Electrochemical Principles of Metallurgy

Thermodynamic reduction with carbon or CO isn't the only route to a free metal. Where a metal ion sits in solution or in a molten salt, electrochemistry takes over — reduction happens either through electrolysis or by adding a more reactive reducing element.

The governing relation

Electrochemical feasibility connects directly back to Gibbs energy:

where $n$ is the number of electrons transferred and $E^\circ$ is the standard electrode potential of the redox couple. More reactive metals have large negative electrode potentials, which is precisely why they're harder to reduce. If the difference between two $E^\circ$ values works out positive, $\Delta G^\circ$ in equation 6.40 comes out negative, and the less reactive metal is displaced out of solution while the more reactive metal dissolves into it — for example:

In straightforward electrolysis, $M^{n+}$ ions are discharged and deposited at the cathode (the negative electrode). Because different metals behave differently once liberated, electrode materials are chosen with care, and a flux is sometimes added simply to make the molten mass conduct better.

Aluminium: the Hall–Héroult process

Purified $Al_2O_3$ is mixed with cryolite ($Na_3AlF_6$) or $CaF_2$, which lowers the melting point of the mixture dramatically and improves its conductivity. The fused mass is then electrolysed (Fig. 6.6) in a steel vessel lined with carbon acting as cathode, with graphite anodes dipping into the melt. The overall cell reaction is:

At the electrodes individually:

Oxygen liberated at the anode attacks the carbon anode itself, which is why the graphite anodes are steadily consumed — roughly 0.5 kg of carbon anode is burnt away for every kg of aluminium produced, an operating cost built into the process.

Copper from low-grade ores and scrap: hydrometallurgy

Where an ore is too low-grade for conventional smelting, copper is instead leached out with acid or even bacteria, and the resulting $Cu^{2+}$ solution is reduced with scrap iron or hydrogen gas:

This is the same displacement idea as equation 6.41, just using $H_2$ in place of a more reactive metal.

Extraction by oxidation: chlorine from brine

Not every "extraction" is a reduction — non-metals are often obtained by oxidation instead. Chlorine, abundant in sea water as common salt, is extracted from brine by electrolysis:

$2Cl^-(aq) + 2H_2O(l) \rightarrow 2OH^-(aq) + H_2(g) + Cl_2(g) \qquad(6.47)$ …