In the Core syllabus, redox reactions are defined by the gain or loss of oxygen atoms. However, many significant reactions—such as the formation of sodium chloride from its elements or the processes occurring inside a battery—do not involve oxygen at all. To unify these reactions under a single theory, chemists use the concept of electron transfer. The IGCSE 0654 Extended syllabus requires a formal definition of oxidation and reduction in terms of electrons and the ability to represent these changes using ionic half-equations.

The most robust definition of redox focuses on the movement of electrons between chemical species.

• Oxidation: The loss of electrons by a substance.
• Reduction: The gain of electrons by a substance.

The OIL RIG Mnemonic

To recall these definitions, chemists universally use the acronym OIL RIG:

• Oxidation Is Loss
• Reduction Is Gain

Electrons are subatomic particles and cannot exist freely in a stable chemical environment. Therefore, if one atom loses an electron, another atom must immediately gain it.

• Redox Integrity: Oxidation and reduction are two sides of the same coin; they always occur together. The total number of electrons lost by the species being oxidized must equal the total number gained by the species being reduced.

Ionic half-equations are used to focus on what happens to a single element during a redox reaction. They explicitly show the electrons ($e^{-}$) involved.

(a) Oxidation Half-Equations

In oxidation, electrons are lost. Consequently, the electrons are written on the product side (right) of the equation.

• Example (Magnesium): $Mg \rightarrow Mg^{2+} + 2e^{-}$
• Example (Iron II to III): $Fe^{2+} \rightarrow Fe^{3+} + e^{-}$

(b) Reduction Half-Equations

In reduction, electrons are gained. The electrons are written on the reactant side (left) of the equation.

• Example (Copper): $Cu^{2+} + 2e^{-} \rightarrow Cu$
• Example (Chlorine): $Cl_2 + 2e^{-} \rightarrow 2Cl^{-}$

The oxidation number of an element is a measure of its "electron control."

• When an element is oxidized (loses negative electrons), its oxidation number increases (becomes more positive).
• When an element is reduced (gains negative electrons), its oxidation number decreases (becomes more negative).

The electronic definition of redox provides a universal framework for understanding chemical change. By identifying the loss and gain of electrons, chemists can explain the behavior of metals and non-metals across all branches of the science, from the extraction of elements to the complex biological electron transport chains that power life. Mastery of half-equations is the key to quantifying these subatomic transitions.