Reduction and Oxidation

Definition of Oxidation and Reduction

Oxidation and reduction are processes that involve the transfer of electrons between substances.

• Oxidation is the loss of electrons by an atom, ion, or molecule.
• Reduction is the gain of electrons by an atom, ion, or molecule.

Because electrons lost by one species are gained by another, oxidation and reduction always happen together in a redox reaction.

For example, when magnesium reacts with oxygen, magnesium atoms lose electrons (oxidised) and oxygen atoms gain electrons (reduced).

Remember: OIL RIG — Oxidation Is Loss, Reduction Is Gain (of electrons).

Oxidation States

Oxidation states (or numbers) are used to keep track of how many electrons an atom has gained or lost compared to its elemental form.

• In an element by itself, the oxidation state is 0 (e.g., O₂, Fe).
• For ions, the oxidation state equals the charge (e.g., Na⁺ is +1, Cl⁻ is −1).
• In compounds, oxidation states are assigned based on known rules (e.g., oxygen is usually −2, hydrogen +1).

During a redox reaction, the oxidation state of an element changes:

• If the oxidation state increases, the element is oxidised (loses electrons).
• If the oxidation state decreases, the element is reduced (gains electrons).

This helps identify which species are oxidised or reduced in a reaction.

For instance, in the reaction between magnesium and oxygen:

Magnesium changes from 0 to +2 (oxidised), oxygen changes from 0 to −2 (reduced).

Redox in Metal Reactivity

Metals tend to lose electrons and become positively charged ions. This means metals are usually oxidised in reactions.

Non-metals tend to gain electrons and become negatively charged ions, meaning they are reduced.

The more reactive a metal is, the more easily it loses electrons (oxidised).

For example, potassium is more reactive than iron because potassium loses electrons more readily.

This explains why metals react differently with substances like oxygen or acids — their tendency to be oxidised varies.

For example, when magnesium reacts with oxygen:

Magnesium atoms lose electrons to form Mg²⁺ ions (oxidised), while oxygen gains electrons to form O²⁻ ions (reduced).

Examples of Redox Reactions

Metal + Oxygen → Metal Oxide

When metals react with oxygen, they form metal oxides. The metal atoms lose electrons (oxidation) and oxygen atoms gain electrons (reduction).

For example, magnesium reacts with oxygen:

Magnesium is oxidised: Mg (0) → Mg²⁺ (+2)

Oxygen is reduced: O₂ (0) → O²⁻ (−2)

The overall reaction is:

$2 \mathrm{Mg} + \mathrm{O}_2 \rightarrow 2 \mathrm{MgO}$

Metal Displacement Reactions

A more reactive metal can displace a less reactive metal from its compound. The more reactive metal is oxidised (loses electrons), and the less reactive metal ion is reduced (gains electrons).

For example, zinc displaces copper from copper sulfate solution:

Zinc is oxidised: Zn (0) → Zn²⁺ (+2)

Copper is reduced: Cu²⁺ (+2) → Cu (0)

Overall reaction:

$\mathrm{Zn} + \mathrm{CuSO}_4 \rightarrow \mathrm{ZnSO}_4 + \mathrm{Cu}$

Extraction of Metals by Reduction

Many metals are extracted from their ores by reduction, where the metal oxide is reduced to the metal by removing oxygen.

In this process, the metal ion gains electrons (reduced), and the reducing agent loses electrons (oxidised).

For example, iron oxide is reduced to iron in a blast furnace:

Iron ions: Fe³⁺ → Fe (gain of electrons, reduction)

Carbon monoxide acts as the reducing agent and is oxidised:

CO → CO₂ (loss of electrons, oxidation)

This shows how redox reactions are essential in metal extraction.

Example: When magnesium burns in oxygen, it forms magnesium oxide. Magnesium atoms lose two electrons each (oxidised), and oxygen atoms gain electrons (reduced).

This is a redox reaction because both oxidation and reduction occur simultaneously.