Electrolysis

Electrolysis

Ions of metals in of groups 1 and 2, along with aluminum, are very difficult to reduce; therefore, it is necessary to prepare these elements by electrolysis, an important process discussed in the tutorial on electrochemistry. Briefly, electrolysis involves using electrical energy to drive unfavorable chemical reactions to completion; it is useful in the isolation of reactive metals in their pure forms. Sodium, aluminum, and magnesium are typical examples.

The Preparation of Sodium

The most important method for the production of sodium is the electrolysis of molten sodium chloride; the set-up is a Downs cell, shown in the figure below. The reaction involved in this process is:

\(\text{2NaCl}(l)\;\underset{600\;\text{°C}}{\overset{\phantom{\rule{0.5em}{0ex}}\text{electrolysis}\phantom{\rule{0.5em}{0ex}}}{\to }}\;\text{2Na}(l)+{\text{Cl}}_{2}(g)\)

The electrolysis cell contains molten sodium chloride (melting point 801 °C), to which calcium chloride has been added to lower the melting point to 600 °C (a colligative effect). The passage of a direct current through the cell causes the sodium ions to migrate to the negatively charged cathode and pick up electrons, reducing the ions to sodium metal. Chloride ions migrate to the positively charged anode, lose electrons, and undergo oxidation to chlorine gas. The overall cell reaction comes from adding the following reactions:

\(\begin{array}{l}\text{at the cathode:}\;{\text{2Na}}^{+}+{\text{2e}}^{-}\;⟶\phantom{\rule{0.4em}{0ex}}\text{2Na}(l)\\ \text{at the anode:}\;{\text{2Cl}}^{-}\;⟶\;{\text{Cl}}_{2}(g)+{\text{2e}}^{-}\\ \text{overall change:}\;{\text{2Na}}^{+}+{\text{2Cl}}^{-}\;⟶\;\text{2Na}(l)+{\text{Cl}}_{2}(g)\end{array}\)

Separation of the molten sodium and chlorine prevents recombination. The liquid sodium, which is less dense than molten sodium chloride, floats to the surface and flows into a collector. The gaseous chlorine goes to storage tanks. Chlorine is also a valuable product.

Pure sodium metal is isolated by electrolysis of molten sodium chloride using a Downs cell. It is not possible to isolate sodium by electrolysis of aqueous solutions of sodium salts because hydrogen ions are more easily reduced than are sodium ions; as a result, hydrogen gas forms at the cathode instead of the desired sodium metal. The high temperature required to melt NaCl means that liquid sodium metal forms.

The Preparation of Aluminum

The preparation of aluminum utilizes a process invented in 1886 by Charles M. Hall, who began to work on the problem while a student at Oberlin University in Ohio. Paul L. T. Héroult discovered the process independently a month or two later in France. In honor to the two inventors, this electrolysis cell is known as the Hall–Héroult cell. The Hall–Héroult cell is an electrolysis cell for the production of aluminum. The figure below illustrates the Hall–Héroult cell.

The production of aluminum begins with the purification of bauxite, the most common source of aluminum. The reaction of bauxite, AlO(OH), with hot sodium hydroxide forms soluble sodium aluminate, while clay and other impurities remain undissolved:

\(\text{AlO}(\text{OH})(s)+\text{NaOH}(aq)+{\text{H}}_{2}\text{O}(l)\;⟶\;\text{Na}[\text{Al}{(\text{OH})}_{4}](aq)\)

After the removal of the impurities by filtration, the addition of acid to the aluminate leads to the reprecipitation of aluminum hydroxide:

\(\text{Na}[\text{Al}{(\text{OH})}_{4}](aq)+{\text{H}}_{3}{\text{O}}^{\text{+}}(aq)\;⟶\;\text{Al}{(\text{OH})}_{3}(s)+{\text{Na}}^{\text{+}}(aq)+{\text{2H}}_{2}\text{O}(l)\)

The next step is to remove the precipitated aluminum hydroxide by filtration. Heating the hydroxide produces aluminum oxide, Al₂O₃, which dissolves in a molten mixture of cryolite, Na₃AlF₆, and calcium fluoride, CaF₂. Electrolysis of this solution takes place in a cell like that shown in the figure below. Reduction of aluminum ions to the metal occurs at the cathode, while oxygen, carbon monoxide, and carbon dioxide form at the anode.

An electrolytic cell is used for the production of aluminum. The electrolysis of a solution of cryolite and calcium fluoride results in aluminum metal at the cathode, and oxygen, carbon monoxide, and carbon dioxide at the anode.

The Preparation of Magnesium

Magnesium is the other metal that is isolated in large quantities by electrolysis. Seawater, which contains approximately 0.5% magnesium chloride, serves as the major source of magnesium. Addition of calcium hydroxide to seawater precipitates magnesium hydroxide. The addition of hydrochloric acid to magnesium hydroxide, followed by evaporation of the resultant aqueous solution, leaves pure magnesium chloride. The electrolysis of molten magnesium chloride forms liquid magnesium and chlorine gas:

\({\text{MgCl}}_{2}(aq)+\text{Ca}{(\text{OH})}_{2}(aq)\;⟶\;\text{Mg}{(\text{OH})}_{2}(s)+{\text{CaCl}}_{2}(aq)\)

\(\text{Mg}{(\text{OH})}_{2}(s)+\text{2HCl}(aq)\;⟶\;{\text{MgCl}}_{2}(aq)+{\text{2H}}_{2}\text{O}(l)\)

\({\text{MgCl}}_{2}(l)\;⟶\;\text{Mg}(l)+{\text{Cl}}_{2}(g)\)

Some production facilities have moved away from electrolysis completely. In the next section, we will see how the Pidgeon process leads to the chemical reduction of magnesium.