Electron transfer: who loses, who gains, who is the agent

By the end of this lesson

By the end you can write half-equations for electron loss and gain, combine them into a full redox equation, and identify the oxidising agent and reducing agent in a reaction.

Two halves of one change

Electrons cannot simply appear or vanish. If one species loses electrons, another must gain them at the same moment.

That is why the two processes always happen together, and why the whole thing is called redox: reduction plus oxidation.

The old definitions used oxygen and hydrogen. The modern ones use electrons, which is far more useful.

Definition

Oxidation and reduction

Oxidation is the loss of electrons. The oxidation number of the element increases.

Reduction is the gain of electrons. The oxidation number of the element decreases.

A useful memory aid is OIL RIG: Oxidation Is Loss, Reduction Is Gain.

Both halves can be written as half-equations, which show only the electrons and the species that change.

Worked example: Worked example: two simple half-equations

Write half-equations for the two changes in

$\text{Zn(s)} + \text{Cu}^{2+}\text{(aq)} \rightarrow \text{Zn}^{2+}\text{(aq)} + \text{Cu(s)}$

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Definition

Oxidising agent and reducing agent

An oxidising agent (oxidant) is a species that accepts electrons and is therefore reduced. It causes the other species to be oxidised.

A reducing agent (reductant) is a species that donates electrons and is therefore oxidised. It causes the other species to be reduced.

The names describe what the substance does to the other substance, not what happens to itself.

Species	What it does	What happens to it
Oxidising agent	takes electrons from something else	is reduced
Reducing agent	gives electrons to something else	is oxidised

Why the oxidising agent is the one reduced

Think of the oxidising agent as a thief of electrons. By taking electrons, it oxidises its victim — but stealing electrons is itself reduction.

So the oxidising agent always appears on the left with a higher oxidation number and ends up reduced.

Common oxidising agents you will meet: $\text{KMnO}_4$ , $\text{K}_2\text{Cr}_2\text{O}_7$ , $\text{Cl}_2$ , concentrated $\text{H}_2\text{SO}_4$ , $\text{H}_2\text{O}_2$ (usually).

Common reducing agents: metals such as $\text{Zn}$ and $\text{Mg}$ , $\text{Fe}^{2+}$ , $\text{SO}_2$ , $\text{H}_2$ , carbon.

Worked example: Worked example: permanganate and iron(II)

For the reaction

$\text{MnO}_4^- + 5\text{Fe}^{2+} + 8\text{H}^+ \rightarrow \text{Mn}^{2+} + 5\text{Fe}^{3+} + 4\text{H}_2\text{O}$

identify what is oxidised, what is reduced, and name the oxidising and reducing agents.

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Quick check

Multiple choiceQuestion 1 · Recall

CalculateQuestion 2 · Recall

Multiple choiceQuestion 3 · Recall

Answer the 3 questions above to continue.