Electrolysis of Molten and Aqueous Compounds
What Is Electrolysis?
Scope note: electrolysis is topic C3.4, examined on Chemistry Paper 1 (Foundation) or Paper 3 (Higher), and assumed knowledge for the Paper 2 topics.
Electrolysis is the process of using electricity to decompose an ionic compound into its elements. The compound being broken down is called the electrolyte, and it must be either molten or dissolved in water so that the ions are free to move and carry charge.
The electrolyte is placed between two electrodes: the cathode (negative electrode) and the anode (positive electrode). Positive ions (cations) migrate to the cathode; negative ions (anions) migrate to the anode. A useful mnemonic is PANIC: Positive Anode, Negative Is Cathode.
At the cathode (negative): positive ions gain electrons — this is reduction. At the anode (positive): negative ions lose electrons — this is oxidation. Remember: electrolysis involves both oxidation and reduction simultaneously.
Electrolysis of Molten Compounds
When a molten ionic compound is electrolysed, only the ions of that compound are present. The metal ion moves to the cathode and is deposited as the metal; the non-metal ion moves to the anode and is released as the non-metal element. For example, molten lead bromide produces lead metal at the cathode and bromine gas at the anode. The key requirement is that the compound is molten — solid ionic compounds cannot conduct electricity because the ions are held in a fixed lattice and are not free to move. Dissolving the compound in water also frees the ions, though this introduces additional complexity because water provides its own ions.
It is important to understand that the electrolyte is not simply “broken down” into smaller pieces — it is chemically decomposed into its constituent elements. This makes electrolysis a powerful tool for extracting reactive metals that cannot be produced by simpler reduction methods. The external electrical supply acts as an electron pump: the power supply removes electrons from the anode (making it positive) and pushes electrons into the cathode (making it negative). Ions in the electrolyte carry charge through the solution between the electrodes, completing the circuit.
For molten lead bromide electrolysis:
Cathode (reduction): \(Pb^{2+} + 2e^- \rightarrow Pb\)
Anode (oxidation): \(2Br^- \rightarrow Br_2 + 2e^-\)
Half-equations must be balanced for both atoms and charge. At the cathode, electrons are gained (shown on the left side of the arrow). At the anode, electrons are released (shown on the right side). The number of electrons must cancel when the two half-equations are combined.
Electrolysis of Aqueous Solutions
When an ionic compound is dissolved in water, the solution also contains \(H^+\) and \(OH^-\) ions from the water itself. This means there is competition at each electrode between different ions. The rules for predicting products are:
- Cathode: If the metal is more reactive than hydrogen, hydrogen gas is produced from \(H^+\) ions. If the metal is less reactive than hydrogen, the metal is deposited.
- Anode: If a halide ion (Cl−, Br−, I−) is present in large amounts, the halogen is produced. If no halide is present, oxygen is produced (from \(OH^-\) ions).
Brine is concentrated sodium chloride solution. During electrolysis of brine, three useful products are obtained:
- Cathode: hydrogen gas is produced (sodium is too reactive to be deposited; \(H^+\) ions are discharged instead).
- Anode: chlorine gas is produced (chloride ions are present in high concentration and are preferentially discharged).
- Remaining solution: sodium hydroxide (NaOH) remains in solution, as the sodium and hydroxide ions are not discharged.
The three products — hydrogen, chlorine, and sodium hydroxide — are all industrially important: chlorine is used to make bleach and PVC; hydrogen is used as a fuel; sodium hydroxide is used to make soap and paper.
Extraction of Aluminium
Aluminium is too reactive to be extracted by reduction with carbon. Instead, it is extracted by electrolysis of molten aluminium oxide (\(Al_2O_3\)). The aluminium oxide is dissolved in molten cryolite to lower the melting point and reduce energy costs.
Aluminium is deposited at the cathode (\(Al^{3+}\) ions gain electrons). Oxygen is released at the anode (\(O^{2-}\) ions lose electrons). The carbon anodes gradually burn away because the oxygen produced reacts with them at high temperature, so they need regular replacement. The process uses large amounts of electricity, making aluminium expensive to produce.
Practical Electrolysis (PAG C2)
In your practical work (PAG C2), you will investigate the electrolysis of aqueous solutions using inert (carbon) electrodes. A typical investigation uses copper sulfate solution or brine. To test the gases produced: hold a lit splint near the cathode to detect hydrogen (squeaky pop); hold damp litmus paper near the anode to detect chlorine (paper bleaches) or use a glowing splint to detect oxygen (relights). Record which gas appears at which electrode and compare your results with predictions from competition rules.
Electroplating is another important application of electrolysis. To electroplate an object with copper: make the object the cathode, use a copper electrode as the anode, and use copper sulfate solution as the electrolyte. Copper atoms at the anode dissolve into solution as copper ions; the same mass of copper deposits on the cathode. The object gains a thin, evenly spread copper layer. Electroplating is used in jewellery (silver plating), corrosion protection (zinc plating on steel, called galvanising) and in electronics (copper tracks on circuit boards).
Higher Tier half-equations for aluminium extraction:
Cathode: \(Al^{3+} + 3e^- \rightarrow Al\)
Anode: \(2O^{2-} \rightarrow O_2 + 4e^-\)