14.4.3: Acid mine drainage

Acid mine drainage forms when water percolates through mining-impacted rock that contains sulfide minerals exposed to oxygen. Oxidation of those minerals produces hydrogen ions, which can drive the pH of the solution to acidic levels. In addition to acid, the water can also contain elevated concentrations of sulfate and numerous metals and metalloids, including arsenic, cadmium, cobalt, chromium, mercury, manganese, molybdenum, nickel, lead, and zinc (Newsome and Falagán, 2021; Nordstrom, 2011), creating a threat to the quality of nearby surface water and groundwater. Most acid mine drainage has pH between 2 and 6 (Nordstrom et al., 2015), although much lower values have been observed. In one of the most extreme cases, pH values as low as -2.5 have been measured in water with elevated concentrations of sulfate \((760,000 \ \text{mg/L})\), iron \((111,000 \ \text{mg/L})\), and several trace elements at the Richmond Mine at Iron Mountain, California, USA (Nordstrom, 2011; Nordstrom et al., 2000). Acid mine drainage exists within most countries worldwide and its source and toxicity has been recognized for thousands of years (Nordstrom, 2011).

Oxidation is thought to initiate with direct reaction between sulfide minerals and oxygen, which can be described with the following reaction in which sulfide minerals are represented by pyrite \(\left(\text{FeS}_{2}\right)\) (Nordstrom et al., 2015): \[\text{FeS}_{2} + 3.5 \ \text{O}_{2} \ (aq) + \text{H}_{2} \text{O} \longleftrightarrow 2 \ \text{SO}_{4}^{2-} + \text{Fe}^{2+} + 2 \ \text{H}^{+}\]

The ferrous iron \((\text{Fe(II)})\) produced by the reaction may subsequently oxidize to ferric iron \((\text{Fe(III)})\): \[\text{Fe}^{2+} + \text{H}^{+} + 0.25 \ \text{O}_{2} \ (aq) \longleftrightarrow \text{Fe}^{3+} + 0.5 \ \text{H}_{2} \text{O}\]

In a solution with near-neutral pH, ferrous iron oxidation can occur rapidly without the aid of microbial catalysis. The ferric iron produced can then form mixed (oxyhydr)oxide phases with variable stoichiometries including ferrihydrite \(\left(\text{Fe}_{5} \text{HO}_{8} \cdot 4\left(\text{H}_{2} \text{O}\right)\right)\), goethite \((\text{FeOOH})\), schwertmannite \(\left(\text{Fe}_{16} \text{O}_{16} (\text{OH})_{y} \left(\text{SO}_{4}\right)_{z} \cdot n \left(\text{H}_{2} \text{O}\right)\right)\), and jarosite \(\left(\text{KFe}_{3} \left(\text{SO}_{4}\right)_{2} (\text{OH})_{6}\right)\) (Nordstrom and Southam, 1997). Such precipitates typically coat the beds of streams receiving mining impacted water, providing a vivid signal of the presence of acid mine drainage (Fig. \(14.7\)).

Figure \(14.7\): Acid mine drainage flowing away from the collapsed entrance to the Davis Mine pyrite mine near Rowe, Massachusetts, USA. The mine was active from 1882 to 1911, after which the mine filled with water and began generating acidic drainage (Becerra et al., 2009). Near the mine entrance, drainage pH is near 2, and outside of the waste rock area, pH ranges from about 4.5 to 6 (Becerra et al., 2009).
Image source: https://commons.wikimedia.org/wiki/File:AMD_at_the_Davis_Mine.jpg

Combining reaction \(\PageIndex{1}\) with reactions for iron oxidation (reaction \(\PageIndex{2}\) and goethite precipitation gives the following overall net reaction (Nordstrom et al., 2015): \[\text{FeS}_{2} + 3.75 \ \text{O}_{2} + 2.5 \ \text{H}_{2}\text{O} \longleftrightarrow 2 \ \text{SO}_{4}^{2-} + \text{FeOOH} + 4 \ \text{H}^{+}\]

Thus, the overall reaction generates four moles of hydrogen ions per mole of pyrite, which works to drive the pH of the environment downward to acidic levels.

Under acidic conditions, the pathway of sulfide mineral oxidation changes. Ferric iron solubility increases and dissolved ferric iron ions become more effective than oxygen as an oxidant for pyrite (Moses et al., 1987). In fact, in systems with pH below 3, ferric iron is the only important oxidizer of pyrite (Konhauser, 2007). An example net reaction can be written as follows: \[\text{FeS}_{2} + 14 \ \text{Fe}^{3+} + 8 \ \text{H}_{2}\text{O} \longleftrightarrow 2 \ \text{SO}_{4}^{2-} + 15 \ \text{Fe}^{2+} + 16 \ \text{H}^{+}\]

In the reaction, ferric iron is reduced to ferrous iron. Thus, for the reaction to continue moving forward, ferric iron needs to be resupplied by ferrous iron oxidation (reaction \(\PageIndex{2}\)). As such, oxygen remains an important ingredient because of its role in ferrous iron oxidation. Alternatively, ferrous iron oxidation could also be coupled with nitrate reduction if nitrate is available (Section 5.3.2).

As the reaction pathways shift with conditions, so do the contributions of microorganisms. Under near-neutral pH conditions, iron-oxidizing microorganisms may be involved (Percak-Dennett et al., 2017) although they are not thought to play a critical role because ferrous iron can rapidly oxidize without catalysis (Nordstrom et al., 2015). However, at acidic pH, ferrous iron oxidation is slow unless microbially catalyzed. Under these conditions, microorganisms can speed up the reaction by five orders of magnitude (Nordstrom et al., 2015). In doing so, they can resupply ferric iron and push pyrite oxidation forward (Fig. \(14.8\)). Without this push, pyrite oxidation would stop because ferrous iron oxidation is too slow at acidic pH (Nordstrom, 2011). Thus, iron-oxidizing microorganisms not only speed up pyrite oxidation but also help sustain it under acidic conditions.

Figure \(14.8\): Schematic illustration pyrite weathering under acidic conditions (e.g., below pH 4.5) by iron- and sulfur-oxidizing microorganisms. Elemental sulfur (𝑆𝑆0) represents one of multiple intermediate sulfur species that can be produced by sulfide oxidation. In addition to oxygen \(\left(\text{O}_{2}\right)\), iron-oxidizing microorganisms can also couple ferrous iron oxidation to nitrate reduction. This figure is based on figure 3 of Newsome and Falagán (2021).
Image source: https://commons.wikimedia.org/wiki/File:AMD_conceptual_model.jpg

In addition to iron oxidizers, sulfide oxidizers are also influential. In fact, pyrite oxidation has been found to be faster in experiments that contain iron- and sulfur-oxidizing microorganisms rather than just iron oxidizers (McGuire et al., 2001; Nordstrom and Southam, 1997). Intermediate reactions during sulfide mineral oxidation produce multiple sulfur species, including elemental sulfur \(\left(\text{S}^{0}\right)\), thiosulfate \(\left(\text{S}_{2} \text{O}_{3}\right)\), polythionates \(\left(\text{S}_{x} \text{O}_{6}^{2-}\right)\), and more. Sulfur oxidizers consume these intermediates and prevent their accumulation in solution and on mineral surfaces (McGuire et al., 2001) (Fig. \(14.8\)). Moreover, reactions catalyzed by sulfur oxidizers generate acid that benefits the acidophilic microorganisms that are responsible for driving sulfide mineral weathering forward (Banfield and Welch, 2000).

Although some microorganism help drive production of acid mine drainage, others can help clean up the contamination (Becerra et al., 2009; Coggon et al., 2012). Acid mine drainage is ultimately caused by oxidative weathering of sulfide minerals. Thus, the negative impacts of acidic drainage can be reversed by re-forming those minerals under conditions where they are not exposed to oxidation. Microorganisms can cause sulfide minerals to form where they catalyze reduction of ferric iron and sulfate (Section 12.1.4). The reactions consume hydrogen ions and thus help neutralize acid (Section 14.2.1). Moreover, the sulfide minerals that form sequester iron and sulfur, as well as hazardous trace elements (Fortin and Beveridge, 1997). Thus, microbial production of sulfide biominerals can lessen impacts of acid mine drainage in much the same way that it can lower arsenic concentrations in naturally contaminated aquifers (Section 14.4.2).