Redox Reactions

Redox reactions: oxidation vs reduction, electron transfer, oxidation state changes, and identifying oxidising and reducing agents in reactions.

SEC G3 Pure Chemistry 2027

Reading settings

On this page

Redox can be described using oxygen, hydrogen, electrons or oxidation states. The best method depends on the information given, but oxidation and reduction always occur together.

1. Definition

A redox reaction is a reaction where oxidation and reduction happen at the same time.

2. Key Ideas

Term	How to recognise it
Oxidation	gain oxygen / lose hydrogen / lose electrons / oxidation state increases
Reduction	lose oxygen / gain hydrogen / gain electrons / oxidation state decreases
Oxidising agent	causes oxidation and is reduced (it gains electrons)
Reducing agent	causes reduction and is oxidised (it loses electrons)

If ions or half-equations are given, use electron transfer. If oxidation states are shown or can be assigned, state the increase or decrease. For metal oxides, oxygen gain or loss may be the shortest valid explanation.

3. Detailed Explanations

Oxidation: electrons lost; oxidation state increases.
Reduction: electrons gained; oxidation state decreases.
The reducing agent donates electrons and is oxidised.
The oxidising agent accepts electrons and is reduced.

The agent name describes what it causes the other species to do.

A. Oxidation and reduction in terms of oxygen

If a substance gains oxygen, it is oxidised. If it loses oxygen, it is reduced.

Example: H₂(g) + PbO(s) → H₂O(g) + Pb(s)

H₂ is oxidised (gains oxygen).
PbO is reduced (loses oxygen).
H₂ is the reducing agent (it causes PbO to be reduced and is oxidised itself).
PbO is the oxidising agent (it causes H₂ to be oxidised and is reduced itself).

B. Oxidation and reduction in terms of hydrogen

In suitable redox reactions, loss of hydrogen indicates oxidation and gain of hydrogen indicates reduction. Check what is transferred: accepting a hydrogen ion in an acid–base reaction is not, by itself, reduction.

Example: H₂S(g) + Cl₂(g) → S(s) + 2HCl(g)

H₂S is oxidised (loses hydrogen).
Cl₂ is reduced (gains hydrogen).

C. Oxidation and reduction in terms of electron transfer

If a substance loses electrons, it is oxidised. If it gains electrons, it is reduced.

Example: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Half-equations make the electron loss and gain explicit: Zn(s) → Zn²⁺(aq) + 2e⁻ Cu²⁺(aq) + 2e⁻ → Cu(s)

Zn is oxidised (loses 2e⁻) → reducing agent.
Cu²⁺ is reduced (gains 2e⁻) → oxidising agent.

An electron-accounting model for Zn + Cu²⁺ → Zn²⁺ + Cu. One zinc atom donates the two electrons accepted by one copper(II) ion. The horizontal arrow shows electron transfer; the vertical arrows compare each species before and after. Symbols do not depict a physical arrangement, relative size or electrons floating freely through the solution. Total charge is +2 both before and after.

Check both atoms and charge in each half-equation. For zinc, the right side has net charge + 2-2 = 0, matching the neutral atom. For copper, the left side has net charge + 2-2 = 0, matching neutral copper. Add the half-equations: the two electrons cancel because electron loss equals electron gain.

D. Oxidation state (oxidation number) method

Oxidation state, also called oxidation number, is a bookkeeping value assigned to an atom. For a one-atom ion it equals the ionic charge. In a covalent molecule it is not the atom’s actual charge: the value helps us track changes in bonding without claiming that free ions or electrons are present. In redox:

oxidation state increases → oxidation
oxidation state decreases → reduction

Revise the calculation method in Assigning and Calculating Oxidation States if you need more practice.

Assign oxidation states.
Identify what increases (oxidised) and decreases (reduced).
The species reduced is the oxidising agent; the species oxidised is the reducing agent.

4. Common Mistakes

Saying “oxidising agent causes oxidation so it is oxidised” (wrong): the oxidising agent is reduced.
Calling every reaction involving an oxygen-containing substance redox: neutralisation can involve oxygen without changing oxidation states. Look for paired oxidation and reduction.
Forgetting that oxidation and reduction must occur together (if one happens, the other must happen too).
Mixing up oxidation state with ionic charge: oxidation state is a bookkeeping number, not always a real charge in covalent molecules.
Naming a whole compound when the reacting species is an ion. In the zinc–copper(II) example, Cu²⁺ is the oxidising agent; spectator ions are not involved.

5. Exam Tips

Oxidation = loss of electrons. Reduction = gain of electrons. (“OIL RIG” is fine, but write electrons in your answer.)

Name the reacting species precisely. If an explanation is requested, connect its agent role to electron gain/loss or an oxidation-state change.
If the question gives ions, write half-equations to show electron transfer.
For practical redox evidence, revise the tests for oxidising and reducing agents.

6. Worked Examples

Problem

H₂(g) + PbO(s) → H₂O(g) + Pb(s), identify what is oxidised and reduced, then name both agents.

Study the worked solution

Track oxygen gain
Method
Identify hydrogen as oxidised.
Reason
H₂ gains oxygen to form H₂O.
Working
H₂: oxidation; therefore H₂ is the reducing agent.
Track oxygen loss
Method
Identify lead(II) oxide as reduced.
Reason
PbO loses oxygen to form Pb.
Working
PbO: reduction; therefore PbO is the oxidising agent.
Check simultaneous changes
Method
Pair the two transfers.
Reason
Oxygen gained by hydrogen comes from lead(II) oxide, so oxidation and reduction occur together.
Working
Reducing agent H₂; oxidising agent PbO.

Problem

For

Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s), write both half-equations and identify the agents.

Place electrons and assign roles

Hints

Hint 1: charges

A neutral zinc atom must lose two electrons to become

Zn²⁺.

Hint 2: agent rule

The electron acceptor is reduced and is the oxidising agent.

View solution step by step

Write zinc oxidation
Method
Place two electrons on the product side.
Reason
Zinc loses two electrons as its oxidation state rises from 0 to +2.
Working
Zn(s) → Zn²⁺(aq) + 2e⁻; Zn is the reducing agent.
Write copper-ion reduction
Method
Place two electrons on the reactant side.
Reason
Cu²⁺ gains two electrons to become neutral copper.
Working
Cu²⁺(aq) + 2e⁻ → Cu(s); Cu²⁺ is the oxidising agent.

Learner claim

For

Zn + Cu²⁺ → Zn²⁺ + Cu, a student writes, “

Cu²⁺ is oxidised because it loses electrons.” Correct the species change and agent role.

View solution step by step

Use the charge change
Method
Give Cu²⁺ two electrons.
Reason
The +2 ion becomes neutral copper only by gaining two negative charges.
Working
Cu²⁺ + 2e⁻ → Cu.
Name process and agent
Method
Call the change reduction and Cu²⁺ the oxidising agent.
Reason
Electron gain is reduction; an oxidising agent accepts electrons from the species it oxidises.
Working
Cu²⁺ is reduced and oxidises zinc.

Examination question

2KI(aq) + Cl₂(aq) → 2KCl(aq) + I₂(aq), identify what is oxidised and reduced and name both agents using oxidation states. [4 marks]

Write the complete oxidation-state analysis

Write your answer

View solution step by step

Iodide oxidation
1 mark
Method
Track iodine from -1 in I⁻ to 0 in I₂.
Reason
An increase in oxidation state is oxidation.
Working
-1 → 0: I⁻ is oxidised.
Chlorine reduction
1 mark
Method
Track chlorine from 0 in Cl₂ to -1 in Cl⁻.
Reason
A decrease in oxidation state is reduction.
Working
0 → -1: Cl₂ is reduced.
Reducing agent
1 mark
Method
Name I⁻ in KI as the reducing agent.
Reason
The species that is oxidised causes the paired reduction.
Working
I⁻ in KI is the reducing agent because it is oxidised.
Oxidising agent
1 mark
Working
Cl₂ is the oxidising agent because it is reduced.

Self-mark with the mark scheme

Compare your response with each mark point. Select a point only when your response contains that evidence.

Self-mark both oxidation-state changes and both agent identities.

Mark points

1 mark: Iodine −1 to 0; iodide oxidised.
1 mark: Chlorine 0 to −1; chlorine reduced.
1 mark: Identifies I−/KI as reducing agent.
1 mark: Identifies Cl2 as oxidising agent.

Classification transfer

Which reaction is redox? Justify both classifications. (A)

NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l); (B)

Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g).

Look for paired electron changes

Reaction A

Not redoxRedox

Reaction B

RedoxNot redox

Hints

Hint 1: do not use reaction label alone

Neutralisation or gas formation does not by itself decide redox; test whether oxidation states change.

Hint 2: track magnesium and hydrogen

In B, compare Mg before and after, then H in

H⁺ and

H₂.

View solution step by step

Test neutralisation
Method
Classify A as not redox.
Reason
No oxidation state changes and no electron transfer occur; H⁺ and OH⁻ form water.
Working
Reaction A: acid–base neutralisation, not redox.
Track magnesium
Method
Oxidise magnesium from 0 to +2.
Reason
Magnesium loses two electrons when forming Mg²⁺.
Working
Mg → Mg²⁺ + 2e⁻.
Track hydrogen and classify
Method
Reduce H⁺ to H₂ and classify B as redox.
Reason
Hydrogen gains the electrons magnesium loses, so the two changes occur together.
Working
2H⁺ + 2e⁻ → H₂; Reaction B is redox.

Next step

Try Agent logic independently.

7. Mind Stretchers

Mind stretcher 1: Agent logic (no guessing)Extension

Question: In a redox reaction, the oxidising agent is always reduced. Explain why this must be true using electrons.

Show Answer

An oxidising agent causes oxidation by taking electrons from another species. If it takes electrons, it gains electrons, so it is reduced. The reducing agent donates those electrons and is oxidised. Electron loss and gain balance.

Mind stretcher 2: Two methods, same conclusionExtension

Try independently: For 2Fe²⁺(aq) + Cl₂(aq) → 2Fe³⁺(aq) + 2Cl⁻(aq), identify both agents. Write the two half-equations and explain why two iron(II) ions are needed for each chlorine molecule.

Show answer and reasoning

Each iron(II) ion loses one electron: Fe²⁺ → Fe³⁺ + e⁻. Iron changes from +2 to +3, so Fe²⁺ is oxidised and is the reducing agent.

One chlorine molecule gains two electrons: Cl₂ + 2e⁻ → 2Cl⁻. Chlorine changes from 0 to −1, so Cl₂ is reduced and is the oxidising agent. Two iron(II) ions supply the two electrons accepted by one chlorine molecule.

Try independently: In NH₃(aq) + H⁺(aq) → NH₄⁺(aq), ammonia gains a hydrogen ion. Does that make this reduction? Use the usual values H = +1 and the total charge to check nitrogen.

Show answer and reasoning

No. Nitrogen is −3 in ammonia: x + 3(+1) = 0. It is also −3 in ammonium: x + 4(+1) = +1. Hydrogen remains +1. This is an acid–base reaction with no oxidation-state changes, so it is not redox. Gaining a proton is different from gaining an electron.

Practise and check

Test yourself on oxidation vs reduction, agents, half-equations, and oxidation state changes.

Open the Redox Chemistry topic check

Syllabus and review details

SEC G3 Pure Chemistry 2027 · 2027
Content structure and subject content, PDF pages 9–24

Last reviewed: 2026-08-29

Redox Reactions

1. Definition

2. Key Ideas

3. Detailed Explanations

A. Oxidation and reduction in terms of oxygen

B. Oxidation and reduction in terms of hydrogen

C. Oxidation and reduction in terms of electron transfer

D. Oxidation state (oxidation number) method

4. Common Mistakes

5. Exam Tips

6. Worked Examples

Oxygen definition (and agents)

Track oxygen gain

Track oxygen loss

Check simultaneous changes

Electron transfer (half-equations)

Write zinc oxidation

Write copper-ion reduction

Spot the wrong sentence (and fix it)

Use the charge change

Name process and agent

Oxidation state method (common exam type)

Iodide oxidation

Chlorine reduction

Reducing agent

Oxidising agent

Is it redox?

Test neutralisation

Track magnesium

Track hydrogen and classify

7. Mind Stretchers

Mind stretcher 1: Agent logic (no guessing)Extension

Mind stretcher 2: Two methods, same conclusionExtension

Practise and check