and water, oxidation of metal sulfides (often pyrite, which is iron-sulfide) within the surrounding rock and overburden generates acidity. Colonies of bacteria and archaea greatly accelerate the decomposition of metal ions, although the reactions also occur in an abiotic environment. These microbes, called extremophiles for their ability to survive in harsh conditions, occur naturally in the rock, but limited water and oxygen supplies usually keep their numbers low. Extremophiles known as acidophiles especially favor the low pH levels of abandoned mines. In particular, Acidithiobacillus ferrooxidans is a key contributor to pyrite oxidation. Metal mines may generate highly acidic discharges where the ore is a sulfide mineral or is associated with pyrite. In these cases the predominant metal ion may not be iron but rather zinc, copper, or nickel. The most commonly mined ore of copper, chalcopyrite, is itself a copper-iron-sulfide and occurs with a range of other sulfides. Thus, copper mines are often major culprits of acid mine drainage. At some mines, acidic drainage is detected within 2–5 years after mining begins, whereas at other mines, it is not detected for several decades. In addition, acidic drainage may be generated for decades or centuries after it is first detected. For this reason, acid mine drainage is considered a serious long-term environmental problem associated with mining. Chemistry The chemistry of oxidation of pyrites, the production of ferrous ions and subsequently ferric ions, is very complex, and this complexity has considerably inhibited the design of effective treatment options. Although a host of chemical processes contribute to acid mine drainage, pyrite oxidation is by far the greatest contributor. A general equation for this process is: 2FeS2{\scriptstyle (s)} + 7O2{\scriptstyle (g)} + 2H2O{\scriptstyle (l)} -> 2Fe^2+ {\scriptstyle (aq)} + 4SO4^{2-}{\scriptstyle (aq)} + 4H+{\scriptstyle (aq)}. The oxidation of the sulfide to sulfate solubilizes the ferrous iron (iron(II)), which is subsequently oxidized to ferric iron (iron(III)): 4Fe^2+{\scriptstyle (aq)} + O2{\scriptstyle (g)} + 4H+{\scriptstyle (aq)} -> 4Fe^3+{\scriptstyle (aq)} + 2H2O{\scriptstyle (l)}. Either of these reactions can occur spontaneously or can be catalyzed by microorganisms that derive energy from the oxidation reaction. The ferric cations produced can also oxidize additional pyrite and reduce into ferrous ions: FeS2{\scriptstyle (s)} + 14Fe^3+{\scriptstyle (aq)} + 8H2O{\scriptstyle (l)} -> 15Fe^2+{\scriptstyle (aq)} + 2SO4^{2-}{\scriptstyle (aq)} + 16H+{\scriptstyle (aq)}. The net effect of these reactions is to release H+, which lowers the pH and maintains the solubility of the ferric ion. Effects Effects on