OxfordAQA International GCSE Chemistry: Redox reactions (9202)

This guide teaches section 3.8.4 Redox reactions from the OxfordAQA International GCSE Chemistry (9202) specification. The version followed here is 6.2, for exams May/June 2018 onwards. In all, 3.8.4 has three statements, labelled a, b and c, and this page covers all of them. The qualification has no tiers, so every idea below is for you whatever grade you want, and the specification says either paper may assess content from any part of it.

There is no required practical in section 3.8.4. The redox ideas here do turn up inside other practicals, such as the copper sulfate electrolysis in 3.3.2, so expect to use them there too.

The OxfordAQA chemistry hub links all the topics in this course, and the course checklist can be printed so you can tick statements off. If you want to find your weak spots first, take a free diagnostic. Afterwards, shorten what you learned using the Redox reactions revision notes and test yourself with the Redox reactions practice questions.

Statements in this unit

Spec ref	Skill
3.8.4 a	Say that a substance gaining oxygen is oxidised and one losing oxygen is reduced
3.8.4 b	Say that a particle losing electrons is oxidised and one gaining electrons is reduced; apply both definitions to reactions across the specification
3.8.4 c	Recognise a redox reaction as one where oxidation and reduction happen at the same time, as in Fe₂O₃ + 3CO → 2Fe + 3CO₂

The guidance under statement b matters most. It asks you to describe reactions found elsewhere in the specification in terms of oxidation and reduction. So this unit is really a toolkit you apply to metals, electrolysis, the halogens and fuels.

3.8.4 a: Oxidation and reduction as oxygen transfer

The first definition is the older one, and it is easy to use whenever oxygen appears in an equation.

• Oxidation is the gain of oxygen by a substance.
• Reduction is the loss of oxygen from a substance.

To use it, look at each reactant in turn and ask: does this substance end up with more oxygen, or less?

Worked example 1: copper(II) oxide and carbon monoxide

Carbon monoxide passed over hot copper(II) oxide gives copper metal:

CuO + CO → Cu + CO₂

• Copper(II) oxide starts with one oxygen atom and ends as copper with none. It has lost oxygen, so CuO is reduced.
• Carbon monoxide has one oxygen atom and becomes carbon dioxide with two. It has gained oxygen, so CO is oxidised.

Both processes happen in the same reaction. That pairing is the subject of statement c.

Worked example 2: burning a fuel

Butane is a hydrocarbon used in camping gas. When it burns completely:

2C₄H₁₀ + 13O₂ → 8CO₂ + 10H₂O

Which substance is oxidised? The carbon in butane ends up joined to oxygen in CO₂, and the hydrogen ends up joined to oxygen in H₂O. So both the carbon and the hydrogen in the fuel are oxidised. This matches statement 3.10.1.2 e, which says the carbon and hydrogen in hydrocarbon fuels are oxidised during combustion.

The same thinking covers metals burning in air. In 4Al + 3O₂ → 2Al₂O₃, aluminium gains oxygen, so aluminium is oxidised.

Where the oxygen definition runs out

Many reactions involve no oxygen at all. Iron burning in chlorine, or chlorine water reacting with potassium iodide, cannot be judged by counting oxygen atoms. For these you need statement b.

3.8.4 b: Oxidation and reduction as electron transfer

The second definition works for any reaction in which electrons move from one particle to another.

• Oxidation means a particle loses electrons.
• Reduction means a particle gains electrons.

To remember which is which, use OIL RIG: Oxidation Is Loss and Reduction Is Gain.

Writing half equations

A half equation shows one particle losing or gaining electrons. Electrons are written e⁻.

• If electrons appear on the right, they have been given out, so the change is oxidation: Mg → Mg²⁺ + 2e⁻
• If electrons appear on the left, they have been taken in, so the change is reduction: Ag⁺ + e⁻ → Ag

Check every half equation two ways. The atoms must balance, and the total charge on each side must be the same. In Mg → Mg²⁺ + 2e⁻ the left has charge 0, and the right has (+2) + (−2) = 0.

The specification (3.3.2 f) also shows a second style for an oxidation half equation, 2Cl⁻ − 2e⁻ → Cl₂, with the electrons subtracted on the left. It means the same as 2Cl⁻ → Cl₂ + 2e⁻.

Worked example 3: iron burning in chlorine

Hot iron wool glows in chlorine gas and forms iron(III) chloride, an ionic solid containing Fe³⁺ and Cl⁻ ions.

2Fe + 3Cl₂ → 2FeCl₃

Step 1. Find what happens to the iron. Each Fe atom becomes an Fe³⁺ ion, so it loses three electrons: Fe → Fe³⁺ + 3e⁻ (oxidation)

Step 2. Find what happens to the chlorine. Each Cl₂ molecule becomes two Cl⁻ ions, so it gains two electrons: Cl₂ + 2e⁻ → 2Cl⁻ (reduction)

Step 3. Match the electrons. The iron half equation gives 3 electrons and the chlorine half equation takes 2. The lowest common number is 6, so double the first and triple the second: 2Fe → 2Fe³⁺ + 6e⁻ 3Cl₂ + 6e⁻ → 6Cl⁻

Step 4. Conclude. Iron is oxidised and chlorine is reduced. Six electrons pass from iron to chlorine for every two iron atoms. Notice that no oxygen is involved, yet the reaction is clearly oxidation and reduction.

Worked example 4: a metal displacement

The specification asks you to describe displacement reactions in terms of oxidation and reduction and to write their ionic equations. A strip of magnesium is placed in silver nitrate solution, and grey crystals of silver grow on it.

Full equation: Mg + 2AgNO₃ → Mg(NO₃)₂ + 2Ag

Step 1. Remove spectator ions. The nitrate ions, NO₃⁻, are in solution before and after. They take no part, so leave them out.

Step 2. Write the ionic equation. Mg + 2Ag⁺ → Mg²⁺ + 2Ag

Step 3. Split into half equations. Mg → Mg²⁺ + 2e⁻ (magnesium loses electrons, so it is oxidised) 2Ag⁺ + 2e⁻ → 2Ag (silver ions gain electrons, so they are reduced)

The more reactive metal is always the one oxidised in a displacement. It gives its electrons to the ions of the less reactive metal. The scrap-iron method for copper in 3.3.1.1 works the same way: Fe + Cu²⁺ → Fe²⁺ + Cu, with iron oxidised and copper ions reduced.

Worked example 5: a halogen displacement

Chlorine water is added to colourless potassium iodide solution and the mixture turns brown as iodine forms.

Full equation: Cl₂ + 2KI → 2KCl + I₂. Without the potassium ions: Cl₂ + 2I⁻ → 2Cl⁻ + I₂

• 2I⁻ → I₂ + 2e⁻: iodide ions lose electrons, so they are oxidised.
• Cl₂ + 2e⁻ → 2Cl⁻: chlorine gains electrons, so it is reduced.

Chlorine is higher in Group 7 than iodine. Its outer electrons are in a lower energy level, so it gains electrons more easily, which is why it can take them from iodide ions (statement 3.7.1 f).

Electrolysis in electron terms

Section 3.3.2 already defines oxidation and reduction as loss and gain of electrons. At the cathode, positive ions gain electrons, so reduction always happens at the cathode. At the anode, negative ions lose electrons, so oxidation always happens at the anode.

For molten calcium bromide, CaBr₂:

• Cathode: Ca²⁺ + 2e⁻ → Ca (reduction)
• Anode: 2Br⁻ → Br₂ + 2e⁻ (oxidation)

3.8.4 c: Redox reactions

When oxidation and reduction happen at the same time, the reaction is called a redox reaction. The word joins REDuction and OXidation.

They must happen together for a simple reason. If one substance gains oxygen, that oxygen came from somewhere. If one particle loses electrons, another particle must gain them. Electrons lost always equal electrons gained.

The specification’s example

Fe₂O₃ + 3CO → 2Fe + 3CO₂

This is the main reaction that makes iron in the blast furnace (3.3.1.1 d).

• Iron(III) oxide → iron: oxygen is lost, so this half of the reaction is reduction.
• Carbon monoxide → carbon dioxide: oxygen is gained, so this half is oxidation.

In the specification this is drawn with a “Reduction” arrow joining Fe₂O₃ to Fe above the equation and an “Oxidation” arrow joining CO to CO₂ below it. You may be asked to label an equation in the same way.

Worked example 6: is it redox or not?

Decide which of these reactions are redox.

Reaction	Oxygen or electrons transferred?	Redox?
Ca + 2HCl → CaCl₂ + H₂	Ca → Ca²⁺ + 2e⁻ and 2H⁺ + 2e⁻ → H₂	Yes
HCl + NaOH → NaCl + H₂O	Ions swap partners; no particle gains or loses electrons	No
CaCO₃ → CaO + CO₂	Ca²⁺ stays Ca²⁺; no oxygen is gained by any substance	No
2K + 2H₂O → 2KOH + H₂	K → K⁺ + e⁻; potassium is oxidised	Yes
C₂H₆O + O₂ → C₂H₄O₂ + H₂O	Ethanol gains oxygen to become ethanoic acid	Yes

Neutralisation and thermal decomposition of carbonates are common traps. They are reactions, but not redox. The last row is the oxidation of ethanol to ethanoic acid from statement 3.10.3.1 c.

Redox across the specification

Statement b expects you to apply the definitions to reactions you meet elsewhere. This table collects them.

Where in the specification	Reaction type	What is oxidised	What is reduced
3.3.1.1 a	Metal with water or dilute acid	Metal atoms (lose electrons)	Hydrogen ions or water (hydrogen forms)
3.3.1.1 b, h	Metal displacement, including scrap iron	The more reactive metal	Ions of the less reactive metal
3.3.1.1 d	Carbon or carbon monoxide with a metal oxide	Carbon or CO (gains oxygen)	The metal oxide (loses oxygen)
3.3.2	Electrolysis	Negative ions at the anode	Positive ions at the cathode
3.7.1 a	Group 1 metal with water or a halogen	Group 1 metal	Water or the halogen
3.7.1 e	Halogen displacement	Halide ions	The more reactive halogen
3.10.1.2 e	Combustion of hydrocarbons	Carbon and hydrogen in the fuel	Oxygen
3.10.3.1 c	Ethanol to ethanoic acid	Ethanol	The oxidising agent or oxygen

For carbon reduction of metal oxides, the specification limits your knowledge of reduction to removal of oxygen. Use the oxygen definition there.

Reactivity links to oxidation too. A more reactive metal loses electrons more easily, so it is oxidised more easily. Going down Group 1, the outer electron sits in a higher energy level each time, so the metal gives it up (is oxidised) more readily and reactivity rises.

Common errors

• Reversing OIL RIG. Losing electrons is oxidation, even though the particle becomes “more positive”.
• Putting electrons on the wrong side. In a reduction the electrons go on the left, as in Pb²⁺ + 2e⁻ → Pb.
• Forgetting that Cl₂, Br₂ and I₂ are molecules. You need 2Cl⁻ and 2e⁻, not Cl⁻ and e⁻.
• Saying “copper sulfate is reduced” in a displacement. It is the copper ions that gain electrons; the sulfate ions are spectators.
• Labelling every reaction as redox. Check for oxygen or electron transfer first.
• Unbalanced charge in a half equation. Add the charges on each side before you move on.
• Saying oxidation happens at the cathode in electrolysis. It is the anode.

Where to go next

• Redox reactions revision notes for quick recall and a self-test
• Redox reactions practice questions with full answers
• Metals study guide for the reactivity series and extraction
• Electrolysis study guide for more half equations at electrodes
• Free diagnostics to check what to revise next

Official syllabus

International GCSE Chemistry (9202) specification published by OxfordAQA; Version 6.2, for exams May/June 2018 onwards. Unit covered: 3.8.4 Redox reactions, statements a to c.