Oxidation and reduction defined in terms of oxygen and of electrons (OIL RIG), identifying redox in displacement and other reactions, and writing half-equations.

What this dot point is asking

CCEA wants you to define oxidation and reduction in terms of both oxygen and electrons, identify what is oxidised and reduced in a reaction, and write half-equations showing electron loss and gain.

Two definitions of oxidation and reduction

For example, when magnesium burns it gains oxygen, so it is oxidised; when iron oxide is reduced by carbon to iron, the iron oxide loses oxygen, so it is reduced. The electron view captures the same idea more deeply: magnesium loses electrons as it forms $\text{Mg}^{2+}$.

OIL RIG

Because electrons cannot just disappear, oxidation and reduction always occur together, which is why the combined process is called redox.

Half-equations

A half-equation must balance for both atoms and charge. The electrons lost in the oxidation half-equation equal the electrons gained in the reduction half-equation, so the two can be combined into the full equation.

Worked example

Examples in context

Example 1. Rusting as oxidation

When iron rusts it gains oxygen (and loses electrons) to form iron oxide, so rusting is an oxidation reaction. Recognising corrosion as redox is the first step to understanding how to prevent it.

Example 2. Extracting metals by reduction

Iron is obtained from its ore by reducing iron oxide with carbon, which removes the oxygen. The whole metal-extraction industry rests on redox chemistry, with reduction freeing the metal from its compound.

Example 3. Electrolysis as forced redox

In electrolysis an electric current drives redox: positive ions are reduced at the cathode by gaining electrons, while negative ions are oxidised at the anode by losing them. Recognising electrolysis as a redox process explains why the products are always neutral elements formed from ions, linking this dot point directly to the electrolysis topic.

Oxidising and reducing agents

It helps to name the species responsible for each change. An oxidising agent is the substance that causes oxidation by taking electrons (and is itself reduced); a reducing agent causes reduction by giving electrons (and is itself oxidised). In the blast furnace, carbon monoxide is the reducing agent because it removes oxygen from the iron oxide. In a displacement reaction the more reactive metal is the reducing agent, because it hands its electrons to the less reactive metal's ions. Spotting which species is the oxidising agent and which is the reducing agent is a common CCEA follow-up to identifying what is oxidised and reduced.

Try this

Q1. State what OIL RIG stands for. [1 mark]

• Cue. Oxidation Is Loss, Reduction Is Gain (of electrons).

Q2. Write the half-equation for the oxidation of magnesium. [1 mark]

• Cue. $\text{Mg} \rightarrow \text{Mg}^{2+} + 2e^-$.