Electroplating, Extracting Metals and Electrolysis in Industry

1From a spoon to a smelter

A shiny tap, a silver-coloured spoon, a soft-drink can, a bottle of bleach and the chlorine in treated drinking water all depend on electrolysis. So does a sheet of aluminium foil: aluminium is the most abundant metal in the Earth's crust, yet it was a costly rarity until 1886, when Charles Martin Hall in the USA and Paul Héroult in France independently found how to extract it by electrolysis. This lesson brings together the uses named in the syllabus.

2Electroplating

Electroplating coats an object with a thin layer of a metal by electrolysis. The rules are always the same (Figure 5.1):

• the object to be plated is the cathode;
• the plating metal is the anode;
• the electrolyte contains ions of the plating metal.

For silver plating: anode $\mathrm{A}\mathrm{g}(\mathrm{s})\to \mathrm{A}{\mathrm{g}}^{+}(\mathrm{a}\mathrm{q})+{\mathrm{e}}^{-}\backslash mathrm\{Ag(s)\; \backslash rightarrow\; Ag^+(aq)\; +\; e^-\}$; cathode $\mathrm{A}{\mathrm{g}}^{+}(\mathrm{a}\mathrm{q})+{\mathrm{e}}^{-}\to \mathrm{A}\mathrm{g}(\mathrm{s})\backslash mathrm\{Ag^+(aq)\; +\; e^-\; \backslash rightarrow\; Ag(s)\}$. Plating is used to stop corrosion (chromium or nickel on steel taps and bicycle parts, tin on the steel of food cans, zinc on screws), to improve appearance (silver and gold on cutlery and jewellery), and to give a cheap metal a costly surface.

3Extracting reactive metals

Metals above carbon in the reactivity series cannot be extracted by heating their oxides with carbon, so they are extracted by electrolysis of a molten compound.

Aluminium. Purified aluminium oxide (from the ore bauxite) melts at over 2000 °C, far too hot to be practical. Instead it is dissolved in molten cryolite, $\mathrm{N}{\mathrm{a}}_3\mathrm{A}\mathrm{l}{\mathrm{F}}_6\backslash mathrm\{Na\_3AlF\_6\}$, at about 950 °C (Figure 5.2).

$$\text{cathode: }\mathrm{A}{\mathrm{l}}^{3+}+3{\mathrm{e}}^{-}\to \mathrm{A}\mathrm{l}(\mathrm{l})\text{anode: }2{\mathrm{O}}^{2-}\to {\mathrm{O}}_2(\mathrm{g})+4{\mathrm{e}}^{-}\backslash text\{cathode:\; \}\; \backslash mathrm\{Al^\{3+\}\; +\; 3e^-\; \backslash rightarrow\; Al(l)\}\; \backslash qquad\; \backslash text\{anode:\; \}\; \backslash mathrm\{2O^\{2-\}\; \backslash rightarrow\; O\_2(g)\; +\; 4e^-\}$$

At this temperature the oxygen reacts with the carbon anodes, $\mathrm{C}(\mathrm{s})+{\mathrm{O}}_2(\mathrm{g})\to \mathrm{C}{\mathrm{O}}_2(\mathrm{g})\backslash mathrm\{C(s)\; +\; O\_2(g)\; \backslash rightarrow\; CO\_2(g)\}$, so the anodes burn away and must be replaced regularly. The process uses enormous amounts of electricity.

Sodium. Sodium is made by electrolysing molten sodium chloride (Lesson 1) in a Downs cell. Calcium chloride is added to lower the melting point from 801 °C to about 600 °C, saving energy; the sodium and chlorine are kept apart so they cannot recombine.

4Chlorine, hydrogen and sodium hydroxide: the chlor-alkali industry

Electrolysing concentrated brine (Lesson 2) gives three useful products at once. In a modern membrane cell, a membrane lets sodium ions through but keeps the chlorine away from the hydrogen and the sodium hydroxide.

$$2\mathrm{N}\mathrm{a}\mathrm{C}\mathrm{l}(\mathrm{a}\mathrm{q})+2{\mathrm{H}}_2\mathrm{O}(\mathrm{l})\to 2\mathrm{N}\mathrm{a}\mathrm{O}\mathrm{H}(\mathrm{a}\mathrm{q})+{\mathrm{H}}_2(\mathrm{g})+\mathrm{C}{\mathrm{l}}_2(\mathrm{g})\backslash mathrm\{2NaCl(aq)\; +\; 2H\_2O(l)\; \backslash rightarrow\; 2NaOH(aq)\; +\; H\_2(g)\; +\; Cl\_2(g)\}$$

Table 5.1: Products of electrolysis and their uses

Electrolysis of water, made conducting by adding an electrolyte such as potassium hydroxide, gives very pure hydrogen and oxygen. When the electricity comes from hydroelectric, solar or wind power, the hydrogen is made without burning fuel.