What you'll learn
This revision guide covers oxidation and reduction (redox) reactions as specified in the AQA GCSE Chemistry specification. You'll learn to identify oxidation and reduction in terms of both oxygen transfer and electron transfer, write half equations, and understand the role of oxidising and reducing agents in chemical reactions.
Key terms and definitions
Oxidation — the gain of oxygen, loss of hydrogen, or loss of electrons by a substance
Reduction — the loss of oxygen, gain of hydrogen, or gain of electrons by a substance
Redox reaction — a chemical reaction in which both oxidation and reduction occur simultaneously
Oxidising agent — a substance that oxidises another substance by accepting electrons (the oxidising agent itself is reduced)
Reducing agent — a substance that reduces another substance by donating electrons (the reducing agent itself is oxidised)
Half equation — an equation showing either the oxidation or reduction process separately, including electrons
Displacement reaction — a reaction where a more reactive element displaces a less reactive element from its compound
OIL RIG — a mnemonic for remembering electron transfer: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons)
Core concepts
Oxidation and reduction in terms of oxygen
The original definitions of oxidation and reduction were based on oxygen transfer. These definitions remain valid and are frequently tested at GCSE level.
Oxidation occurs when a substance gains oxygen. For example:
- When magnesium burns in air: 2Mg + O₂ → 2MgO
- The magnesium has been oxidised to magnesium oxide
Reduction occurs when a substance loses oxygen. For example:
- When iron oxide is reduced using carbon: 2Fe₂O₃ + 3C → 4Fe + 3CO₂
- The iron oxide has been reduced to iron
- The carbon has been oxidised to carbon dioxide
In many reactions, one substance is oxidised whilst another is reduced. The substance that gains oxygen is oxidised, and the substance that provides the oxygen is the oxidising agent.
Oxidation and reduction in terms of electrons
The modern definition of oxidation and reduction focuses on electron transfer. This is more comprehensive as it applies to reactions that don't involve oxygen.
OIL RIG is the key mnemonic:
- Oxidation Is Loss (of electrons)
- Reduction Is Gain (of electrons)
When a metal reacts with a non-metal:
- The metal atoms lose electrons and are oxidised
- The non-metal atoms gain electrons and are reduced
For example, when sodium reacts with chlorine:
- Sodium: Na → Na⁺ + e⁻ (oxidation — loss of electrons)
- Chlorine: Cl₂ + 2e⁻ → 2Cl⁻ (reduction — gain of electrons)
Writing half equations
Half equations show either oxidation or reduction separately, with electrons shown as reactants or products. This is an essential skill for GCSE Chemistry.
Rules for writing half equations:
- Write the formula of the substance before and after the change
- Balance the atoms (except oxygen and hydrogen initially)
- Add electrons to balance the charges
- Check that both atoms and charges are balanced
Example — oxidation half equations:
- Magnesium: Mg → Mg²⁺ + 2e⁻
- Iron(II): Fe²⁺ → Fe³⁺ + e⁻
- Aluminium: Al → Al³⁺ + 3e⁻
Example — reduction half equations:
- Oxygen: O₂ + 4e⁻ → 2O²⁻
- Copper: Cu²⁺ + 2e⁻ → Cu
- Chlorine: Cl₂ + 2e⁻ → 2Cl⁻
The electrons lost in oxidation must equal the electrons gained in reduction. When combining half equations, you may need to multiply one or both equations to balance the electrons.
Identifying oxidation and reduction in reactions
In any redox reaction, you must identify which substance is oxidised and which is reduced.
Method 1 — Using oxygen transfer: Look for which substance gains oxygen (oxidised) and which loses oxygen (reduced).
Method 2 — Using electron transfer: Look at the charges or oxidation states:
- If the charge becomes more positive (or less negative), oxidation has occurred
- If the charge becomes more negative (or less positive), reduction has occurred
Example: Zn + Cu²⁺ → Zn²⁺ + Cu
- Zinc goes from 0 to +2 charge: loses electrons, so oxidised
- Copper goes from +2 to 0 charge: gains electrons, so reduced
Oxidising agents and reducing agents
Understanding the role of oxidising and reducing agents is crucial for explaining redox reactions.
An oxidising agent:
- Accepts electrons from another substance
- Causes oxidation in another substance
- Is itself reduced in the process
- Examples: oxygen, chlorine, copper(II) ions, acidified potassium manganate(VII)
A reducing agent:
- Donates electrons to another substance
- Causes reduction in another substance
- Is itself oxidised in the process
- Examples: hydrogen, carbon, carbon monoxide, metals, sulfur dioxide
In the reaction: Zn + Cu²⁺ → Zn²⁺ + Cu
- Zinc is the reducing agent (it reduces copper ions)
- Copper ions are the oxidising agent (they oxidise zinc)
Displacement reactions as redox reactions
Displacement reactions occur when a more reactive element displaces a less reactive element from a compound. All displacement reactions are redox reactions.
Metal displacement reactions:
When zinc is added to copper sulfate solution: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
Or more simply: Zn + Cu²⁺ → Zn²⁺ + Cu
Half equations:
- Oxidation: Zn → Zn²⁺ + 2e⁻
- Reduction: Cu²⁺ + 2e⁻ → Cu
Zinc is more reactive than copper, so it displaces copper from the solution. Zinc is oxidised (loses electrons) and copper ions are reduced (gain electrons).
Halogen displacement reactions:
Chlorine is more reactive than bromine, so chlorine displaces bromine from bromide solutions: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂
Half equations:
- Oxidation: 2Br⁻ → Br₂ + 2e⁻
- Reduction: Cl₂ + 2e⁻ → 2Cl⁻
Bromide ions are oxidised and chlorine is reduced.
Extraction of metals by reduction
The extraction of metals from their ores often involves reduction. Carbon is commonly used as a reducing agent to extract metals less reactive than carbon.
Example — iron extraction:
Iron oxide is reduced in the blast furnace: Fe₂O₃ + 3CO → 2Fe + 3CO₂
Analysis:
- Iron oxide loses oxygen, so iron is reduced
- Carbon monoxide gains oxygen, so carbon is oxidised
- Carbon monoxide is the reducing agent
More reactive metals (such as aluminium) cannot be extracted by carbon. These require electrolysis, where reduction occurs at the cathode: Al³⁺ + 3e⁻ → Al
Worked examples
Example 1: Identifying oxidation and reduction
Question: Aluminium reacts with iron oxide in the thermit reaction: 2Al + Fe₂O₃ → Al₂O₃ + 2Fe
(a) Identify which substance has been reduced. Explain your answer in terms of oxygen. [2 marks]
(b) Write a half equation to show what happens to aluminium atoms in this reaction. [2 marks]
Solution:
(a) Iron (oxide) / Fe₂O₃ has been reduced [1 mark] because it has lost oxygen / oxygen has been removed [1 mark]
(b) Al → Al³⁺ + 3e⁻ [2 marks] Award 1 mark for Al³⁺ shown as product Award 2nd mark for correct number of electrons (3e⁻) on right side
Example 2: Displacement reactions
Question: A student adds magnesium ribbon to copper sulfate solution. The blue solution fades and a brown solid forms.
(a) Write the ionic equation for this reaction. [2 marks]
(b) Identify which substance is the oxidising agent. Explain your answer in terms of electrons. [2 marks]
Solution:
(a) Mg + Cu²⁺ → Mg²⁺ + Cu [2 marks] Award 1 mark for correct reactants and products Award 2nd mark for balanced equation with correct charges
(b) Copper ions / Cu²⁺ is the oxidising agent [1 mark] because it accepts/gains electrons (from magnesium) / causes oxidation [1 mark]
Example 3: Half equations
Question: Iron can be extracted from iron(III) oxide using carbon monoxide.
Fe₂O₃ + 3CO → 2Fe + 3CO₂
(a) Complete the half equation for the reduction of iron ions: Fe³⁺ + _____ → _____ [2 marks]
(b) Explain why carbon monoxide is described as a reducing agent in this reaction. [2 marks]
Solution:
(a) Fe³⁺ + 3e⁻ → Fe [2 marks] Award 1 mark for correct number of electrons (3e⁻) Award 1 mark for Fe as product
(b) Carbon monoxide removes oxygen from / reduces iron oxide [1 mark] Carbon monoxide is itself oxidised (in the process) [1 mark]
Alternative: Carbon monoxide donates electrons / causes reduction [1 mark]
Common mistakes and how to avoid them
Confusing OIL RIG — Remember that oxidation is LOSS of electrons, not gain. The mnemonic helps: Oxidation Is Loss. Students often reverse this, particularly under exam pressure.
Forgetting to balance charges in half equations — Always check that the total charge is the same on both sides. For example, in Mg → Mg²⁺ + 2e⁻, the left side has 0 charge and the right side has +2 + (2×-1) = 0 charge.
Mixing up oxidising and reducing agents — The oxidising agent is reduced (it accepts electrons), while the reducing agent is oxidised (it donates electrons). This seems backwards but makes sense when you think about the effect each has on the other substance.
Writing molecular formulas instead of ionic — In displacement reactions, focus on the ions changing, not the spectator ions. Write Cu²⁺ + 2e⁻ → Cu, not CuSO₄ + 2e⁻ → Cu.
Incorrect electron numbers in half equations — The number of electrons must equal the change in charge. If Al goes from 0 to Al³⁺, it must lose 3 electrons, not 1 or 2.
Not using state symbols when required — Some exam questions specifically ask for state symbols. Remember (s), (l), (g), and (aq).
Exam technique for "Chemical changes: oxidation and reduction in terms of electrons and oxygen"
Command word "Identify" — State which substance is oxidised/reduced. One-word answers with the substance name are often sufficient for 1 mark, but the question may require an explanation for the second mark.
Command word "Explain" — You must give a reason. For oxidation/reduction, state what happens in terms of oxygen OR electrons, depending on the question context. "Magnesium is oxidised because it loses electrons" scores marks; "Magnesium is oxidised" alone may not.
Half equations — These are worth 2 marks typically. You get 1 mark for the correct species and 1 mark for balancing (including electrons). Show your working by writing electrons clearly as e⁻.
Mark allocation guide — In a 4-mark question on redox, expect to: identify both oxidation and reduction (2 marks) and explain each in terms of oxygen or electrons (2 marks). Plan your answer to cover all mark points.
Quick revision summary
Oxidation is the gain of oxygen or loss of electrons (OIL). Reduction is the loss of oxygen or gain of electrons (RIG). In redox reactions, oxidation and reduction occur simultaneously. Oxidising agents accept electrons and are reduced; reducing agents donate electrons and are oxidised. Half equations show electron transfer in oxidation or reduction processes separately. Displacement reactions are redox reactions where a more reactive element displaces a less reactive one. Metal extraction often involves reduction of metal ores using carbon or electrolysis.