Electrolysis of Sodium Chloride: Molten, Dilute and Concentrated

1Three products from salt water

Pass a current through concentrated salt water and you obtain three valuable chemicals at once: chlorine, used to disinfect drinking water and to make PVC; sodium hydroxide, used to make soap, paper and aluminium oxide; and hydrogen. The chlor-alkali industry is one of the largest uses of electricity in chemistry. But change the concentration of the salt, or melt it instead of dissolving it, and the products change. This lesson explains why.

2Molten sodium chloride

Only $\mathrm{N}{\mathrm{a}}^{+}\backslash mathrm\{Na^\{+\}\}$ and $\mathrm{C}{\mathrm{l}}^{-}\backslash mathrm\{Cl^\{-\}\}$ are present, so they must be discharged:

$$\text{cathode: }\mathrm{N}{\mathrm{a}}^{+}+{\mathrm{e}}^{-}\to \mathrm{N}\mathrm{a}(\mathrm{l})\text{anode: }2\mathrm{C}{\mathrm{l}}^{-}\to \mathrm{C}{\mathrm{l}}_2(\mathrm{g})+2{\mathrm{e}}^{-}\backslash text\{cathode:\; \}\; \backslash mathrm\{Na^\{+\}\; +\; e^\{-\}\; \backslash rightarrow\; Na(l)\}\; \backslash qquad\; \backslash text\{anode:\; \}\; \backslash mathrm\{2Cl^\{-\}\; \backslash rightarrow\; Cl\_2(g)\; +\; 2e^\{-\}\}$$

Sodium chloride melts at 801 °C; industrially (the Downs cell) calcium chloride is added to lower the melting point to about 600 °C, saving energy. This is how sodium metal is made.

3Concentrated aqueous sodium chloride (brine)

Now water is present too. At the cathode, water is reduced far more easily than $\mathrm{N}{\mathrm{a}}^{+}\backslash mathrm\{Na^\{+\}\}$ ($E^{\circ }E^\{\backslash circ\}$ −0.83 V against −2.71 V), so hydrogen forms, never sodium. At the anode, chloride ions are oxidised:

$$\text{cathode: }2{\mathrm{H}}_2\mathrm{O}(\mathrm{l})+2{\mathrm{e}}^{-}\to {\mathrm{H}}_2(\mathrm{g})+2\mathrm{O}{\mathrm{H}}^{-}(\mathrm{a}\mathrm{q})\text{anode: }2\mathrm{C}{\mathrm{l}}^{-}(\mathrm{a}\mathrm{q})\to \mathrm{C}{\mathrm{l}}_2(\mathrm{g})+2{\mathrm{e}}^{-}\backslash text\{cathode:\; \}\; \backslash mathrm\{2H\_2O(l)\; +\; 2e^\{-\}\; \backslash rightarrow\; H\_2(g)\; +\; 2OH^\{-\}(aq)\}\; \backslash qquad\; \backslash text\{anode:\; \}\; \backslash mathrm\{2Cl^\{-\}(aq)\; \backslash rightarrow\; Cl\_2(g)\; +\; 2e^\{-\}\}$$

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

Equal volumes of hydrogen and chlorine form (both need 2 electrons per molecule), and $\mathrm{N}{\mathrm{a}}^{+}\backslash mathrm\{Na^\{+\}\}$ and $\mathrm{O}{\mathrm{H}}^{-}\backslash mathrm\{OH^\{-\}\}$ are left in solution. Phenolphthalein turns pink around the cathode; chlorine bleaches damp litmus paper at the anode.

4Dilute aqueous sodium chloride

With few chloride ions, water is oxidised at the anode instead, and oxygen is the main product:

$$\text{anode: }2{\mathrm{H}}_2\mathrm{O}(\mathrm{l})\to {\mathrm{O}}_2(\mathrm{g})+4{\mathrm{H}}^{+}(\mathrm{a}\mathrm{q})+4{\mathrm{e}}^{-}\backslash text\{anode:\; \}\; \backslash mathrm\{2H\_2O(l)\; \backslash rightarrow\; O\_2(g)\; +\; 4H^\{+\}(aq)\; +\; 4e^\{-\}\}$$

The cathode still gives hydrogen, so overall it is water that is electrolysed, and the salt solution slowly becomes more concentrated. At intermediate concentrations the anode gas is a mixture of oxygen and chlorine.

5Why does concentration matter?

From $E^{\circ }E^\{\backslash circ\}$ alone ($+1.23+1.23$ V for ${\mathrm{O}}_2\mathrm{/}{\mathrm{H}}_2\mathrm{O}\backslash mathrm\{O\_2/H\_2O\}$, $+1.36+1.36$ V for $\mathrm{C}{\mathrm{l}}_2\mathrm{/}\mathrm{C}{\mathrm{l}}^{-}\backslash mathrm\{Cl\_2/Cl^\{-\}\}$), water should always be oxidised first. Two effects change this in concentrated brine: (1) the very high concentration of $\mathrm{C}{\mathrm{l}}^{-}\backslash mathrm\{Cl^\{-\}\}$ makes chloride oxidation easier, and (2) oxygen has a large overpotential on the anodes used, so in practice chlorine forms more readily. In dilute solution, too few chloride ions reach the anode, and oxygen wins.

6The membrane cell

Figure 13.2: A membrane cell. The membrane keeps chlorine away from the sodium hydroxide.

If chlorine met the sodium hydroxide, they would react: $\mathrm{C}{\mathrm{l}}_2+2\mathrm{N}\mathrm{a}\mathrm{O}\mathrm{H}\to \mathrm{N}\mathrm{a}\mathrm{C}\mathrm{l}+\mathrm{N}\mathrm{a}\mathrm{C}\mathrm{l}\mathrm{O}+{\mathrm{H}}_2\mathrm{O}\backslash mathrm\{Cl\_2\; +\; 2NaOH\; \backslash rightarrow\; NaCl\; +\; NaClO\; +\; H\_2O\}$. That is how household bleach is made on purpose, but in a chlor-alkali plant it would waste both products. A cation-exchange membrane allows only $\mathrm{N}{\mathrm{a}}^{+}\backslash mathrm\{Na^\{+\}\}$ ions to pass to the cathode side, where they join the $\mathrm{O}{\mathrm{H}}^{-}\backslash mathrm\{OH^\{-\}\}$ ions to give pure sodium hydroxide solution.