Oxidation and reduction: four ways to see them: WASSCE Chemistry

Iron roofing sheets rusting, charcoal burning, bleach removing stains, a car battery delivering current: all are redox reactions, in which one substance is oxidised while another is reduced. This lesson builds up the four ways of recognising oxidation and reduction, and the oxidising and reducing agents that cause them.

Learning objectives

By the end of this lesson you should be able to:

define oxidation and reduction in terms of oxygen, hydrogen, electrons and oxidation number
explain why oxidation and reduction always occur together
identify oxidising and reducing agents in a reaction
write half-equations for redox reactions
describe tests for oxidising and reducing agents
explain displacement reactions as redox reactions

Four definitions

Definition	Oxidation	Reduction
Oxygen	Gain of oxygen	Loss of oxygen
Hydrogen	Loss of hydrogen	Gain of hydrogen
Electrons	Loss of electrons	Gain of electrons
Oxidation number	Increase in oxidation number	Decrease in oxidation number

A useful memory aid for the electron definition is OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons).

Worked example

Copper(II) oxide heated in hydrogen.

\text{CuO(s)} + \text{H}_2\text{(g)} \rightarrow \text{Cu(s)} + \text{H}_2\text{O(l)}

\begin{gathered}\text{CuO(s)} + \text{H}_2\text{(g)} \\[6pt] \rightarrow \text{Cu(s)} + \text{H}_2\text{O(l)}\end{gathered}

CuO loses oxygen: it is reduced to copper.
H₂ gains oxygen: it is oxidised to water.

Hydrogen sulphide and chlorine.

\text{H}_2\text{S(g)} + \text{Cl}_2\text{(g)} \rightarrow 2\text{HCl(g)} + \text{S(s)}

\begin{gathered}\text{H}_2\text{S(g)} + \text{Cl}_2\text{(g)} \\[6pt] \rightarrow 2\text{HCl(g)} + \text{S(s)}\end{gathered}

H₂S loses hydrogen: it is oxidised to sulphur.
Cl₂ gains hydrogen: it is reduced to HCl.

The electron definition is the most general: it covers reactions with no oxygen or hydrogen at all.

Worked example

Sodium burning in chlorine.

2\text{Na(s)} + \text{Cl}_2\text{(g)} \rightarrow 2\text{NaCl(s)}

\begin{gathered}2\text{Na(s)} + \text{Cl}_2\text{(g)} \\[6pt] \rightarrow 2\text{NaCl(s)}\end{gathered}

Na → Na⁺ + e⁻: sodium loses electrons, so it is oxidised.
Cl₂ + 2e⁻ → 2Cl⁻: chlorine gains electrons, so it is reduced.

Redox: always together

Electrons lost by one substance must be gained by another, so oxidation and reduction always happen together. Such a reaction is called a redox reaction.

Oxidising and reducing agents

Definition

An oxidising agent causes another substance to be oxidised; it is itself reduced (it gains electrons).
A reducing agent causes another substance to be reduced; it is itself oxidised (it loses electrons).

In CuO + H₂ → Cu + H₂O, CuO is the oxidising agent and H₂ is the reducing agent.

Common oxidising agents	Common reducing agents
Oxygen, O₂	Hydrogen, H₂
Chlorine and the halogens	Carbon and carbon(II) oxide
Acidified potassium manganate(VII), KMnO₄	Reactive metals (Na, Mg, Zn, Fe)
Acidified potassium dichromate(VI), K₂Cr₂O₇	Hydrogen sulphide, H₂S
Concentrated nitric and sulphuric acids	Sulphur(IV) oxide, SO₂ (as sulphite)
Hydrogen peroxide, H₂O₂	Iron(II) salts; potassium iodide

Tests

Test for	Reagent	Positive result
A reducing agent	Acidified potassium manganate(VII)	Purple → colourless
A reducing agent	Acidified potassium dichromate(VI)	Orange → green
An oxidising agent	Potassium iodide solution	Colourless → brown (iodine released); turns starch blue-black
An oxidising agent	Hydrogen sulphide	Yellow deposit of sulphur

Half-equations

A redox reaction can be split into two half-equations, one for oxidation and one for reduction. Each must balance in atoms and charge.

Worked example

Zinc in copper(II) sulphate solution. The blue colour fades and brown copper forms on the zinc.

\text{Zn(s)} + \text{CuSO}_4\text{(aq)} \rightarrow \text{ZnSO}_4\text{(aq)} + \text{Cu(s)}

\begin{gathered}\text{Zn(s)} + \text{CuSO}_4\text{(aq)} \\[6pt] \rightarrow \text{ZnSO}_4\text{(aq)} + \text{Cu(s)}\end{gathered}

Oxidation half-equation: Zn(s) → Zn²⁺(aq) + 2e⁻
Reduction half-equation: Cu²⁺(aq) + 2e⁻ → Cu(s)
Ionic equation (adding them; the electrons cancel and SO₄²⁻ is a spectator):

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

\begin{gathered}\text{Zn(s)} + \text{Cu}^{2+}\text{(aq)} \\[6pt] \rightarrow \text{Zn}^{2+}\text{(aq)} + \text{Cu(s)}\end{gathered}

Zinc is the reducing agent; Cu²⁺ is the oxidising agent.

Displacement reactions

A more reactive element displaces a less reactive one from its compounds. These are redox reactions.

Metals: Mg, Zn and Fe displace copper from copper(II) salts, because they are more reactive.
Halogens: a more reactive halogen displaces a less reactive one from its halide:

\text{Cl}_2\text{(aq)} + 2\text{Br}^-\text{(aq)} \rightarrow 2\text{Cl}^-\text{(aq)} + \text{Br}_2\text{(aq)}

\begin{gathered}\text{Cl}_2\text{(aq)} + 2\text{Br}^-\text{(aq)} \\[6pt] \rightarrow 2\text{Cl}^-\text{(aq)} + \text{Br}_2\text{(aq)}\end{gathered}

The solution turns orange-brown as bromine forms. Chlorine is reduced and bromide ions are oxidised.

In Sierra Leone

Household bleach sold in Sierra Leone contains sodium hypochlorite (sodium chlorate(I)), a strong oxidising agent. It removes stains by oxidising coloured compounds into colourless ones, and it kills germs, which is why diluted bleach is used to disinfect water and surfaces. Never mix bleach with acid-based toilet cleaners: the redox reaction releases poisonous chlorine gas.

Common mistake

The oxidising agent is reduced, and the reducing agent is oxidised. Students often say the oxidising agent is oxidised. It does the oxidising, so it must itself gain the electrons.

Key points

Oxidation: gain of oxygen, loss of hydrogen, loss of electrons, increase in oxidation number. Reduction: the opposite.
OIL RIG: Oxidation Is Loss, Reduction Is Gain of electrons.
Oxidation and reduction occur together (redox).
The oxidising agent is reduced; the reducing agent is oxidised.
Tests: reducing agents decolourise KMnO₄ (purple → colourless) and turn K₂Cr₂O₇ green; oxidising agents release iodine from KI.
Displacement reactions of metals and halogens are redox reactions.

WASSCE examination focus

WAEC tests this

Expect to define oxidation and reduction in more than one way, to identify the oxidising and reducing agents in an equation, to write half-equations, and to describe tests for oxidising and reducing agents with colour changes. Displacement of copper by zinc and of bromine by chlorine are standard examples.

Quick check

Define oxidation in terms of electrons.
In Fe₂O₃ + 3CO → 2Fe + 3CO₂, which substance is the reducing agent?
What colour change shows that a substance is a reducing agent when added to acidified potassium dichromate(VI)?
Write the two half-equations for Mg + 2H⁺ → Mg²⁺ + H₂.

Answers:

Oxidation is the loss of electrons.
Carbon(II) oxide, CO: it gains oxygen and reduces the iron(III) oxide.
Orange to green.
Oxidation: Mg → Mg²⁺ + 2e⁻; reduction: 2H⁺ + 2e⁻ → H₂.

Oxidation and reduction: four ways to see them

Learning objectives

Four definitions

Redox: always together

Oxidising and reducing agents

Tests

Half-equations

Displacement reactions

Key points

WASSCE examination focus

Quick check

Keep going with oxidation and reduction reactions

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