Iron (Fe) is obtained from rock that is extracted from open pit mines and then
crushed. The process used to obtain the pure metal from the crushed rock
produces solid waste, called tailings, which are stored in disposal areas near
the mines. The tailings pose a serious environmental risk because they contain
sulfides, such as pyrite ( \(\mathrm{FeS}_{2}\) ), which oxidize in air to
produce metal ions and \(\mathrm{H}^{+}\) ions that can enter into surface water
or ground water. The oxidation of \(\mathrm{FeS}_{2}\) to \(\mathrm{Fe}^{3+}\) is
described by the unbalanced chemical equation below.
\(\mathrm{FeS}_{2}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{g})+\mathrm{H}_{2}
\mathrm{O}(\mathrm{l}) \longrightarrow\)
\(\quad
\mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{SO}_{4}^{2-}(\mathrm{aq})+\mathrm{H}^{+}(\mathrm{aq})
\quad(\text { not balanced })\)
Thus, the oxidation of pyrite produces \(\mathrm{Fe}^{3+}\) and \(\mathrm{H}^{+}\)
ions that can leach into surface or ground water. The leaching of
\(\mathrm{H}^{+}\) ions causes the water to become very acidic. To prevent
acidification of nearby ground or surface water, limestone
\(\left(\mathrm{CaCO}_{3}\right)\) is added to the tailings to neutralize the
\(\mathrm{H}^{+}\) ions:
\(\mathrm{CaCO}_{3}(\mathrm{s})+2 \mathrm{H}^{+}(\mathrm{aq})
\underset{\mathrm{Ca}^{2+}}{\longrightarrow}(\mathrm{aq})+\mathrm{H}_{2}
\mathrm{O}(\mathrm{l})+\mathrm{CO}_{2}(\mathrm{g})\)
(a) Balance the equation above for the reaction of \(\mathrm{FeS}_{2}\) and
\(\mathrm{O}_{2}\). [ Hint: Start with the half-equations
\(\mathrm{FeS}_{2}(\mathrm{s}) \rightarrow\)
\(\left.\mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{SO}_{4}^{2-}(\mathrm{aq}) \text {
and } \mathrm{O}_{2}(\mathrm{g}) \rightarrow \mathrm{H}_{2} \mathrm{O}(1)
.\right]\) (b) What is the minimum amount of \(\mathrm{CaCO}_{3}(\mathrm{s})\)
required, per kilogram of tailings, to prevent contamination if the tailings
contain \(3 \%\) S by mass? Assume that all the sulfur in the tailings is in the
form \(\mathrm{FeS}_{2}\).
