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TECHNICAL REPORTS

Neutralization Potential Determination of Siderite (FeCO₃) Using Selected Oxidants

^a Railroad Commission of Texas, Surface Mining and Reclamation Division, P.O. Box 12967 Austin, TX 78711-2967
^b USDA-ARS, 808 East Blackland Road, Temple, TX 76502
^c Department of Soil and Crop Sciences, MS 2474, Texas A&M University, College Station, TX 77843-2474

^* Corresponding author (gn-white{at}tamu.edu)

Received for publication May 12, 2005.

	ABSTRACT

TOP ABSTRACT INTRODUCTION MATERIALS AND METHODS RESULTS AND DISCUSSION CONCLUSIONS REFERENCES

 
Siderite (FeCO₃) is commonly found in coal overburden and, when present, can cause interference in the determination of neutralization potential (NP). Under acidic testing conditions, FeCO₃ reacts to neutralize acid, which contributes to the NP. However, continued weathering of FeCO₃ (oxidation of Fe²⁺ and hydrolysis of Fe³⁺) produces a neutral to slightly acidic solution. The effects of hydrogen peroxide (H₂O₂), potassium permanganate (KMnO₄), and O₂ on the laboratory measurement of NP of siderite samples taken from overburden were examined. All oxidation treatments lowered the NP values of the siderite samples as compared with the standard USEPA method. However, oxidation with H₂O₂ produced variable results depending on the amount of H₂O₂ added. Neutralization potential values obtained after oxidation treatments were highly correlated with Mn concentration. Reaction products (i.e., 2-line ferrihydrite) of siderite samples with H₂O₂ and KMnO₄ were not representative of natural siderite weathering. Oxidation with O₂ produced the lowest NP values for siderite samples. The reaction products produced by oxidation with O₂ most closely represent those intermediate products formed when siderite is exposed to atmospheric weathering conditions. Oxidation with O₂ also proved to be the most reproducible method for accurately assessing NP when siderite is present in overburden samples.

Abbreviations: AAS, atomic adsorption spectroscopy • ABA, acid–base account • AFM, acid-forming materials • EA, exchangeable acidity • GR1(CO₃), green rust • NP, neutralization potential • PA, potential acidity • XRD, X-ray diffraction analysis

	INTRODUCTION

TOP ABSTRACT INTRODUCTION MATERIALS AND METHODS RESULTS AND DISCUSSION CONCLUSIONS REFERENCES

 
PRESENT reclamation procedures in Texas allow for the use of mixed overburden as a topsoil substitute. Selectively handled overburden may be used to replace topsoil if it is of equal or better quality than the original topsoil and is the best available material to support revegetation. Determination of a suitable topsoil substitute is based on chemical and physical analyses of the overburden and undisturbed topsoil. The analyses include, but are not limited to, pH, net acidity, net alkalinity, nutrient status, particle size analysis, and texture class (Railroad Commission of Texas, 1998). Care must be taken to prevent acid soils and/or acid drainage due to the oxidation of acid-forming materials (AFM) that are commonly found in reduced overburden and in association with lignite seams. To assess residual soil acidity and the potential of acid drainage, the acid–base account (ABA) for overburden samples is commonly determined using Eq. [1] or Eq. [2] where NP is the neutralization potential, PA is the potential acidity, and EA is the exchangeable acidity.

[1]

[2]

When the ABA is positive, one assumes there is an adequate amount of NP to neutralize any acidity present in the mine-soil. If the ABA in negative, the overburden may produce acid mine soils and/or acid drainage and should be buried more than 1.2 m (4 feet) below the surface.

Neutralization potential measures easily weatherable minerals (i.e., olivine and micas), including carbonates, present in an overburden (Doolittle et al., 1992). The NP of an overburden is determined using USEPA Method 600/2-78-054 3.2.3 (Sobek et al., 1978). The sample is first given a fizz rating (none, slight, moderate, or severe) based on visual and audible ratings after treating the sample with 1:3 HCl. Samples are then treated with a known amount of HCl, determined by the fizz rating, and heated until the reaction is complete (sample settles on bottom of flask). The excess acid is titrated with NaOH to a pH of 7.0. Neutralization potential is reported in terms of calcium carbonate (CaCO₃) equivalent and is calculated directly from the amount of acid consumed by the sample.

Siderite (FeCO₃) is frequently present in anoxic overburden and can cause false positive or unrealistically high NP values. Siderite is susceptible to oxidation when exposed to atmospheric weathering conditions, which can be described by Eq. [3] to [6] (Frisbee and Hossner, 1995; Lindsay, 1979).

[3]

[4]

[5]

[6]

Continued weathering of FeCO₃ produces a neutral reaction when lepidocrocite (-FeOOH) or goethite (-FeOOH) is the reaction product. Atmospheric weathering conditions favor alteration of siderite to goethite through dissolution and ensuing oxidation or precipitation (Senkayi et al., 1986). However, under acidic testing conditions, FeCO₃ reacts to neutralize acid, which contributes to the NP (Eq. [3]). Siderite affects NP because Fe²⁺ released during acid digestion does not undergo complete oxidation and hydrolysis (Eq. [4] and [5]) during NP evaluation (Doolittle et al., 1992; O'Shay et al., 1990).

Alterations to the USEPA method are needed to accurately assess the NP of overburdens containing siderite. Skousen et al. (1997) studied the effects of H₂O₂ on NP values for overburden soils from Pennsylvania and West Virginia. Hydrogen peroxide is an effective oxidizing agent at pH values below 5.8 as shown in Eq. [6] (Diehl, 1970).

[7]

Hydrogen peroxide serves to oxidize Fe²⁺ after samples are digested with HCl during NP testing. Hydrogen peroxide theoretically acts to oxidize Fe²⁺ to Fe³⁺, which is subsequently hydrolyzed to an iron hydroxide (Fe(OH)₃) as shown in Eq. [8–11] (Lindsay, 1979; Schwertmann and Cornell, 1991).

[8]

[9]

[10]

[11]

The overall reaction (Eq. [11]) shows that, with Fe(OH)₃ as the reaction product, 1 mol of FeCO₃ consumes 1 mole of H⁺ on dissolution, oxidation, and hydrolysis versus 2 mol H⁺ consumed by nonoxidized samples. Therefore, NP values of oxidized samples should decrease accordingly. Skousen et al. (1997) found that H₂O₂ decreased NP values of samples containing siderite due to complete oxidation of Fe²⁺ and hydrolysis of Fe³⁺.

Analytical laboratories commonly use H₂O₂ for determining NP when siderite is assumed to be present in overburden samples. First, samples are analyzed according to the USEPA method. After titration to pH 7.0, samples are treated with 5 mL of 30% H₂O₂, and then titrated again to pH 7.0. Neutralization potential values are calculated from the total amount of NaOH needed during the first and second titrations.

Potassium permanganate is a strong oxidizing agent (Eq. [12]) that serves as its own color indicator during oxidation–reduction titrations (Diehl, 1970).

[12]

The MnO₄^– ion is intensely purple, and the Mn²⁺ ion is colorless. As long as some reducing agent is being titrated, the solution is colorless, since the purple MnO₄^– ion being added is converted to Mn²⁺. However, when the reducing agent has been consumed, the next drop of titrant will turn the solution light purple or pink. Therefore, the end point occurs one drop beyond the stoichiometric point when all of the reducing agent has been consumed (Zumdahl, 1989).

The expected reactions of KMnO₄ with FeCO₃ are as follows:

[13]

[14]

[15]

Equation [15] indicates that 0.6 mol of H⁺ are consumed for every mole of FeCO₃ oxidized and hydrolyzed. Oxidation is expected to decrease NP values by reducing the amount of acidity consumed by FeCO₃ from 2 H⁺ to 0.6 H⁺ mol^–1.

Oxygen is an effective natural oxidant. The standard reduction potential of O₂ is shown in Eq. [8] (Zumdahl, 1989).

[16]

Pure O₂ can be bubbled through the sample solution to provide the O₂ for Fe²⁺ oxidation. The reactions for the dissolution and oxidation of FeCO₃ under NP testing conditions using O₂ as an oxidant should be identical to natural weathering conditions (Eq. [3]–[6]).

The objective of this research was to examine the effects of the oxidants; H₂O₂, KMnO₄, and O₂ on the reaction of siderite during NP determination and to study the resulting NP values. These experiments were designed to evaluate potential methods for NP determination in overburden samples containing siderite.

	MATERIALS AND METHODS

TOP ABSTRACT INTRODUCTION MATERIALS AND METHODS RESULTS AND DISCUSSION CONCLUSIONS REFERENCES

 
Sample Analysis
Two siderite samples (A and B) were collected from lignite mine overburdens, crushed, and ground in an agate mortar to pass a 250-µm (60-mesh) sieve. The mineral composition of the siderite samples was determined using X-ray diffraction analysis (XRD). X-ray diffraction analysis was also performed on reagent-grade MnCO₃ and FeCO₃ reference samples. Minerals were identified from XRD patterns. Inorganic C was determined using the simple titrimetric method by Bundy and Bremner (1972). The method was altered to allow samples to digest in 2.0 M HCl for 2 d to ensure complete dissolution of minerals. Iron, Ca, Mn, and Mg concentrations of subsamples were determined using atomic adsorption spectroscopy (AAS) after samples were digested with 2.0 M HCl and filtered. The cation concentrations were used to calculate theoretical NP values using stoichiometric relationships for the dissolution of FeCO₃, CaCO₃, MgCO₃, and MnCO₃ under NP testing conditions. For example, Eq. [8] to [11] as described above were utilized to calculate the theoretical NP of 2 g of FeCO₃ after acid dissolution and oxidation with H₂O₂. The theoretical NP value was directly calculated from the theoretical amount of NaOH needed to neutralize the amount of acid produced during the reaction of FeCO₃ with HCl and H₂O₂.

Neutralization Potential Determination
Subsamples of the siderite samples were weighed into flasks to contain 0.05, 0.10, and 0.20 g of siderite sample (equal to 2.5, 5.0, and 10.0% of a 2-g sample). Neutralization potential was determined on siderite subsamples and reagent-grade MnCO₃, FeCO₃, and CaCO₃ using USEPA Method 600/2-78-054 3.2.3 (Sobek et al., 1978).

All samples were digested with 20 mL of 0.1 M HCl (because of the "none" fizz rating) and heated until reaction was complete (sample residue settled on bottom of flask). Samples were brought to 125 mL volume with deionized water, then boiled for 1 min. Samples were allowed to cool to room temperature (25°C), then titrated to pH 7.0 with 0.1 M NaOH using a pH meter and combination electrode. Neutralization potential was reported in grams CaCO₃ equivalent kg^–1 substrate and was determined directly from the amount of excess acid titrated. Solution color (precipitate formed) was determined before and during titration with NaOH using the Munsell Soil Color Charts. Precipitates formed during titration were filtered, dried, and analyzed using XRD spectroscopy.

Neutralization potential was also determined on siderite subsamples and reagent-grade MnCO₃, FeCO₃, and CaCO₃ following digestion in acid and oxidation of Fe²⁺ with H₂O₂, KMnO₄, and O₂. Alterations to the USEPA method are listed in Table 1. Hydrogen peroxide, KMnO₄, or O₂ were incorporated as oxidants and summarized as follows.

Potassium Permanganate
Samples digested in 0.1 M HCl were cooled to room temperature and treated with 1.0 M KMnO₄ (pH 7). Potassium permanganate was slowly added to the samples until a pink color formed, indicating an excess of oxidant and complete oxidation of Fe²⁺. Samples were then titrated with 0.1 M NaOH to pH 7.0. Precipitate color was not determined during titration due to dominance of pink color from KMnO₄.

Oxygen
Samples digested in 0.1 M HCl were treated with pure O₂. The flask containing the digested sample was equipped with a small fish-tank aerator that was attached to a tank of O₂. Gas was vigorously bubbled through the sample solution while titrating to pH 7.0 with 0.1 M NaOH. Oxygen was bubbled through the sample for the duration of the titration. The length of time that the solution was in contact with the O₂ varied depending on the amount of acidity consumed during digestion (i.e., the mineralogy of the sample). Solution color (precipitate formed) was determined before and during titration with NaOH using the Munsell Soil Color Charts. Precipitates formed during titration were filtered, dried, and analyzed using XRD spectroscopy.

Statistical analysis was performed using one-way ANOVA (SPSS, 1997).

	RESULTS AND DISCUSSION

TOP ABSTRACT INTRODUCTION MATERIALS AND METHODS RESULTS AND DISCUSSION CONCLUSIONS REFERENCES

 
Sample Identification
X-ray diffraction results combined with Fe, Mn, Ca, and Mg analysis indicate that Sample A is composed largely of siderite (with Fe partially substituted by Mn, Ca, and Mg) and quartz. Samples A and B have enough Mn substitution to be classified as manganoan siderite and also contain quartz (Fig. 1). Peaks located around 26.6° 2 are characteristic of quartz. Peaks around 32° 2 are indicative of siderite. A peak at approximately 31.5° 2 indicates the presence of manganoan siderite in Samples A and B. The dominant peak that is characteristic of FeCO₃ is shifted to the left of the dominant peak for reagent-grade FeCO₃ toward the MnCO₃ peak, suggesting Mn substitution.

View larger version (23K): [in this window] [in a new window]   Fig. 1. X-ray diffractograms showing major peaks for Samples A and B as compared with reagent-grade rhodochrosite (MnCO3) and siderite (FeCO3). S = siderite, M = manganoan siderite, R = rhodochrosite, and Q = quartz.

View larger version (33K): [in this window] [in a new window]   Fig. 3. Experimental neutralization potential (NP) values of siderite samples determined using USEPA Method 600/2-78-054 3.2.3 (Sobek et al., 1978) and oxidation treatments as compared with carbonate and cation concentrations.

Neutralization potential values of 0.1 g of reagent-grade CaCO₃ were significantly higher than NP values of 0.1 g FeCO₃ and 0.1 g MnCO₃ (Fig. 2). Rhodochrosite (MnCO₃) NP values were 60% of CaCO₃–NP values, while FeCO₃–NP values were only 20% of CaCO₃–NP values.

View larger version (38K): [in this window] [in a new window]   Fig. 2. Neutralization potential values for 0.1 g of reagent-grade CaCO3, MnCO3, and FeCO3. Error bars indicate 1 SD.

View larger version (21K): [in this window] [in a new window]   Fig. 4. Comparison of neutralization potential (NP) values obtained using USEPA Method 600/2-78-054 3.2.3 (Sobek et al., 1978) as compared with oxidation treatments. Error bars indicate 1 SD.

View larger version (29K): [in this window] [in a new window]   Fig. 5. Neutralization potential (NP) values of siderite samples determined using 5 or 10 mL of H2O2 on filtered and nonfiltered samples after digestion with 0.1 M HCl. Error bars indicate 1 SD.

View larger version (35K): [in this window] [in a new window]   Fig. 6. Experimental neutralization potential (NP) values of H2O2 treatments as compared to theoretical values obtained using stoichiometric relationships describing the dissolution of Fe, Ca, Mg, and Mn carbonates under acidic testing conditions.

View larger version (17K): [in this window] [in a new window]   Fig. 7. X-ray diffraction patterns for the dissolution products formed from siderite using O2, KMnO4, and H2O2. The sharp peaks on the O2 pattern correspond to residual siderite.

Oxidation with KMnO₄ decreased NP values of reagent-grade FeCO₃ slightly (Fig. 2). Data in Fig. 2 show that CaCO₃ and MnCO₃ NP values were not significantly affected by oxidation.

Due to the deep purple color resulting from the reaction of KMnO₄ with the precipitates formed during titration, precipitate colors were not analyzed after oxidation and titration. X-ray diffraction patterns show the formation of 2-line ferrihydrite after treatment with KMnO₄. The XRD data also suggests the presence of noncrystalline or poorly ordered Fe-hydroxides (Fig. 7). Hydrolysis of acidic solutions of Fe(III) salts can produce hematite, akaganeite, goethite, and ferrihydrite. Which product forms and its degree of crystallinity is dependent on temperature of reaction, pH, rate of hydrolysis, and the nature of the anions present. Unless the reaction conditions are controlled precisely, a mixture of products will form (Schwertmann and Cornell, 1991). The reaction products of KMnO₄ and FeCO₃ are poorly ordered and therefore not representative of atmospheric oxidation of FeCO₃.

Oxygen
After oxidation with O₂, NP values are no longer relative to Fe concentration, suggesting that FeCO₃ was oxidized and hydrolyzed during testing, eliminating the FeCO₃ interference (Fig. 3D). Remaining NP values were dependent on Mg and Mn concentration. Data in Fig. 4 show that NP values significantly decreased after samples were oxidized with O₂. Neutralization potential values were not significantly different between the H₂O₂ and O₂ treatments for all samples containing <0.2 g (10%) of sample. The NP values obtained after the O₂ treatment were lower than values from the H₂O₂ treatment for samples containing 10% siderite indicating that there may not have been enough oxidant for oxidation and hydrolysis of all the Fe in the sample. The decrease in NP values after oxidation is highly correlated with Fe present in the sample (r² = 0.99).

The NP values of reagent-grade FeCO₃ decreased after oxidation with O₂ as compared with NP values obtained after oxidation with H₂O₂ (Fig. 2), indicating that O₂ had a greater oxidizing effect on the Fe present in the sample than H₂O₂. The NP values of reagent-grade MnCO₃ were less after treatment with H₂O₂ as compared with NP values obtained after treatment with O₂. Data in Fig. 2 show that CaCO₃ NP values were not significantly affected by oxidation with O₂.

During NP testing, green precipitate forms on the initial addition of NaOH, which is converted to a red precipitate on exposure to O₂ indicating the Fe²⁺ in solution is oxidized. Titration with NaOH should proceed slowly so that O₂ has ample time to oxidize Fe²⁺ in the sample and to avoid missing the endpoint at pH 7.0. The color of precipitates formed after complete oxidation with O₂ and titration with NaOH were dark reddish brown (2.5YR 3/4), suggesting that Fe²⁺ was oxidized to Fe³⁺ and hydrolyzed to ferrihydrite, lepidocrocite (-FeOOH), or an amorphous Fe oxide during testing. The transformation of FeCO₃ to -FeOOH or -FeOOH after oxidation with O₂ is consistent with atmospheric weathering of FeCO₃ under natural conditions (Senkayi et al., 1986; Frisbee and Hossner, 1995).

The XRD pattern of precipitates formed after oxidation with O₂ and subsequent titration to pH 7.0 suggests the formation of Fe oxides with a poorly crystalline structure or the formation of mixed reaction products (Fig. 7). Peaks at approximately 32 °2 and 52 °2 denote siderite, which was not completely dissolved in 0.1 M HCl. Incomplete dissolution of siderite is not a problem since dissolution of FeCO₃ and oxidation and hydrolysis of Fe²⁺ should produce a neutral reaction under natural weathering conditions. Neutralization potential analysis conditions, when using O₂ as an oxidant, are similar to laboratory conditions for the synthesis of -FeOOH (Schwertmann and Cornell, 1991). Oxidation and hydrolysis of Fe(II) salts produces lepidocrocite, goethite, magnetite, and ferrihydrite. Goethite and -FeOOH form at pH values between 6 and 7. With rapid oxidation, using H₂O₂, feroxyhyte (-FeOOH) is produced (Schwertmann and Cornell, 1991). Laboratory synthesis of -FeOOH involves dissolving FeCl₂ salt in water, adjusting the pH to 6.7 to 6.9 with NaOH, then adding an air supply to bubble through the solution. The initial gray color formed changes to orange with time. The entire process takes 1 to 3 h. The degree of crystallinity of minerals formed depends on reaction conditions. Mixtures of minerals form, rather than a single product, unless reaction conditions are carefully controlled. Reaction pH must remain constant to obtain a crystalline product. During NP testing, the pH is steadily increased with NaOH and O₂ is bubbled through the solution throughout titration. In contrast to the laboratory conditions for the synthesis of -FeOOH, the titration used in NP determinations is completed in <1 h and this is insufficient time to allow crystalline phases to form.

	CONCLUSIONS

TOP ABSTRACT INTRODUCTION MATERIALS AND METHODS RESULTS AND DISCUSSION CONCLUSIONS REFERENCES

 
Neutralization potential values of samples containing siderite were the highest when USEPA Method 600/2-78-054 3.2.3 was used. Oxidation and hydrolysis of Fe with H₂O₂, KMnO₄, O₂ lowered NP values compared with USEPA Method 600/2-78-054 3.2.3. Oxidation with H₂O₂ (5 mL, filtered) and O₂ produced similar NP values and were not highly correlated with Fe concentration (based on the correlation of determination values). The O₂ method was relatively faster to perform and might be preferred in an analytical laboratory setting. Neutralization potential values were reduced after oxidation with KMnO₄; however, they were still highly correlated with Fe concentration.

The authors agree that the preferred method for NP determination for samples containing siderite is as follows:

1. Weigh 2 g of a soil sample that has been processed and sieved to <60 mesh (250 µm) into a 200-mL beaker.
   
2. Add 20 mL of 0.1 M HCl and heat on a hot plate at 121°C for 15 min or until reaction appears to be complete (sample will settle on the bottom of the flask at this point).
   
3. Add distilled water to a volume of roughly 125 mL into the beaker and cover with a watch glass.
   
4. Boil the contents of the beaker for 1 min.
   
5. After the samples have boiled for 1 min, turn the hot plate off and remove the samples from the hot plate to cool.
   
6. Remove the watch glass and rinse the liquid on watch glass into the beakers with the samples. Analysis should not continue until the samples have returned to room temperature.
   
7. After the samples have cooled to room temperature, place the 200-mL beaker containing the sample on the stir plate and equip with stir bar.
   
8. Begin stirring sample while adding NaOH to the burette. Take a calibrated reading of the burette.
   
9. Rinse pH probe on a calibrated meter with deionized water in a separate cup, then place the probe in the sample.
   
10. Note the pH of the sample before beginning titration.a. If the pH is greater than 6.0, the sample must be reweighed and digested using 40 mL of HCl.b. If the pH is approximately 2.0 to 3.0, the sample is ready to be titrated.
    
11. Equip the flask containing the digested sample with a small fish-tank aerator attached to a tank of O₂ and turn on tank slowly so as to gently bubble O₂ though the sample.
    
12. Begin to slowly titrate the sample with NaOH while it is being aerated and stirred. Bubble O₂ through the sample for the duration of the titration. The length of time that the solution is in contact with the O₂ will vary depending on the amount of acidity consumed during digestion (i.e., the amount of time needed to neutralize remaining acidity with NaOH).
    
13. When the sample maintains a pH between 6.95 and 7.05 for 30 s, take a reading of the burette and record the result.
    

Neutralization potential is calculated using Eq. [17]:

[17]

After all oxidation treatments, NP values were highly correlated with Mn content. This is of concern since hydroxyl ions may be consumed during the oxidation and hydrolysis of Mn released during weathering of Mn carbonates, which will not result in the consumption of acidity (Dixon et al., 1982). Therefore, false positive NP values may result when MnCO₃ or other carbonates substituted with Mn are present in overburden. Further investigation on the effects of acid digestion and oxidation of Mn, as well as other minerals, during NP testing is warranted.

	REFERENCES

TOP ABSTRACT INTRODUCTION MATERIALS AND METHODS RESULTS AND DISCUSSION CONCLUSIONS REFERENCES

 

• Bundy, L.G., and J.M. Bremner. 1972. A simple titrimetric method for determination of inorganic carbon in soils. Soil Sci. Soc. Am. Proc. 36:273–275.
• Diehl, H. 1970. Quantitative analysis. Oakland Street Science Press, Ames, IA.
• Dixon, J.B., L.R. Hossner, A.L. Senkayi, and K. Egashira. 1982. Mineralogical properties of lignite overburden as they relate to mine spoil reclamation. p. 169–191. In Acid sulfate weathering. SSSA, Madison, WI.
• Doolittle, J.J., N.M. Frisbee, and L.R. Hossner. 1992. Evaluation of acid–base accounting techniques used in surface-mine reclamation. In Proc. 9th Meeting of the American Society for Surface Mining and Reclamation, Duluth, MN. 14–18 June 1992. Am. Soc. for Surface Mining and Reclamation, Lexington, KY.
• Frisbee, N.M., and L.R. Hossner. 1995. Siderite weathering in acidic solutions under carbon dioxide, air, and oxygen. J. Environ. Qual. 24:856–860.
• Legrand, L., S. Savoye, A. Charusse, and R. Messina. 2000. Study of oxidation products formed on iron in solutions containing bicarbonate/carbonate. Electrochim. Acta 46:111–117.
• Lindsay, W.L. 1979. Chemical equilibria in soils. John Wiley & Sons, New York.
• O'Shay, T., L.R. Hossner, and J.B. Dixon. 1990. A modified hydrogen peroxide oxidation method for determination of potential acidity in pyritic overburden. J. Environ. Qual. 19:778–782.[ISI]
• Railroad Commission of Texas. 1998. Coal mining regulations. Railroad Commission of Texas, Surface Mining and Reclamation Division, Austin, TX.
• Schwertmann, U., and R.M. Cornell. 1991. Iron oxides in the laboratory: Preparation and characterization. Cambridge, New York.
• Schwertmann, U., H. Fechter, R.M. Taylor, and H. Stanjek. 1995. A lecture and demonstration for student on iron oxide formation. p. 11–14. In G.J. Churchman et al. (ed.) Clays controlling the environment. Proc. 10th Int. Clay Conf., Adelaide, Australia. 18–23 July 1993. CSIRO Publ., Melbourne, Australia.
• Senkayi, A.L., J.B. Dixon, and L.R. Hossner. 1986. Todorokite, goethite, and hematite: Alteration products of siderite in east Texas lignite overburden. Soil Sci. 142:36–42.
• Skousen, J., J. Renton, H. Brown, P. Evans, B. Leavitt, K. Brady, L. Cohen, and P. Ziemkiewicz. 1997. Neutralization potential of overburden samples containing siderite. J. Environ. Qual. 26:673–681.
• Sobek, A.A., W.A. Schuller, J.R. Freeman, and R.M. Smith. 1978. Field and laboratory methods applicable to overburdens and minesoils. Industrial environmental research laboratory. USEPA Rep. 600/2-78-054. Office of Research and Dev., USEPA, Cincinnati, OH.
• SPSS. 1997. SigmaStat for Windows Version 2.03. SPSS, Chicago, IL.
• Zumdahl, S.S. 1989. Chemistry. 2nd ed. p. 164–165. D.C. Heath and Co., Lexington, MA.

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