Magnesium is a valuable, lightweight metal. It is used as a structural metal
and in alloys, in batteries, and in chemical synthesis. Although magnesium is
plentiful in Earth's crust, it is cheaper to "mine" the metal from seawater.
Magnesium forms the second most abundant cation in the sea (after sodium);
there are about \(1.3 \mathrm{~g}\) of magnesium in \(1 \mathrm{~kg}\) of
seawater. The method of obtaining magnesium from seawater employs all three
types of reactions discussed in this chapter: precipitation, acid-base, and
redox reactions. In the first stage in the recovery of magnesium, limestone
\(\left(\mathrm{CaCO}_{3}\right)\) is heated at high temperatures to produce
quicklime, or calcium oxide \((\mathrm{CaO})\) :
$$
\mathrm{CaCO}_{3}(s) \longrightarrow \mathrm{CaO}(s)+\mathrm{CO}_{2}(g)
$$
When calcium oxide is treated with seawater, it forms calcium hydroxide
\(\left[\mathrm{Ca}(\mathrm{OH})_{2}\right]\), which is slightly soluble and
ionizes to give \(\mathrm{Ca}^{2+}\) and \(\mathrm{OH}^{-}\) ions:
$$
\mathrm{CaO}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow
\mathrm{Ca}^{2+}(a q)+2 \mathrm{OH}^{-}(a q)
$$
The surplus hydroxide ions cause the much less soluble magnesium hydroxide to
precipitate:
$$
\mathrm{Mg}^{2+}(a q)+2 \mathrm{OH}^{-}(a q) \longrightarrow
\mathrm{Mg}(\mathrm{OH})_{2}(s)
$$
The solid magnesium hydroxide is filtered and reacted with hydrochloric acid
to form magnesium chloride \(\left(\mathrm{MgCl}_{2}\right)\)
\(\mathrm{Mg}(\mathrm{OH})_{2}(s)+2 \mathrm{HCl}(a q) \longrightarrow\)
$$
\mathrm{MgCl}_{2}(a q)+2 \mathrm{H}_{2} \mathrm{O}(l)
$$
After the water is evaporated, the solid magnesium chloride is melted in a
steel cell. The molten magnesium chloride contains both \(\mathrm{Mg}^{2+}\) and
\(\mathrm{Cl}^{-}\) ions. In a process called electrolysis, an electric current
is passed through the cell to reduce the \(\mathrm{Mg}^{2+}\) ions and oxidize
the \(\mathrm{Cl}^{-}\) ions. The halfreactions are
$$
\begin{aligned}
\mathrm{Mg}^{2+}+2 e^{-} \longrightarrow \mathrm{Mg} \\
2 \mathrm{Cl}^{-} \longrightarrow \mathrm{Cl}_{2}+2 e^{-}
\end{aligned}
$$
The overall reaction is
$$
\mathrm{MgCl}_{2}(l) \longrightarrow \mathrm{Mg}(s)+\mathrm{Cl}_{2}(g)
$$
This is how magnesium metal is produced. The chlorine gas generated can be
converted to hydrochloric acid and recycled through the process.
(a) Identify the precipitation, acid-base, and redox processes.
(b) Instead of calcium oxide, why don't we simply add sodium hydroxide to
precipitate magnesium hydroxide?
(c) Sometimes a mineral called dolomite (a combination of \(\mathrm{CaCO}_{3}\)
and \(\mathrm{MgCO}_{3}\) ) is substituted for limestone
\(\left(\mathrm{CaCO}_{3}\right)\) to bring about the precipitation of magnesium
hydroxide. What is the advantage of using dolomite?
(d) What are the advantages of mining magnesium from the ocean rather than
from Earth's crust?
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