Electrolysis is a process that uses direct current (DC) to drive a chemical reaction that would not occur spontaneously. This technique involves passing an electric current through an electrolyte—either a molten ionic compound or an aqueous solution—causing the ions to move and react at the electrodes.

The word "electrolysis" originates from the Greek words "electron" (amber, associated with electricity) and "lysis" (to loosen or split). This method was first investigated by Michael Faraday in the early 19th century, who formulated the fundamental laws of electrolysis. These laws state that the amount of substance liberated at each electrode is directly proportional to the quantity of electricity passed through the electrolyte.

Electrolysis is widely used in industries for extracting and refining metals, manufacturing chemicals, and producing gases.

Electrolysis of Molten Compounds

In molten compounds, the electrolysis process involves the movement of ions when the compound is heated to a high temperature until it melts. In this state, the ionic bonds are broken, allowing the ions to move freely.

The process involves cations migrating towards the cathode (negative electrode) and being discharged by gaining electrons, while anions move towards the anode (positive electrode) and are discharged by losing electrons.

For example, in the electrolysis of molten lead bromide (PbBr₂):

• At the cathode: lead ions (Pb²⁺) gain electrons (reduction) and are discharged as solid lead (Pb)
• At the anode: bromide ions (Br⁻) are discharged by losing electrons (oxidation), releasing bromine gas (Br₂) at the anode

The products are thus collected at the electrodes, with molten lead accumulating at the bottom and bromine gas being released. This setup requires inert electrodes, such as graphite or platinum, to prevent the electrodes from reacting with the products.

Molten electrolysis is crucial in extracting metals like aluminium, where aluminium oxide is electrolysed to produce aluminium metal and oxygen gas.

Electrolysis of Aqueous Solutions

In aqueous solutions, both the dissolved ionic compound and water can provide ions. Water dissociates into H⁺ and OH⁻ ions, which can compete with the ions of the dissolved compound to be discharged at the electrodes. The actual ions discharged depend on their reactivity and concentration.

At the cathode, reduction occurs as cations gain electrons and are discharged. The order of discharge preference is based on the reactivity of the ions:

1. Least reactive metals (e.g., copper (Cu²⁺), silver (Ag⁺)) are discharged first because they more readily accept electrons to form neutral atoms.
2. Hydrogen ions (H⁺) are discharged as hydrogen gas (H₂) if the solution contains more reactive metal ions.
3. More reactive metals (e.g., sodium (Na⁺), potassium (K⁺)) are not typically discharged in aqueous solutions because water is reduced instead, producing hydrogen gas.

At the anode, oxidation occurs as anions lose electrons and are discharged. The discharge order typically follows:

1. Halide ions (Cl⁻, Br⁻, I⁻) are preferentially discharged as their respective halogen gases (Cl₂, Br₂, I₂), because they require less energy to oxidise than other anions.
2. If no halide ions are present or in low concentration, hydroxide ions (OH⁻) are discharged as oxygen (O₂) gas.
3. Other anions, such as sulfate (SO₄²⁻), are generally not discharged as they are more challenging to oxidise compared to OH⁻.

For instance, in the electrolysis of sodium chloride solution (NaCl), chlorine gas (Cl₂) is discharged at the anode due to the presence of chloride ions, while hydrogen gas (H₂) is discharged at the cathode.

Electrolysis of Copper

Aluminium Extraction

Half Equations and Redox Reactions