Environ. Sci. Technol. 2000, 34, 2288-2292
Potential Inhibition of Acid Formation in Pyritic Environments Using Calcium Sulfite Byproduct
SO32- + 1/2O2 f SO42-
YUELI HAO AND WARREN A. DICK* School of Natural Resources, The Ohio State University, 1680 Madison Avenue, Wooster, Ohio 44691
Pyrite (FeS2) oxidation in coal mine spoils and coal refuse produces copious amounts of acid that can cause severe environmental damage. Prevention of acid formation by calcium sulfite (CaSO3), a flue gas desulfurization (FGD) byproduct of the scrubbing reaction used to remove sulfur dioxide from flue gases when coal is burned, was studied. First-order rate constants of CaSO3 oxidation by O2 ranged from approximately 0.0164 min-1 (initial pH 8.0, 25 °C, 21% O2) to 0.3807 min-1 (initial pH 5.0, 15 °C, 21% O2). Overall, the rate constants dramatically increased when the initial pH decreased from 6.5 to 5.0, the temperature increased from 5 to 15 °C, and the O2 content in air increased from 0.2% to 21%. When 6.4 g of CaSO3 or 10 g of CaSO3containing FGD byproduct were mixed with 50 g of coal refuse in small columns (2.5 × 13 cm), the total acidity in the leachate water collected during 27 weeks was reduced by 45% and 64%, respectively, as compared to the untreated control. Mixing CaSO3 with coal mine spoil or coal refuse before extensive oxidation occurs reduces acid production because the CaSO3 reacts with any O2 introduced into the spoil or refuse materials before it can react with pyrite to form acid.
Introduction Surface mining of coal often exposes pyritic spoil material and washing of coal, to remove sulfur, also creates refuse materials with enhanced concentrations of pyrite. Pyrite oxidation can be described using the following three reactions (1):
FeS2 + 3.5O2 + H2O f Fe2+ + 2SO42- + 2H+
(1)
2Fe2+ + 0.5O2 + 2H+ f 2Fe3+ + H2O
(2)
FeS2 + 14Fe3+ + 8H2O f 15Fe2+ + 16H+ + 2SO42- (3) Acid mine drainage produced as a result of pyrite oxidation can cause severe damage to the environment. Minimizing contact of O2 with pyrite-containing materials is a common approach to inhibit pyrite oxidation (1, 2). However, complete removal of O2 is difficult. This may be possible if we mix the pyrite-containing materials with a readily available strong reductant. Sulfite is a strong reductant and has been used for many years as an O2 scavenger for the preservation of foods and drugs (3, 4). The reaction of sulfite with O2 to produce sulfate: * Corresponding author phone: (330)263-3877; fax: (330)263-3658; e-mail:
[email protected]. 2288
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has a large equilibrium constant (K ) 6.66 × 1040 at 25 °C and 1 atm) (5), and it is also rapid with first-order rate constants ranging from 23 to 38 s-1 at pH < 4 and with second-order rate constants ranging from 8 × 104 to 2.1 × 104 M-1 s-1 when the temperature is varied between 20 to 50 °C and pH > 4 (6, 7). The sulfite salts used for food and drug preservation include Na2SO3, K2SO3, and (NH4)2SO3. These salts are all relatively expensive and only used in small quantities. Calcium sulfite is produced in large quantities during the desulfurization of coal combustion flue gases in many coalpowered electricity-producing utilities (8). This CaSO3 can exist either as a rather pure material or as a component of flue gas desulfurization (FGD) byproduct material, which is generally a mixture of fly ash, unreacted sorbent (generally lime or limestone), and CaSO3/CaSO4. Most of the FGD is currently disposed of in landfills. Until recently there was little incentive to develop beneficial uses of this FGD byproduct. However, the increasing volume of FGD produced and associated landfilling costs have stimulated research on potential uses of FGD (8, 13) that can provide economic benefit to both the producer of the FGD and the end user. We propose to mix FGD containing CaSO3 with coal mine spoil or coal wash refuse to remove O2, thus inhibiting pyrite oxidation and acid production. Unlike the other sulfite salts, CaSO3 has a low solubility in water (0.054 g L-1 at 25 °C), and its dissolution rate greatly depends on pH and temperature (9, 10). This may limit its ability to scavenge O2 in water. However, little work has been done using experimental conditions commonly found in coal mine spoil or coal wash refuse where pyrite is present. To pyritic environments, it is possible to mix scrubber sludge to create heterogeneous conditions with high concentrations of CaSO3. The reaction conditions in these environments will thus vary greatly with time and location because of variations in pH, temperature, and the amount of O2 and sulfite. To determine if CaSO3 or a CaSO3-containing FGD can be added to coal wash refuse to inhibit pyrite oxidation, the oxidation rates of CaSO3 by O2 needs to be measured under a broad range of conditions commonly encountered in coal wash refuse disposal environments. The purpose of this study was to determine the ability of CaSO3 to react with O2 under various conditions expected to be found at a coal refuse disposal site. This information provides estimates of the potential for CaSO3-containing FGD to be mixed with coal wash refuse and thus inhibit acid formation and improve water quality that originates from such sites.
Experimental Section Synthesis of Calcium Sulfite. Calcium sulfite was synthesized by mixing 6% sulfurous acid with calcium carbonate. The excess sulfurous acid was removed by heating gradually in a water bath to a temperature not more than 90 °C. The precipitate of CaSO3 was washed with a large quantity of double-deionized water and then dried at 60 °C. Thermogravimetric analysis indicated that the product contained more than 96% CaSO3, which was also ascertained to be hannebachite (CaSO3‚0.5H2O) by X-ray diffraction analysis. Oxidation of CaSO3 in Aqueous Slurry. Oxidation Rate. A slurry was prepared by mixing CaSO3 with water. To ensure that reactions were not catalyzed by bacteria, the slurries were autoclaved at 121 °C for 20 min. While working in a 10.1021/es9904235 CCC: $19.00
2000 American Chemical Society Published on Web 04/25/2000
FIGURE 1. Effect of initial pH on oxidation of CaSO3 in an aqueous slurry at 21% O2 and 25 °C. Error bar ) 1 SD. The rates follow first-order kinetics with rate constants shown in Table 1. sterile hood, 1 mL of this slurry was placed in each well of a 12-well microplate (Becton Dickinson & Co., Lincoln Park, NJ). Concentrated (18 M) H2SO4 or 10 M NaOH was added to the slurries to obtain the desired pH values, and the zero time of reaction was considered as the instance the acid or based was added. The microplates were shaken on a reciprocal water bath shaker at 80 rpm. The initial experimental conditions were as follows: (i) CaSO3 concentrations (g L-1) of 10, 20, and 40; (ii) pH of 2.0, 3.5, 5.0, 6.5, and 8.0; (iii) incubation temperature (°C) of 5, 15, and 25; and (iv) O2 content in air of 0.2% and 21%. Under alkaline pH conditions, the experiments could last more than 1 month, and sterile double-deionized water was added weekly to replace that lost by evaporation. Oxidation of CaSO3 was conducted in an anaerobic chamber (0.5 m3) in which O2 content in the gas phase was adjusted to 0.2% by flushing with N2 before the chamber door was tightly closed and the experiment started. The O2 content in the gas phase was controlled by intermittent flushes of N2. The O2 content was measured at the start and every 3 h during the experiment by injection of a 5-mL gas sample from the anaerobic chamber into a gas chromatograph (Water Dimension Gas Chromatography, electron capture detector, Tremetrics Inc., Austin, TX) (11). Standard concentrations of O2 gas were prepared by mixing O2 and N2 gases at various ratios. To determine the oxidation rate of CaSO3, total CaSO3 remaining was analyzed (in triplicate) by addition of excess 1.0 M iodine solution and then back-titrated with 0.1 M sodium thiosulfate solution using starch as an indicator. The oxidation kinetics were described as either first-order reaction with the differential equation -dc/dt ) kc or zero-order reaction with the differential equation -dc/dt ) k, where c is CaSO3 concentration at time t and k is the rate constant. The k value was obtained by curve fitting or linear regression of the data obtained in the experiment using SigmaPlot (version 3.0, Jandel Scientific, San Rafael, CA). Concentration of Dissolved O2. This experiment was conducted to determine the level of O2, available to drive pyrite oxidation, found in the presence of CaSO3 and at various pH values. Dissolved O2 concentration in 100 mL of 10 g L-1 CaSO3 aqueous slurry was determined using a YSI 5905 BOD probe and a YSI M58 dissolved oxygen meter (Yellow Springs Instrument Co., Inc., Yellow Springs, OH). The O2 probe and meter have a detection limit of 0.01 mg L-1. The pH of the slurry was initially adjusted to pH 12 with
TABLE 1. Rate Constants and Regression Coefficients of First-Order Kinetic Equations of CaSO3 Oxidation by O2 in an Aqueous Slurry O2 content (%)
temp (°C)
21
25
21
25
21
25
21
15
0.2
25
initial CaSO3 concn (g L-1)
initial pH
rate constant (min-1)
r2
10 10 10 20 20 20 40 40 40 10 20 40 10
5.0 6.5 8.0 5.0 6.5 8.0 5.0 6.5 8.0 5.0 5.0 5.0 5.0
0.2511 0.0918 0.0164 0.3179 0.0604 0.0963 0.1898 0.0252 0.0330 0.3807 0.3317 0.1153 0.0334
0.85 0.88 0.60 0.98 0.80 0.97 0.93 0.52 0.95 0.96 0.97 0.93 0.89
10 M NaOH and then lowered to pH 5 by about one unit at a time by addition of concentrated (≈18 M) H2SO4, and dissolved O2 was determined. Release of SO2. Total S in the slurries was determined at the start and at the end of the reactions. The CaSO3 was first treated with iodine solution to convert all sulfite to sulfate, and then concentrated (≈70%) HNO3 was added before analyses by inductively coupled plasma-atomic emission spectroscopy (ICP-AES). Differences in total S between the start and the end of the reaction were attributed to SO2 release. Release of SO2 was confirmed by a color change in a solution that contained 5% glycerol, 0.1 M NaOH, and color reagents (formaldehyde and rosaniline) (12). This solution was placed inside a closed bottle containing the CaSO3 slurry. Treatment of Coal Refuse with FGD. Unweathered (fresh) coarse coal refuse was obtained from the American Electric Power Plant (Coshocton, OH). Approximately 36% of the refuse had particle sizes ranging between 1 and 10 mm in diameter, and 64% were greater than 10 mm in diameter. It was sieved through a screen with 1.25-cm openings, airdried, and ground into particles less than 2 mm. Total S content in the coal refuse was 115 g kg-1 (11.5%) as determined by an aqua regia-HF digestion and inductively coupled plasma (ICP) emission spectroscopy (13). X-ray diffraction analyses indicated that the fresh coal refuse VOL. 34, NO. 11, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 2. Effect of temperature on oxidation of CaSO3 in an aqueous slurry at initial pH 5.0 and 21% O2. Error bar ) 1 SD. At 15 and 25 °C, the rates follow first-order kinetics with rate constants shown in Table 1. At 5 °C, the rates follow zero-order kinetics. contained gypsum, kaolinite, marcasite, pyrite, and quartz (13). Pyrite content of the coal refuse was analyzed by the method of Begheijn et al. (14) and was 134 g kg-1, which was equivalent to 97.8 g kg-1 S (85% of total S). Plastic syringes (2.5 cm i.d. and 13 cm long) were used as columns. There were three treatments with each having 3 replicates (total 9 columns). The control columns contained 50 g of ground coal refuse. The columns containing the CaSO3 treatment were a mixture of 50 g of ground coal refuse and 6.4 g of CaSO3. The columns of FGD byproduct treatment contained a mixture of 50 g of ground coal refuse and 10 g of FGD byproduct containing an equivalent amount (6.4 g) of CaSO3, 1.0 g of CaCO3, 1.7 g of fly ash, and 0.9 g of CaSO4‚ 2H2O (15). The FGD was obtained from the American Electric Power Plant located near Coshocton, OH, and was weathered for several weeks at the disposal site near the power plant. Before use, the FGD was dried at 105 °C for 24 h. Glass wool was applied at the bottom (1 g) and the top (0.2 g) of the syringes to facilitate water movement through the columns. The columns were leached weekly with 20 mL of doubledeionized water for the first 13 weeks and every 2 weeks for the last 14 weeks (total 27 weeks). The leaching rate was 1 mL/h and was controlled using a mechanical extractor (model 24-01, Centurion International Inc., Lincoln, NC). Leachates were analyzed for pH (glass electrode) and titratable acidity (end pH 8.3). The total amounts of acidity were calculated for the weekly (weeks 1-13, 13 leachates) and every 2 weeks (weeks 13-27, 8 leachates) periods and also for the entire period (weeks 1-27, 20 leachates). Data Analyses. First-order rate constants or zero-order rates of CaSO3 oxidation were determined using SigmaPlot 3.0 (Jandel Scientific, San Rafael, CA). For the column leaching experiment, total acidity calculated for the entire leaching period (27 weeks) was used to perform t-tests in order to compare the effects among the various treatments. Significant differences were assigned based on a 95% level of confidence (p < 0.05).
Results and Discussion Oxidation Rate of CaSO3 in an Aqueous Slurry by O2. Concentrations of CaSO3 decreased rapidly when oxidation was allowed to proceed under normal atmospheric conditions (≈21% O2) and a temperature of 25 °C. This decrease was greatly affected by the initial pH (Figure 1, note different time scales). The pH changed only slightly (generally less 2290
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than 0.2 pH unit) during the time measurements were made. At an initial pH of 5.0, the concentration of CaSO3 after 24 h of oxidation dropped below 0.5 g L-1 regardless of the initial concentrations of CaSO3. When the initial pH of the CaSO3 slurry was 6.5 or 8.0, the concentration of CaSO3 after 14 days of oxidation when the initial concentrations of CaSO3 were 10, 20, and 40 g L-1 was approximately 3, 5, and 22 g L-1, respectively. The changes of rate constants were larger when the initial pH decreased from 6.5 to 5.0 than from 8.0 to 6.5 (Table 1). At a pH of 5.0 and 21% O2, a change of temperature from 25 to 15 °C had little effect on oxidation rate of CaSO3 (Figure 2). The first-order rate constants at 25 °C were, in general, slightly lower than those at 15 °C (Table 1). This might be caused by higher solubilities of CaSO3 and O2 in water at lower temperatures (9). However, when the temperature was further decreased from 15 to 5 °C, the oxidation rate of CaSO3 was greatly decreased. There was great variation in reaction rates among experimental replicates when the temperature was 5 °C. Decreasing O2 content from its normal atmospheric level (≈21% O2) to 0.2%, at an initial pH of 5.0 and temperature of 25 °C, also significantly slowed the oxidation rate of CaSO3 (Table 1). At 0.2% O2, the concentration of CaSO3 dropped approximately below 0.5 g L-1 after 46 h as compared to 5 h at 21% O2 (Figure 3). Calcium sulfite greatly decreased the concentration of dissolved O2 in water exposed to atmospheric conditions at 25 °C (Figure 4). This experiment was repeated at several different times, and only the results of one experiment are presented. Even at pH 8.0, under which the solubility of CaSO3 is approaching its lowest level and the rate of sulfite oxidation rate is reduced as compared to when pH is acidic, the concentration of dissolved O2 in a slurry of 10 g L-1 CaSO3 was below 0.01 mg L-1, which was the lowest concentration that our instrument could detect. The reduction of the O2 content to such low levels can very effectively inhibit pyrite oxidation. The combination of low O2 contents and low solubility and oxidation of CaSO3 at pH 8 suggests that mixing alkaline materials with CaSO3 extends the stability of CaSO3 in pyritic environments and thus enhances its ability to control acid formation. At extremely low pH values, the oxidation of CaSO3 released gaseous S (Table 2). The form of this sulfur was SO2, which was identified by the specific color reaction between
TABLE 2. Effects of Initial pH and Initial CaSO3 Concentrations on the Amounts of Gaseous S Release from Aqueous CaSO3 Slurry during 24-h Oxidation at 21% O2, 25 °C initial pH
initial CaSO3 concn (g L-1)
gaseous S releasea (% of initial S)
2.0 3.5 5.0 2.0 3.5 5.0 2.0 3.5 5.0
10 10 10 20 20 20 40 40 40
33.1 14.1 1.40 34.9 21.0 -0.20 43.9 21.7 9.30
a Calculated as (A - B) × 100/A where A and B are the initial and final S concentrations, respectively.
FIGURE 3. Effect of O2 content in air on oxidation of CaSO3 in an aqueous slurry at initial pH 5.0 and 25 °C. Error bar ) 1 SD. The rates follow first-order kinetics with rate constants shown in Table 1.
FIGURE 5. Decreased CaSO3 concentrations in an aqueous slurry at three initial pH values (2.0, 3.5, and 5.0), 21% O2, and 25 °C. Error bar ) 1 SD.
FIGURE 4. Effect of pH on dissolved O2 concentrations in an aqueous slurry of 10 g L-1 CaSO3 at 21% O2 and 25 °C. SO2 and rosaniline reagent. At initial pH values of 2.0 and 3.5, the amounts of SO2 release were significant at all three initial concentrations of CaSO3, and the maximum amount of S loss reached 44% of the total initial S. The amount of SO2 release increased with increasing initial concentration of CaSO3 and with decreasing pH. The relationship between the amount of SO2 release and the initial pH was linear at all three concentrations of CaSO3 (r 2 > 0.97) with about 10% increase of SO2 release per unit decrease in pH. The amounts of SO2 released from CaSO3 slurry were different at the initial pH values of 2.0, 3.5, or 5.0, but the rates of decrease of CaSO3 concentrations in the slurry were not much different between the initial pH values at a given CaSO3 concentration (Figure 5). This indicates that when initial pH is decreased, the dissociation of sulfite to SO2 (as shown below):
H2SO3 f SO2 + H2O is increasingly favored over oxidation of sulfite to sulfate.
The goal of a scrubber reaction at an coal-fired utility is to prevent SO2 release into the air. If in the process of mixing the FGD with an extremely acid coal refuse or spoil material to inhibit pyrite oxidation SO2 is again released to the atmosphere, we have potentially reversed the scrubber reaction. Therefore, it is important that alkaline materials be added into pyritic environments to adjust pH to at least 5.0 or above to reduce SO2 formation. Furthermore, as pH approaches neutral, sulfite can also consume or react with hydrogen ions according to
HSO3- T SO32- + H+ (pKa ) 7.25) Inhibition of Acid Formation in Coal Wash Refuse by CaSO3. Acid leaching from fresh coarse coal refuse with a strong potential to rapidly produce acid was significantly inhibited by CaSO3 (p e 0.0014) and CaSO3-containing FGD (p e 0.0001). During the initial 4 weeks, the leachate acidity of the control columns increased from 160 to 680 mmol of H+ L-1, and the leachate pH dropped from 2.61 to 1.12 (Figure 6). However, the leachate acidity of the CaSO3-treated columns was below 30 mmol of H+ L-1 during the initial 4 weeks, and the leachate pH increased from 2.82 to 3.93. Throughout the 27 weeks of leaching, the CaSO3 treatment always produced lower leachate acidity than the control (at least 130 mmol of H+ L-1 lower for the initial 13 weeks) and VOL. 34, NO. 11, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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ash and the CaSO4‚2H2O may also enhance the persistence of CaSO3 by retarding oxygen diffusion from the atmosphere into the coal refuse. The strong inhibitory effects on acid leaching from the coal refuse caused by CaSO3 and CaSO3-containing FGD began to disappear after 13 weeks. This was possibly due to the strong potential of acid formation in the coal refuse, the leaching procedure applied, and the close contact of sulfite with O2 in the atmosphere. The optimum conditions for using CaSO3 to prevent acid formation would be to bury it with coal mine spoil and coal refuse to avoid direct contact of the sulfite with O2 in the atmosphere. Addition of sufficient amounts of alkaline material, such as CaCO3, together with CaSO3 to adjust pH would be beneficial to optimize the inhibitory effect of the CaSO3 on acid formation by reducing sulfite leaching and SO2 formation.
Acknowledgments The research was funded in part by the Ohio Coal Development Office, Columbus, OH, as Grant CDO/D-931-008 in cooperation with the Dravo Lime Company (Pittsburgh, PA) and the OARDC/The Ohio State University (Wooster, OH).
Literature Cited
FIGURE 6. Change of acidity (mmol of H+) and pH in leachates from coal refuse columns treated with CaSO3 or CaSO3-containing flue gas desulfurization byproduct. reduced the total leachate acidity of the control by 45% (74% for the initial 13 weeks). Calcium sulfite-containing FGD, which contained CaSO3 (64%), CaCO3 (10%), fly ash (17%), and CaSO4‚2H2O (9%), showed much stronger inhibitory effects on acid leaching from the coal refuse columns than did CaSO3 (Figure 6). The leachate acidity of the FGD treatment was near zero from weeks 2 to 10. The total acidity of the FGD treatment was only 1.8% of that of the CaSO3 treatment during the initial 13 weeks. The leachate pH of the FGD treatment reached a maximum of 8.14 in the sixth week and remained above 5.35 for the entire initial 13 weeks. This was at least 2.5 units higher than the leachate pH of the CaSO3 treatment. According to the pH effect on dissolution and oxidation kinetics of CaSO3, discussed in the previous section, CaCO3 might enhance the persistence of CaSO3 in coal refuse by increasing pH. However, even at increased pH levels, the oxidation rate of CaSO3 by O2 can still be fast enough to maintain a low oxygen level. The fine particle sizes of the fly
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(1) Evangelou, V. P. Pyrite Oxidation and its Control; CRC Press: Boca Raton, FL, 1995. (2) Skousen, J. G.; Ziemkiewicz, P. F. Acid Mine Drainage Control and Treatment; Skousen, J. G., Ziemkiewicz, P. F., Eds.; West Virginia University and National Mine Land Reclamation Center: Morgantown, WV, 1995; pp 45-56. (3) Modderman, J. P. J. Assoc. Off. Anal. Chem. 1986, 69, 1-3. (4) Schroeter, L. C. J. Pharm. Sci. 1963, 52, 559-563. (5) Weast, R. C., Ed. CRC Handbook of Chemistry and Physics; The Chemical Rubber Co.: Cleveland, OH, 1969. (6) Pasiuk-Bronikowska, W.; Ziajka, J. Chem. Eng. Sci. 1985, 40, 1567-1572. (7) Pasiuk-Bronikowska, W.; Ziajka, J. Chem. Eng. Sci. 1988, 44, 915-920. (8) Dick, W. A.; Stehouwer, R. C.; Bigham, J. M.; Wolfe, W. E.; Hao, Y.-L.; Adriano, D.; Beeghly J.; Haefner R. J. In Beneficial Uses of Land Applied Agricultural, Industrial and Municipal ByProducts; Power, J. F., Dick, W. A., Eds.; Soil Science Society of America: Madison, WI, 2000; Chapter 18. (9) Masson, M. R., Ed. Solubility Data Series, 1st ed.; Pergamon Press: Oxford, England, 1986; Vol. 26, pp 187-195. (10) Tseng, P. C.; Rochelle, G. T. Environ. Proc. 1986, 5, 34-40. (11) Mosier, A. R.; Mack, L. Soil Sci. Soc. Am. J. 1980, 44, 1121-1123. (12) Urone, P. F.; Boggs, W. E. Anal. Chem. 1951, 23, 1517-1519. (13) Stehouwer, R.; Dick, A.; Bigham, J.; Forster, R.; Hitzhusen, F.; McCly, E.; Traina, S.; Wolfe, W.; Haefner, R. Report TR-105264; American Electric Power Company: Palo Alto, CA, 1995. (14) Begheijn, L. Th.; van Breemen, N.; Velthorst, E. J. Commun. Soil Sci. Plant Anal. 1978, 9, 873-882. (15) Hao, Y. Ph.D. Dissertation, The Ohio State University, Columbus, OH, 1998.
Received for review April 14, 1999. Revised manuscript received March 13, 2000. Accepted March 13, 2000. ES9904235
