Electrolysis. An electrolytic cell is the reverse of a galvanic cell: electrical energy from an external source drives a non-spontaneous redox reaction. The electrode connected to the positive terminal is the anode (oxidation), and the one connected to the negative terminal is the cathode (reduction). For example, in molten $\text{NaCl}$, $\text{Na}^+$ is reduced to sodium at the cathode and $\text{Cl}^-$ is oxidised to chlorine at the anode.
Faraday's laws of electrolysis. Faraday quantified electrolysis in two laws:
- First law: the mass of a substance deposited or liberated at an electrode is proportional to the quantity of charge passed: $m \propto Q$, i.e. $m = ZQ = ZIt$, where $Z$ is the electrochemical equivalent, $Q = It$ the charge ($I$ in amperes, $t$ in seconds).
- Second law: when the same charge passes through different electrolytes, the masses deposited are proportional to their equivalent masses. One mole of electrons ($1\,F = 96500\,\text{C}$) deposits one gram-equivalent of any substance.
So $m = \dfrac{M \times I \times t}{n \times F}$, where $M$ is the molar mass and $n$ the number of electrons per ion.
Products of electrolysis depend on the electrode material and on which species is easier to discharge (its electrode potential and overpotential). In aqueous $\text{NaCl}$ (brine), $\text{H}_2$ is liberated at the cathode and $\text{Cl}_2$ at the anode (overpotential favours chlorine over oxygen), leaving $\text{NaOH}$ — the basis of the chlor-alkali industry.
Batteries. A battery is a galvanic cell (or several in series) used as a practical source of electrical energy.
- Primary cells cannot be recharged — the reaction is irreversible. The dry cell (Leclanché cell) has a zinc anode, a graphite cathode surrounded by $\text{MnO}_2$ and carbon, with a moist $\text{NH}_4\text{Cl}/\text{ZnCl}_2$ paste; it gives about $1.5\,\text{V}$. The mercury cell ($\text{Zn}$/$\text{HgO}$) gives a steady $1.35\,\text{V}$ and is used in hearing aids and watches.
- Secondary cells can be recharged. The lead storage battery (car battery) has a $\text{Pb}$ anode and a $\text{PbO}_2$ cathode in $\sim38\%$ $\text{H}_2\text{SO}_4$; on discharge both electrodes form $\text{PbSO}_4$, and charging reverses this. The nickel–cadmium (Ni–Cd) cell uses $\text{Cd}$ and $\text{NiO(OH)}$, has a long life and is rechargeable.
Fuel cells. A fuel cell converts the energy of combustion of a fuel directly into electricity, with no thermal step, so it is highly efficient (~70%) and clean. In the hydrogen–oxygen fuel cell, $\text{H}_2$ is oxidised at the anode and $\text{O}_2$ reduced at the cathode in $\text{KOH}$; the only product is water. These powered the Apollo space programme.
Corrosion. Corrosion is the slow electrochemical eating-away of a metal by its environment — most familiarly the rusting of iron, hydrated iron(III) oxide $\text{Fe}_2\text{O}_3\cdot x\text{H}_2\text{O}$. On a wet iron surface, one spot acts as an anode ($\text{Fe} \rightarrow \text{Fe}^{2+} + 2e^-$) and another as a cathode where oxygen is reduced ($\text{O}_2 + 4\text{H}^+ + 4e^- \rightarrow 2\text{H}_2\text{O}$); the $\text{Fe}^{2+}$ is then oxidised by air to rust. Water and oxygen are both essential.
Prevention works by breaking this mini-cell: barrier methods (painting, greasing, electroplating with tin or chromium), galvanising (a sacrificial zinc coat), and cathodic protection — connecting the iron to a more reactive metal such as magnesium or zinc that corrodes preferentially as a sacrificial anode.