Electrode Reactions, Molten Electrolysis and Aluminium Extraction

Why an electrolyte must have moving ions

Electrolysis is the decomposition of an electrolyte by an electric current. It uses electrical energy to drive chemical reactions, separating a compound into simpler substances.

An ionic compound contains positively and negatively charged ions. In a solid, these ions are held in fixed positions in a lattice, so they cannot move through the substance to carry charge. When the compound melts, its ions become free to move and it can conduct electricity. Dissolving an ionic compound in water can also release moving ions, although water introduces additional possible electrode reactions.

The conducting liquid is the electrolyte. Two conducting electrodes dip into it and connect to a direct-current power supply. Charge is carried by moving ions within the electrolyte and by electrons in the external wires.

What happens at each electrode

The cathode is the negative electrode. It attracts positive ions, which gain electrons at its surface. Gain of electrons is reduction. Metal ions become neutral metal atoms; hydrogen can also form at the cathode when suitable hydrogen-containing ions are present.

The anode is the positive electrode. It attracts negative ions, which lose electrons at its surface. Loss of electrons is oxidation. With inert electrodes, the products at the anode are non-metals.

An inert electrode conducts electricity without taking part in the chemical reactions under the conditions used. Platinum or graphite can serve this purpose in suitable electrolytes. Remember the electron-transfer distinction as OIL RIG: Oxidation Is Loss, Reduction Is Gain.

Predicting products from a molten binary compound

A binary ionic compound contains only two elements. When it is molten, there is no water to introduce competing reactions. Identify its positive and negative ions: the metal forms at the cathode and the non-metal forms at the anode, provided the electrodes do not react with the products.

For example, molten lead(II) bromide, $\mathrm{PbBr_2}$, contains $\mathrm{Pb^{2+}}$ ions and $\mathrm{Br^-}$ ions. Lead ions move towards the cathode and gain electrons to form lead. Bromide ions move towards the anode and lose electrons to form bromine, .

The overall decomposition is:

$$\mathrm{PbBr_2 \rightarrow Pb + Br_2}$$

The same method applies to other molten binary salts: molten sodium chloride produces sodium and chlorine, while molten aluminium oxide produces aluminium and oxygen at the electrode surfaces.

Writing half-equations that conserve atoms and charge

A half-equation describes the electron transfer at one electrode. It must conserve both the number of atoms of each element and the total electrical charge.

At the cathode, each lead ion has a charge of $2+$. It needs two electrons to become a neutral lead atom:

$$\mathrm{Pb^{2+} + 2e^- \rightarrow Pb}$$

There is one lead atom on each side. The total charge on the left is $+2-2=0$, matching the neutral product. Lead ions are reduced because they gain electrons.

At the anode, bromine forms molecules containing two bromine atoms. Start with two bromide ions to supply those atoms, then place the electrons lost on the right:

$$\mathrm{2Br^- \rightarrow Br_2 + 2e^-}$$

Both sides contain two bromine atoms and have a total charge of $-2$. Bromide ions are oxidised because they lose electrons.

To construct a half-equation, write the ion and product, balance the atoms, then add electrons to balance charge. Finally, check that electron gain or loss matches the reaction you are describing. Change coefficients, not the chemical formulae of substances.

Extracting aluminium from its ore

Aluminium is more reactive than carbon, so carbon cannot be used to reduce aluminium oxide to aluminium. Instead, electrical energy is used to reduce aluminium ions by electrolysis.

The aluminium ore bauxite is first purified to obtain aluminium oxide, $\mathrm{Al_2O_3}$. Pure aluminium oxide melts above $2000\,^\circ\mathrm{C}$, making it expensive to melt. It is therefore dissolved in . This allows the to remain liquid at a lower temperature, reducing the energy needed to keep the ions mobile.

The industrial cell has a steel shell lined with graphite. The graphite lining is the negative cathode, while large graphite blocks dipping into the electrolyte are the positive anodes.

Aluminium ions gain three electrons each at the cathode:

$$\mathrm{Al^{3+} + 3e^- \rightarrow Al}$$

The aluminium ions are reduced. Molten aluminium collects at the bottom of the cell and is removed periodically. Fresh aluminium oxide is added to replace that used up.

Oxide ions lose electrons at the anode, forming oxygen:

$$\mathrm{2O^{2-} \rightarrow O_2 + 4e^-}$$

The oxide ions are oxidised. Each loses two electrons, so two oxide ions release four electrons altogether.

Why the carbon anodes need replacing

The graphite anodes are not completely inert in the hot industrial cell. Oxygen formed at their surfaces reacts with the carbon:

$$\mathrm{C + O_2 \rightarrow CO_2}$$

This consumes the anodes, so they must be replaced regularly. It also releases carbon dioxide. Aluminium extraction uses large amounts of electricity, which is a major cost: cryolite reduces the heating requirement but does not remove the need for electrical energy to drive electrolysis.

Electrode reactions

• Electrolysis decomposes an electrolyte using an electric current.
• Molten or dissolved ionic compounds conduct because their ions can move; solid ionic compounds do not.
• Cathode (−): positive ions gain electrons — reduction; metals or hydrogen form.