# 10.6 Electrolysis  (Page 2/7)

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Now consider the cathode. Three reductions could occur:

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

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 −7 M . At such low concentrations, reaction (iii) is unlikely and reaction (iv) occurs. The overall reaction is then

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

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.

## Electroplating

An important use for electrolytic cells is in electroplating    . Electroplating results in a thin coating of one metal on top of a conducting surface. Reasons for electroplating include making the object more corrosion resistant, strengthening the surface, producing a more attractive finish, or for purifying metal. The metals commonly used in electroplating include cadmium, chromium, copper, gold, nickel, silver, and tin. Common consumer products include silver-plated or gold-plated tableware, chrome-plated automobile parts, and jewelry. We can get an idea of how this works by investigating how silver-plated tableware is produced ( [link] ).

In the figure, the anode consists of a silver electrode, shown on the left. The cathode is located on the right and is the spoon, which is made from inexpensive metal. Both electrodes are immersed in a solution of silver nitrate. As the potential is increased, current flows. Silver metal is lost at the anode as it goes into solution.

$\text{anode: Ag}\left(s\right)\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}{\text{Ag}}^{\text{+}}\left(aq\right)+{\text{e}}^{\text{−}}$

The mass of the cathode increases as silver ions from the solution are deposited onto the spoon

${\text{cathode: Ag}}^{\text{+}}\left(aq\right)+{\text{e}}^{\text{−}}\phantom{\rule{0.2em}{0ex}}⟶\phantom{\rule{0.2em}{0ex}}\text{Ag}\left(s\right)$

The net result is the transfer of silver metal from the anode to the cathode. The quality of the object is usually determined by the thickness of the deposited silver and the rate of deposition.

## Quantitative aspects of electrolysis

The amount of current that is allowed to flow in an electrolytic cell is related to the number of moles of electrons. The number of moles of electrons can be related to the reactants and products using stoichiometry. Recall that the SI unit for current ( I ) is the ampere (A), which is the equivalent of 1 coulomb per second (1 A = 1 $\frac{\text{C}}{\text{s}}$ ). The total charge ( Q , in coulombs) is given by

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