Acidity is initially generated from the oxidation of pyrite (FeS2; Section 1.3.1) in disturbed overburden where sulphuric acid (H2SO4) is created. de Joux (2003) conducted acid-generating column tests from coal and overburden collected on the Denniston Plateau. He found that coal and mudstone generated leachate containing the lowest pH at 2.8 and 2.5-2.7, respectively. Sandstone leachate contained pH values between 5.3 and 5.5.
The primary metals associated with AMD within the Brunner Coal Measures are Fe and Al. Iron is released during the decomposition of pyrite as a result of the acid production and drop in pH. Aluminium leaches from the ubiquitous micaceous and feldspathic-rich rocks within the carbonaceous mudstones and sandstone (Black et al., 2005). Other metals including Cu, Ni, Zn, Cd, As, Pb and Mn also dissolve from minerals within the parent bedrock and overburden when exposed to acidity. Weber et al. (2006) found that the most likely source of Ni was from pyrite within the Kaiata mudstone.
Numerous factors including site-specific geology, geochemistry and mining practices (e.g. opencast versus underground mining) influence AMD chemistry. Pope et al. (2006) showed variable AMD chemistry generated within the Brunner Coal Measures (pH 2.41-3.78). Pope et al. (2006) also delineated water chemistry from opencast and underground mines and as a function of their geology. There was generally a higher Al:Fe molar concentration ratio in opencast mines (Al/Fe = 1.02-25.91) compared with underground mines (Al/Fe = 0.22-4.06). This was attributed to the greater disturbances of overburden and sediments in open mine pits, which allow greater reaction time of sulphuric acid (H2SO4) produced from pyrite oxidation with Al-containing minerals (Pope et al. 2006; 2010). Newman (1988) and Pope et al. (2006; 2010) also report that sediment surrounding coal seams is often feldspar depleted and contains less Al compared with sediments stratigraphically further from the coal seams. During underground mining, most sediment disturbance occurs near the coal seam(s), whereas overburden extracted during open-pit mining includes sediment stratigraphically further from the coal seam(s).
Sediment, typically quantified in the aqueous matrix as total suspended solids (TSS) and/or turbidity, is also a common contaminant associated with AMD. Sediment transport occurs during erosion of disturbed land, especially during high intensity rainfall events. During the mining process, vegetation which otherwise stabilises soil is removed. Rock fragments are also created during blasting, which are removed to overburden dumps and can be prone to erosion, especially on steep slopes. Minimising erosion and subsequent sediment transport via advection in surface water runoff poses a substantial challenge for the mining industry. Jack (2006) showed that erosion can be reduced and better managed by slope reduction of stockpiles at Stockton Coal Mine, although natural armouring significantly reduced sediment loss over a short-time frame.
1.3.1 Iron Chemistry
Equations 1.1-1.4 summarise the commonly accepted chemical reactions associated with pyrite oxidation and Fe hydrolysis (Skousen, 1996; Rose and Cravotta, 1998; Pennsylvania Department of Environmental Protection, 1999; Ford, 2003; Watzlaf et al., 2004).
2 FeS2 + 7 O2 + 2 H2O → 2 Fe2+ + 4 SO42- + 4 H+ (1.1) pyrite + oxygen + water → ferrous iron + sulphate + hydrogen cations (proton acidity)
4 Fe2+ + O
2 + 4 H+→ 4 Fe3+ + 2 H2O (1.2) ferrous iron + oxygen + hydrogen cations (acidity) → ferric iron + water
4 Fe3+ + 12 H
2O ↔ 4 Fe(OH)3↓ + 12 H+ (1.3) ferric iron + water → ferric hydroxide (ppt) + hydrogen cations (proton acidity)
FeS2 + 14 Fe3+ + 8 H2O → 15 Fe2+ + 2 SO42- + 16 H+ (1.4) pyrite + ferric iron + water → ferrous iron + sulphate + hydrogen cations (proton acidity) Equation 1.1 denotes the weathering of pyrite, which is oxidised once exposed to oxygen and water. This results in the release of proton acidity in the form of H+ ions. Two moles of H+ ions are released for every mole of pyrite oxidised during this reaction.
Equation 1.2 represents the rate-determining step where ferrous iron (Fe2+) is oxidized to ferric iron (Fe3+) in the presence of oxygen (Singer and Stumm, 1970; Lowson, 1982). One mole of proton acidity is consumed for every mole of Fe2+ oxidised. The reaction rate is pH dependent and occurs more readily at pH≥3.5 where the reaction can readily proceed abiotically. At lower pH values (commonly <4.0), iron-oxidizing bacteria such as Acidithiobacillus ferrooxidans (formerly called Thiobacillus ferrooxidans) contribute significantly to the oxidative process and can increase reaction rates by a factor up to 1E6 (Singer and Stumm, 1970; Kleinmann and Crerar, 1979; Kirby et al., 1999; Watzlaf et al., 2004).The Fe3+ formed as a result of Fe2+ oxidation can either react to form ferric hydroxide (Equation 1.3) or act as a strong oxidising agent causing further dissolution of pyrite and release of proton acidity (Equation 1.4).
Iron hydrolysis occurs in Equation 1.3. By-products of this reaction include one mole of ferric hydroxide precipitate and three moles of proton acidity for every mole of Fe3+ hydrolysed. Not all Fe3+ formed in Equation 1.2 precipitates as ferric hydroxide in Equation 1.3, especially at pH<4. Under these conditions, Equation 1.3 is most accurately written as an equilibrium reaction proceeding both forward and reverse. These reactions typically occur relatively rapidly in mine workings due to low pH so solubility equilibrium is often achieved.
The presence of ferric hydroxide was observed in AMD-impacted streams in the vicinity of the Brunner Coal Measures at pH values as low as 2.5. The presence of acid-tolerant algae at these locations may potentially contribute to ferric hydroxide formation at such low pHs. Microenvironments are possibly created at and near the algae-water interface where algae supersaturates the water with dissolved oxygen (DO) and some alkalinity may also be produced during algae photosynthesis via consumption of carbon dioxide (CO2) from carbonic acid (H2CO3; Kadlec and Knight, 1996; Kadlec and Wallace, 2009). The mechanisms of the physiological operation of algae at such low pHs are uncertain but it is suspected that they are somehow able to process free carbon dioxide (Harding and Boothryd, 2004). The self-catalytic nature of biologically mediated ferric hydroxide precipitation also likely contributes to the ongoing formation of ferric hydroxide. Ferric Fe that does not precipitate as ferric hydroxidein Equation 1.3 acts as the oxidising reactant in Equation 1.4. The oxidation of pyrite by Fe3+ occurs very rapidly and is considered the fast step in pyrite dissolution (Lowson, 1982). Ferric Fe acts as a stronger oxidising agent than DO (Younger et al., 2002). As a result, pyrite oxidation by Fe3+ shown in Equation 1.4 proceeds quicker than pyrite oxidation via DO shown in Equation 1.1. For every mole of pyrite oxidized by Fe3+, 16 moles of proton acidity are released.
Equations 1.1, 1.2 and 1.4 are cyclic, self propagating and will proceed until pyrite, Fe3+ or other oxidizing metals or oxygen are used up. An overall reaction summary of the pyrite oxidation process is shown in Equation 1.5.
4 FeS2 + 15 O2 + 14 H2O → 4 Fe(OH)3↓ + 8 SO42- + 16 H+ (1.5) pyrite + oxygen + water → ferric hydroxide (ppt) + sulphate + hydrogen cations (proton acidity)
Overall, four moles of H+ cations are generated for every mole of pyrite oxidised in the net sequence of reactions.
1.3.2 Aluminium Chemistry
The primary sources of Al at Stockton Coal Mine include potassium feldspar or microcline (KAlSi3O8), muscovite (KAl2(AlSi3O10)(F,OH)2) and kaolinite (Al2Si2O5(OH)4) (Black et al., 2005). When potassium feldspar is exposed to proton acidity, silicic acid (H4SiO4) and kaolinite are produced as shown in Equation 1.6 (Younger et al., 2002; Watzlaf et al., 2004). Kaolinite is further degraded in the presence of proton acidity, resulting in trivalent Al ions as shown in Equation 1.7. Precipitation of Al3+ as aluminium hydroxide (AlOH
3), which typically occurs at a pH of 4.7 or greater, releases proton Lewis acidity in the form of hydrogen ions as shown in Equation 1.8.
2 KAlSi3O8 + 2H+ + 9 H2O → 4 H4SiO4 + Al2Si2O5(OH4) + 2 K+ (1.6) potassium feldspar + hydrogen ions (proton acidity) + water → silicic acid + kaolinite Al2Si2O5(OH4) + 6 H+→ 2 Al3+ + 2 H4SiO4 + H2O (1.7)
kaolinite + hydrogen ions (proton acidity) → aluminium ions + silicic acid + water Al3+ + 3 H
2O → Al(OH)3↓ + 3 H+ (1.8) aluminium ions + water → aluminium hydroxide (ppt) + hydrogen ions (proton acidity) Iron and Al dissolution are typically more complicated than depicted in the reactions presented in Equations 1.1-1.8. Additional elements such as Ca, Mg, Na, K and Si and lattice structure influence mineral dissolution and mechanisms for releasing and neutralising proton acidity (Weber, 2003).
1.3.3 Acidity
Acidity is defined as the ability of a water to neutralise a base (Watzlaf et al., 2004; American Public Health Association (APHA), 2005). It is most often reported in units mg/L as calcium carbonate (CaCO3) since the formation of bicarbonate (HCO3-) from calcium carbonate dissolution (Section 1.5.1.1) represents the most common mechanism for creating alkalinity and neutralising acidity in nature (Tchobanoglous and Schroeder, 1985). Acidity can be measured empirically via titration with a sodium hydroxide (NaOH) solution following Standard Method 2310B (APHA, 2005) to a pH endpoint of 3.7, which represents acidity (pH 3.7), or to a pH endpoint of 8.3, which represents total acidity (pH 8.3). Acidity can also be calculated by summing metal and proton acidity as shown in Equation 1.9 where CFe2+, CFe3+, CAl, CCu, CNi and CZn represent their respective dissolved metal concentrations in mg/L (modified from Hedin et al., 1994a; Younger et al., 2002; PIRAMID Consortium, 2003; Watzlaf et al., 2004).
Aciditycalc (mg/L as CaCO3) = 50.045(2 CFe2+/55.85 + 3 CFe3+/55.85+ 3 CAl /26.98 (1.9) + 2 CCu /63.55 + 2 CNi /58.71 + 2 CZn /65.38 + 1000(10-pH))
Equation 1.9 can also be modified to include acidity contributions from additional metals that may be present in MIWs such as Mn and As. Cravotta and Kirby (2004) concluded that measured acidity can be used to avoid calculation discrepancies related to the speciation of metals in the dissolved or precipitated state and the ionic state of Fe as Fe3+ and Fe2+. Cravotta and Kirby (2004) also reported that at pH>2.2, Fe3+ complexes with hydroxide anions (OH-), thus, it may be appropriate to modify Equation 1.9 further such that all Fe in a mine water is represented as Fe2+, avoiding the reliance on speciating between the two Fe cationic states. Cravotta and Kirby (2004) and Hedin (2006b) found that calculated acidity (Equation 1.9) and total acidity (pH 8.3) values were comparable and could be used as a reliable quality assurance/quality control (QA/QC) check; however, Means and Hilton (2004) found poor association between calculated acidity and total acidity (pH 8.3) for three of four
MIWs evaluated. Magnesium hydrolysis during analysis for two of the samples contributed to higher acidity measured from titration compared with calculated values. Incomplete hydrolysis of Mn during titration contributed to biased low total acidity (pH 8.3) compared to calculated acidity for one of the samples. This demonstrates how site specific mine-water chemistry can influence the relationship between calculated and measured acidities.