The Electrolysis of Aqueous Sodium Chloride

The Electrolysis of Aqueous Sodium Chloride

The electrolysis of aqueous sodium chloride is the more common example of electrolysis because more than one species can be oxidized and reduced. Considering the anode first, the possible reactions are

\(\begin{array}{l}(\text{i})\phantom{\rule{0.2em}{0ex}}2{\text{Cl}}^{\text{−}}(aq)\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}{\text{Cl}}_{2}(g)+2\phantom{\rule{0.2em}{0ex}}{\text{e}}^{\text{−}}\phantom{\rule{4em}{0ex}}{E}_{\text{anode}}^{°}=\text{+1.35827 V}\\ (\text{ii})\phantom{\rule{0.2em}{0ex}}2{\text{H}}_{2}\text{O}(l)\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}{\text{O}}_{2}(g)+4{\text{H}}^{\text{+}}(aq)+4{\text{e}}^{\text{−}}\phantom{\rule{4em}{0ex}}{E}_{\text{anode}}^{°}=\text{+1.229 V}\end{array}\)

These values suggest that water should be oxidized at the anode because a smaller potential would be needed—using reaction (ii) for the oxidation would give a less-negative cell potential. When the experiment is run, it turns out chlorine, not oxygen, is produced at the anode. The unexpected process is so common in electrochemistry that it has been given the name overpotential. The overpotential is the difference between the theoretical cell voltage and the actual voltage that is necessary to cause electrolysis. It turns out that the overpotential for oxygen is rather high and effectively makes the reduction potential more positive. As a result, under normal conditions, chlorine gas is what actually forms at the anode.

Now consider the cathode. Three reductions could occur:

\(\begin{array}{l}(\text{iii})\phantom{\rule{0.2em}{0ex}}{\text{2H}}^{\text{+}}(aq)+2{\text{e}}^{\text{−}}\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}{\text{H}}_{2}(g)\phantom{\rule{4em}{0ex}}{E}_{\text{cathode}}^{°}=\text{0 V}\\ (\text{iv})\phantom{\rule{0.2em}{0ex}}{\text{2H}}_{2}\text{O}(l)+2{\text{e}}^{\text{−}}\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}{\text{H}}_{2}(g)+2{\text{OH}}^{\text{−}}(aq)\phantom{\rule{4em}{0ex}}{E}_{\text{cathode}}^{°}=\text{−0.8277 V}\\ (\text{v})\phantom{\rule{0.2em}{0ex}}{\text{Na}}^{\text{+}}(aq)+{\text{e}}^{\text{−}}\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}\text{Na}(s)\phantom{\rule{4em}{0ex}}{E}_{\text{cathode}}^{°}=\text{−2.71 V}\end{array}\)

Reaction (v) is ruled out because it has such a negative reduction potential. Under standard state conditions, reaction (iii) would be preferred to reaction (iv). However, the pH of a sodium chloride solution is 7, so the concentration of hydrogen ions is only 1\(×\) 10⁻⁷M. At such low concentrations, reaction (iii) is unlikely and reaction (iv) occurs. The overall reaction is then

\(\text{overall: 2}{\text{H}}_{2}\text{O}(l)+2{\text{Cl}}^{\text{−}}(aq)\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}{\text{H}}_{2}(g)+{\text{Cl}}_{2}(g)+2{\text{OH}}^{\text{−}}(aq)\phantom{\rule{4em}{0ex}}{E}_{\text{cell}}^{°}=-2.186 V\)

As the reaction proceeds, hydroxide ions replace chloride ions in solution. Thus, sodium hydroxide can be obtained by evaporating the water after the electrolysis is complete. Sodium hydroxide is valuable in its own right and is used for things like oven cleaner, drain opener, and in the production of paper, fabrics, and soap.