps, 1953; Wiersma and Rimstidt, 1984).

First-order kinetics indicate that the change in FeS2 concentration with time is a function of the FeS2 concentration at that particular time. This relationship is expressed mathematically in Eq. [1]:

[1]

The first-order rate constant (k) can be determined graphically from the integrated form of this equation:

[2]

The slope of the line produced when the natural logarithm of the FeS2 concentration is plotted against time is equal to the negative of the first-order rate constant. The intercept is equal to the natural logarithm of the FeS2 concentration at time zero, [FeS2]0. The half-life, the time where [FeS2] = 1/2[FeS2]0, can be calculated from Eq. [3]:

[3]

Zero-order kinetics mean the rate of FeS2 oxidation is constant over time and, therefore, independent of FeS2 concentration. A plot of FeS2 concentration versus time is linear. The slope of the line derived from the plotted data provides an estimate of the oxidation rate.

The abiotic oxidation of pyrite increases with increasing pH (Brown and Jurinak, 1989a,b; Smith et al., 1968). Abiotic oxidation has also been found to increase with increasing solution Fe3+ to Fe2+ ratio (Garrels and Thompson, 1960). Evangelou et al. (1983)( 1985) found that pyrite oxidation followed transport-controlled kinetics as described by Jurinak et al. (1977) for the dissolution of salt in a saline soil. Other researchers have found that abiotic oxidation of pyrite followed zero-order kinetics (Moses and Herman, 1991; Brown and Jurinak, 1989a,b; Vlek and Lindsay, 1978). These data showed that the production of Fe3+ and/or SO2-4 over time was linear and in agreement with the zero-order rate law. Pugh et al. (1984) found a linear relationship between time and the logarithm of SO2-4 production and concluded that the biotic oxidation of FeS2 followed first-order kinetics. Wiersma and Rimstidt (1984) also found that pyrite oxidation followed first-order kinetics but with respect to Fe3+ solution concentration.

When microbial oxidation is involved, researchers generally find that SO2-4 production from FeS2 oxidation follows first-order kinetics (Pugh et al., 1984; Stiller et al., 1989). Iron- and sulfur-oxidizing microorganisms associated with acid mine drainage, especially T. ferrooxidans, substantially increase FeS2 oxidation. The effect is more pronounced at pH values less than 4 (Arkesteyn, 1980; Clark, 1966; Evangelou et al., 1983; Lau et al., 1970; Lorenz and Tarpley, 1963; Pugh et al., 1984; Singer and Stumm, 1970; Taylor et al., 1984; Temple and Delchamps, 1953; Watzalf and Hammack, 1989). Pyrite oxidation has been found to be independent of pH at values less than 4 (McKibben and Barnes, 1986; Singer and Stumm, 1970).

Crushed limestone is the most common material used to raise soil pH and to correct the negative acid–base status of potentially acidic minesoil. Limestone is predominantly calcite (CaCO3), and may also contain some Mg-substituted calcite and/or dolomite [CaMg(CO3)2]. Inclusions of dolomite and Mg-calcite complicate thermodynamic and kinetic considerations because they have different solubility and reactivity in H2O than calcite (Morse, 1983; Thomas and Hargrove, 1984). The pH of a calcareous soil in H2O saturated with air ranges from 7.3 to 8.5 (Lindsay, 1979). Alkalinity (CO2-3, HCO-3, OH-) for the neutralization of acidity is produced as CaCO3 dissolves.

It has been traditional to treat CaCO3 dissolution kinetics empirically:

[4]
where Rd is the dissolution rate, kd is the rate constant, n is the empirical reaction order, and {Omega} is the saturation state = /Ksp. By fitting research data to the logarithmic form of this equation, Eq. [5], linear regression can be used to determine the reaction order from the slope and the rate constant from the intercept (Morse, 1983):

[5]

From a mechanistic point of view, there are two major divisions of rate control in CaCO3 dissolution. In solution far from equilibrium, transport control is generally accepted as more important. Near equilibrium, surface reactions are believed to be the rate-controlling mechanism. Constituents, such as Mg2+, Fe3+, H2PO-4, and HPO2-4, adsorb to the surface of CaCO3 and retard reaction rate. The transition between transport and surface control is not sharp and there is a transition zone where several reaction mechanisms share rate control (Morse, 1983). Rickard and Sjoberg (1983) investigated CaCO3 dissolution kinetics close to equilibrium. They determined that the rate was first-order and controlled by a mixture of surface and transport reactions.

Numerous researchers have documented the slowing of FeS2 oxidation with increasing pH (Nicholson et al., 1990; Clark, 1966; Evangelou et al., 1985; Kleinmann et al., 1981; Lorenz and Tarpley, 1963; Sturey et al., 1982). In all cases the addition of limestone increased the pH of the system to greater than 4. The reduction in FeS2 oxidation was attributed to the elevated pH being unfavorable and reducing the catalytic effect of T. ferrooxidans. Another effect of an elevated pH is the precipitation of iron oxides and a lower solution Fe3+ to Fe2+ ratio. Nicholson et al. (1990) observed iron oxide coatings precipitated directly on pyrite surfaces in carbonate-buffered solutions. The decreased rate of pyrite oxidation was attributed to slower diffusion of O2 through the oxide coating.


   MATERIALS AND METHODS
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
RESULTS AND DISCUSSION
 CONCLUSIONS
 REFERENCES
 
Three samples of overburden strata were excavated from an active mine highwall using a dragline. Each overburden sample was crushed to pass through a 1-mm sieve using a chipmunk ore crusher, mixed thoroughly, freeze-dried, and stored in sealed polyethylene bottles. Two of the overburden samples contained siderite (FeCO3) and were deemed unsuitable. To provide two overburden materials for this study, the FeS2 from one of the strata containing siderite was separated using tetrabromoethane (sg = 2.97) in heavy-liquid separation (Arora et al., 1978). The isolated FeS2 was thoroughly mixed into a portion of the overburden that did not contain siderite. Six kilograms of the two resultant overburden materials were crushed to pass a 0.25-mm sieve and thoroughly mixed using a roll-over, V-cone-shaped, twin-shell blender. The two overburden materials were then characterized for selected chemical characteristics, acid–base properties, and FeS2 morphologies. Exchangeable acidity was measured as outlined by Thomas (1982), neutralization potential by the method of Sobek et al. (1978), potential acidity from pyritic iron (American Society for Testing and Materials, 1968), cation exchange capacity by Na saturation (Chapman, 1965), and carbonate by the Bundy and Bremner method (1972). The two overburden materials will be identified as the 1.9% FeS2 overburden and the 4.1% FeS2 overburden based on the FeS2 percentage that was stoichiometrically calculated from pyritic iron (American Society for Testing and Materials, 1968).

Six rates of CaCO3 were applied to each overburden material. Reagent-grade CaCO3 was added to theoretically neutralize 0, 25, 50, 75, 100, and 125% of the acid–base account deficit (Table 1).


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  		Table 1. Lime treatments and projected acid–base values.

 
Ten grams of treated overburden were diluted with 5 g of pure quartz sand and packed into a polyethylene leaching vial (Doolittle and Hossner, 1997). The sand was added to improve aeration and drainage. Leaching vials were constructed from 20-mL scintillation vials. Four 3-mm-diameter holes were drilled in the bottom of each vial to allow drainage. The holes were covered with a 2.1-cm glass fiber filter to retain particles greater than 2.7 µm. After the overburden treatment was lightly packed into the vial, the surface was covered with a thin layer of fine glass wool to help absorb the impact of added water and prevent surface sealing and crusting. The treated overburden was inoculated with more than 109 active cells of T. ferrooxidans and brought to field capacity with deionized water. Each treatment was replicated 64 times.

Overburden in each vial was leached with 10 mL of deionized water at 7-d intervals for 378 d. This is equivalent to an annual infiltration of 1079 mm of water. The amount of water applied was based on an average annual rainfall of 1218 mm at a meteorological monitoring site (Dugas, 1983) located near the surface mine where the overburden samples were collected. A 50 kPa vacuum was applied to speed the leaching process.

Two replications of each treatment in the leaching vials were removed from the system at regular intervals. A saturated paste was made of the sample within the vial. The pH of the paste was measured using a combination glass electrode. Samples were then freeze-dried immediately. The replications were removed at 7-d intervals for the first 91 d and then at 14-d intervals for the remainder of the study. The leaching cycles were continued for 378 d. Freeze-dried samples were analyzed for pyritic iron (American Society for Testing and Materials, 1968) and residual carbonate (Bundy and Bremner, 1972).

The FeS2 and other heavy minerals in overburden samples were isolated from light minerals using heavy-liquid separation (Arora et al., 1978). Scanning electron microscopy was conducted using a T330A scanning electron microscope (JEOL, Akishima, Japan). Selected particles were also analyzed using a JEOL JSM35 scanning electron microscope equipped with a Tracor Northern energy dispersive X-ray analyzer (EDAX) (Thermo Noran, Middleton, WI).


   RESULTS AND DISCUSSION
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
 RESULTS AND DISCUSSION
CONCLUSIONS
 REFERENCES
 
X-ray diffraction identified pyrite and marcasite as the only heavy minerals and iron sulfides present in the two overburden samples (Table 2). Chemical analyses indicated an initial pH of 6.2 and a negative acid–base account. There was good agreement in potential acidity measured by the hydrogen peroxide method (O'Shay et al., 1990) and the method of the American Society for Testing and Materials (1968). In addition there was excellent agreement between calculated values for acid–base account and acid–base balance.


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  		Table 2. Selected chemical properties and acid–base characteristics of the two overburden samples.

 
Residual Iron Sulfide and Carbonate
The general trend for residual iron sulfide and carbonate data on samples containing 1.9 and 4.1% FeS2 was similar. All of the inherent CaCO3 was dissolved from the 0% acid–base deficit treatment within 25 d (Fig. 1) . Most of the FeS2 was oxidized within 100 d. The FeS2 and CaCO3 contents of the 100% acid–base deficit treatment decreased by only a small amount over 378 d.



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  		Fig. 1. Residual FeS2 and CaCO3 in 1.9% FeS2 overburden treated with CaCO3 to neutralize 0 or 100% of the acid–base account deficit (A/Ba).

 
Rate of Iron Sulfide Oxidation
Oxidation of FeS2 in overburden material followed two different rate laws depending on the pH of the system. In general, when the pH was less than 4, FeS2 oxidation followed first-order kinetics. At pH values greater than 4, FeS2 oxidation followed zero-order kinetics.

The pH of the 0% acid–base deficit overburden treatment dropped to less than 4 within the first 7 d. Residual FeS2 concentration data fit a first-order plot (Fig. 2) . The first-order rate constant by the least square regression line was estimated to be 0.03 d-1. The half-life of FeS2 in this treatment was calculated to be 23.1 d. Iron(III) remains in solution when pH is less than 4 and is the primary oxidant of FeS2; therefore, the measured apparent first-order rate constant is for the Fe3+ mechanism of FeS2 oxidation catalyzed by T. ferrooxidans.



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  		Fig. 2. First-order plot of the loss of FeS2 in the 1.9% FeS2 overburden treated with CaCO3 to neutralize 0% of the acid–base account deficit.

 
Oxidation of FeS2 in the 25% acid–base deficit treatment followed zero-order kinetics for the first 100 d for the 4.1% FeS2 overburden (Fig. 3) . There was a significant drop in FeS2 content between 100 and 378 d. The pH of the system was less than 4 after 100 d and oxidation was described by first-order kinetics.



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  		Fig. 3. Zero-order (0–100 d) and first-order (100–378 d) plots of the loss of FeS2 in the 4.1% FeS2 overburden treated with CaCO3 to neutralize 25% of the acid–base account deficit.

 
The pH of the 25% acid–base deficit treatment for the 1.9% FeS2 treatment remained greater than or near 4 and FeS2 oxidation was described by two zero-order equations (data not shown). During the first 200 d, the pH was near 6. Linear regression estimated the oxidation rate of FeS2 at 0.09 µmol g-1 d-1. A rate of 0.46 µmol g-1 d-1 was observed during the next 100 d when the pH decreased from 5 to 4, with a few replications decreasing to less than 3. The relatively low solubility of Fe3+ at pH values measured in this treatment indicated that the primary oxidant was probably O2.

All the treatments that were limed to 50% of the acid–base deficit or greater had a significantly slower FeS2 oxidation rate that followed zero-order kinetics (Fig. 4) . There was no statistically significant difference between reaction rates for these CaCO3 treatments ({alpha} = 0.05). The average FeS2 oxidation rate of the 50, 75, 100, and 125% lime treatments was 0.01 µmol g-1 d-1. The pH of all treatments remained much greater than 7, indicating O2 was the primary oxidant. Extrapolation of these data indicated it would take approximately 30 yr for FeS2 to oxidize in this pyritic overburden if the pH remained greater than 4.



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  		Fig. 4. Zero-order plot of the loss of FeS2 in the 4.1% FeS2 overburden treated with CaCO3 to neutralize 50, 75, 100, and 125% of the acid–base account deficit (% A/Ba).

 
When overburden pH values are less than 4, Fe3+ is more soluble and considered to be the primary oxidant of FeS2 (Clark, 1966; Singer and Stumm, 1968; Smith et al., 1968; Wiersma and Rimstidt, 1984). The major difference between this research and other studies is the natural overburden matrix in which this oxidation study was conducted. Published oxidation rates cited in Table 3 were determined in stirred solutions (Wiersma and Rimstidt, 1984) or in solutions with gases containing O2 bubbled through them (Pugh et al., 1984). A moist overburden matrix was chosen for this research to more closely simulate the conditions found at a surface-mine reclamation site with moderate rainfall. Reactions in an agitated suspension would be expected to occur more rapidly since reactants, products, and microorganisms are more mobile and have a higher probability of interacting. Wiersma and Rimstidt (1984) observed the fastest first-order oxidation rate, but their experimental design did not include T. ferrooxidans, which Singer and Stumm (1970) reported to increase FeS2 oxidation by 106–fold. They achieved these relatively high oxidation rates using an acid (pH = 2) solution with an initial concentration of 10-3 M Fe3+. The catalytic effect of T. ferrooxidans converting Fe2+ to Fe3+ would not be significant during the relatively short period of time this system was observed.


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  		Table 3. Comparison of observed zero-order FeS2 and first-order oxidation rates of this study with other published rates.

 
First-order FeS2 oxidation rates reported by Pugh et al. (1984) are an order of magnitude slower than those reported by Wiersma and Rimstidt (1984), but are an order of magnitude faster than those measured in this study. The experiment of Pugh et al. (1984), in which FeS2 was added to an acidified (pH = 3) suspension of T. ferrooxidans with bubbled air for 14 d, was somewhat longer and closer to a natural overburden system than that of Wiersma and Rimstidt (1984).

First-order FeS2 oxidation rates measured in this research are more applicable for modeling FeS2 oxidation in natural overburden systems (pH < 4) in udic moisture regimes than previously published rates. Data from Wiersma and Rimstidt (1984) illustrate maximum possible FeS2 oxidation rates under extremely acid conditions and that sedimentary FeS2 is 4 to 7 times more reactive than hydrothermal FeS2. Pugh et al. (1984) demonstrated the relative reactivity of sedimentary FeS2 polymorphs to be: marcasite > framboidal pyrite > massive pyrite.

When expressed on a pure FeS2 basis (Table 3), the zero-order FeS2 oxidation rates measured in this research were consistent with rates reported by Vlek and Lindsay (1978). Thiobacillus ferrooxidans is relatively inactive at pH values greater than 7 (Manning and Cook, 1972), therefore, a significant difference in oxidation rate due to its presence would not be expected. Oxygen is the primary oxidant of FeS2 at pH values greater than 4.5 (Clark, 1966; Singer and Stumm, 1968; Smith et al., 1968; Wiersma and Rimstidt, 1984). Interaction between O2 and FeS2 in an overburden matrix and an aerated sand–water suspension must be similar since the observed oxidation rates are comparable.

The zero-order FeS2 oxidation rate was observed to increase as the overburden pH decreased from a range of 8 to 7 to a range of 6 to 4. This was contrary to the observation of Brown and Jurinak (1989a) (Table 3); however, their data were from sterile, stagnant solutions containing pyrite. The difference was probably due to increased T. ferrooxidans activity in the inoculated overburden as the pH decreased. It was also interesting to note that Brown and Jurinak (1989a) observed zero-order FeS2 oxidation at a pH of 3.5. In the current study, iron sulfide oxidation was found to follow first-order kinetics when pH was less than 4. The difference may again be due to T. ferrooxidans, but Pugh et al. (1984) and Wiersma and Rimstidt (1984) observed first-order kinetics for FeS2 oxidation in sterile solution at low pH.

Rate of Calcium Carbonate Dissolution
The total CaCO3 content gradually decreased with time in each of these treatments (Fig. 5) . The dissolution of the added CaCO3 from both the 1.9 and 4.1% FeS2 overburden materials was linear with time. Plummer et al. (1978) and Rickard and Sjoberg (1983) reported that the dissolution of CaCO3 in aqueous media, far from saturation, was first-order. This research does not confirm or conflict with that information. The loss of applied CaCO3 was determined by measuring residual CaCO3 in the overburden matrix. Dissolution of CaCO3 into the water front of each leaching cycle most probably followed first-order kinetics; however, this could not be determined from the data generated in this study. The observed linear dissolution of CaCO3 from the overburden matrix (Fig. 6) was consistent with most carbonate minerals in soil.



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  		Fig. 5. Dissolution rate of CaCO3 in the 4.1% FeS2 overburden treated with CaCO3 to neutralize 25, 50, 75, 100, and 125% of the acid–base account deficit (A/Ba).

 


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  		Fig. 6. Scanning electron micrographs of representative weathered pyrite particles separated from the overburden following 378 d of incubation and initially containing 4.1% FeS2. (A) Weathered pyrite particle from the 25% acid–base account deficit treatment. (B) Weathered pyrite particle from the 0% acid–base account deficit treatment. (C) Large weathered pyrite particle from the 25% acid–base account deficit treatment. (D) Large pyrite particle from the 125% acid–base account deficit treatment. There is no apparent evidence of pitting and the morphology is similar to pyrite particles at time zero.

 
The rate of dissolution of CaCO3 appears to decrease with increasing lime rate; however, there is no significant difference ({alpha} = 0.05) between the slopes (Fig. 5). A difference in CaCO3 dissolution rates between the 50, 75, 100, and 125% acid–base deficit treatments would not be anticipated based on the CaCO3 dissolution mechanism contributions described by Plummer et al. (1978). Since FeS2 oxidation was very slow, these treatments were under conditions where H2O is primarily responsible for CaCO3 dissolution (pH > 5.6, PCO2 = 10-3.53 MPa). All the treatments were leached with the same amount of H2O; therefore, the same amount of CaCO3 should have dissolved with each leaching cycle. The CaCO3 in the 25% acid–base deficit treatment should have dissolved significantly faster because of a the more rapid FeS2 oxidation rate and lower pH.

The acid–base deficit of treatments receiving higher rates of CaCO3 (lime rate = 50, 75, 100, and 125% of the acid–base deficit) was more negative with time. This decrease indicated that neutralization potential in the form of CaCO3 was leaving the system faster than the FeS2 was releasing H2SO4. Therefore, each of the overburden materials treated with CaCO3 would eventually become acidic. This trend was consistent with similar studies by Caruccio (1978), Caruccio et al. (1981), and Geidel (1979). These authors reported that the oxidation of FeS2 did not change with flushing frequency; however, the dissolution of indigenous and added CaCO3 increased with flushing frequency.

A prediction of when each treatment would become acidic is shown in Table 4. These data were calculated using the FeS2 oxidation rates and CaCO3 dissolution rates measured at the termination of this project and assumed that reaction rates did not change as the pH of the system decreased. Iron sulfide oxidation and CaCO3 dissolution increased as pH decreased; therefore, these estimates are conservative and probably overestimate the length of time for CaCO3 and FeS2 removal. By extrapolating the CaCO3 dissolution rates from the 125% acid–base deficit treatment of the 1.9% FeS2 overburden, it would take approximately 9.9 yr for all the applied lime to dissolve. This is 19.4 yr short of the life expectancy of the FeS2 in the system. Therefore, 54% of the FeS2 will still remain when the CaCO3 is gone and acid mine soil will eventually form. The same trend is shown in the 125% acid–base deficit treatment of the 4.1% FeS2 overburden, but only a 3-yr difference is predicted. After the CaCO3 is dissolved out of this overburden material, 23% of the initial FeS2 is predicted to still remain. Since dissolution by infiltrating water limits the residence time of the applied CaCO3 in these treated overburdens, a possible alternative would be to increase the lime rate and/or use a lime source that is predominantly dolomite. Dolomite [CaMg(CO3)2] has a significantly slower reaction rate and is considered to be relatively unreactive in percolating water (Chou et al., 1989; Morse, 1983). The slower reactivity would give the lime a longer residence time in the overburden.


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  		Table 4. Extrapolation of residual FeS2 and CaCO3 based on respective oxidation and dissolution rates measured at the termination of the experiment.

 
Scanning Electron Micrographs of Weathered Pyrite
The residual FeS2 was separated from the treated overburden matrix at the end of the oxidation study using heavy-liquid separation techniques (Arora et al., 1978). The yield of FeS2 was consistent with treatment results described earlier. It was difficult to find residual FeS2 particles in the 0% acid–base deficit treatments; however, the high-lime treatments yielded almost as much FeS2 as the time-zero samples.

Scanning electron micrographs of a characteristic cross-section of pyrite particles separated from the overburden mixtures leached for 378 d are shown in Fig. 6. Energy dispersive X-ray analysis confirmed that only particles containing iron and sulfur were photographed.

Residual marcasite or framboidal pyrite particles were not identified in the heavy-liquid separates from any of the leached treatments. The overall size distribution of the residual pyrite particles was larger than in the fresh overburden. The size distribution between the time zero and residual pyrite particles was not quantified, but during visual observations very few pyrite particles less than 25 µm in diameter were observed. It was apparent that the more reactive iron sulfides, marcasite, framboidal pyrite, and small pyrite particles, oxidized more quickly, leaving large, low surface area residual pyrite particles. Surface morphology of pyrite separated from the fresh overburden samples and pyrite separated following 378 d of incubation from CaCO3 treatments with greater than 50% acid–base deficit was similar (Fig. 6D).


   CONCLUSIONS
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
 RESULTS AND DISCUSSION
 CONCLUSIONS
REFERENCES
 
Additions of CaCO3 affected FeS2 oxidation by raising the pH of the system and slowing the rate of FeS2 oxidation. Liming to 25% of the acid–base deficit delayed the time when the overburden pH decreased to less than 4; therefore, increasing the time required for FeS2 oxidation to change from zero-order to the first-order rate law. The addition of lime did not affect the subsequent half-life of FeS2 after the pH decreased to less than 4. Liming to greater than 50% of the acid–base deficit did not significantly affect the FeS2 zero-order oxidation rate over the 378 d of this experiment.

Dissolution of added CaCO3 increased with increasing FeS2 oxidation rate. Dissolution of CaCO3 appeared to decrease with increased lime application rate, but there was no statistically significant difference ({alpha} = 0.05) between the measured rates. The dissolution of CaCO3 was faster in the 50, 75, 100, and 125% acid–base deficit overburden than the release of potential acidity from FeS2. When the difference in the rate of FeS2 oxidation between potential acidity and carbonate alkalinity of the 125% acid–base deficit treatments was extrapolated, there was a 3- to 19-yr difference between the time when all of the applied CaCO3 is dissolved out of the system and all the FeS2 is oxidized. This trend is also expressed in the acid–base deficit measurements that become more negative with time. Therefore, each of the CaCO3–treated overburden materials would eventually become acidic.

Liming overburden or topsoil substitutes containing potential acidity to an acid–base deficit of zero may not be the best long-term solution for the prevention of acid mine spoil and acid mine drainage in moderately high rainfall environments. The lime treatment will delay the rapid oxidation of FeS2. A less soluble carbonate source or higher rates of lime application will increase the residence time of the applied limestone. Long-term studies and field research are needed to substantiate this conclusion.


   REFERENCES
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
 RESULTS AND DISCUSSION
 CONCLUSIONS
 REFERENCES
 


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