AQA A-Level Chemistry: Oxidation, reduction and redox equations (7405) – Revision Notes

These revision notes cover section 3.1.7, Oxidation, reduction and redox equations, of the AQA AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for AS and A-level exams from June 2016 onwards. The whole section is AS content and is also part of the full A-level; none of it is marked “A-level only”. The specification places it in Paper 1 at both AS and A-level, and A-level Paper 3 can test any content.

For full explanations and step-by-step worked examples, read the study guide. When you are ready to test yourself under exam conditions, use the practice questions. The AQA A-Level Chemistry hub and printable checklist list every other section.

3.1.7 Definitions

Term	Definition (specification wording)
Oxidation	the process of electron loss
Reduction	the process of electron gain
Oxidising agent	an electron acceptor (it is reduced)
Reducing agent	an electron donor (it is oxidised)
Redox reaction	electrons are transferred from the reducing agent to the oxidising agent

In terms of oxidation state:

• oxidation state goes up → the element is oxidised
• oxidation state goes down → the element is reduced

Oxidation state rules

Species or element	Oxidation state
Uncombined element	0
Neutral compound (sum)	0
Ion (sum)	charge on the ion
Group 1 / Group 2 / Al in compounds	+1 / +2 / +3
F in compounds	−1
H	+1 (−1 in metal hydrides, e.g. NaH, CaH₂)
O	−2 (−1 in peroxides, e.g. H₂O₂; +2 in OF₂)
Cl	−1 (positive when bonded to O or F)

Roman numerals in names = oxidation state: sulfate(VI), manganate(VII), dichromate(VI), chlorate(I), iron(II), iron(III).

Writing convention: oxidation state +2; ionic charge 2+.

Oxidation states worth recognising on sight

These all follow from the rules above. Knowing them saves time, but in the exam you should still be able to show the working.

Element	Species and oxidation state
Manganese	MnO₄⁻ +7, MnO₂ +4, Mn²⁺ +2
Chromium	Cr₂O₇²⁻ +6, CrO₄²⁻ +6, Cr³⁺ +3
Sulfur	SO₄²⁻ and H₂SO₄ +6, SO₂ +4, S₈ 0, H₂S −2
Nitrogen	NO₃⁻ +5, NO₂ +4, NO +2, NH₃ and NH₄⁺ −3
Chlorine	ClO₃⁻ +5, ClO⁻ +1, Cl₂ 0, Cl⁻ −1
Iodine	IO₃⁻ +5, I₂ 0, I⁻ −1
Oxygen	H₂O −2, H₂O₂ −1, O₂ 0
Iron	Fe₂O₃ and Fe³⁺ +3, Fe²⁺ +2

Method in steps

Working out an oxidation state

1. Write x for the unknown element.
2. Fill in the known values from the rules.
3. Set the sum equal to 0 (compound) or the ion’s charge.
4. Solve. A fraction (such as +2.5 for S in S₄O₆²⁻) is an average and is acceptable.

Small reminder: Cl in ClO₃⁻ → x + 3(−2) = −1, so x = +5.

Deciding what is oxidised and reduced in an equation

1. Write the oxidation state of every element on both sides.
2. Cross out any element whose oxidation state does not change (spectator ions such as Na⁺, K⁺, SO₄²⁻ often fall here).
3. Up = oxidised (that species is the reducing agent). Down = reduced (that species is the oxidising agent).
4. Name the species, not just the element, when asked for an agent.

Small reminder: Zn + 2Fe³⁺ → Zn²⁺ + 2Fe²⁺. Zn goes 0 → +2 (oxidised, reducing agent). Fe goes +3 → +2 (reduced, so Fe³⁺ is the oxidising agent).

Writing a half-equation in acid

1. Balance the element that changes oxidation state.
2. Balance O with H₂O.
3. Balance H with H⁺.
4. Balance charge with e⁻.
5. Check electrons = change in oxidation state × atoms changing.

Small reminder:

NO₃⁻ + 2H⁺ + e⁻ → NO₂ + H₂O

N goes from +5 in NO₃⁻ to +4 in NO₂: one electron gained. Reduction.

Combining two half-equations

1. Scale so electrons lost = electrons gained.
2. Add.
3. Cancel e⁻, then any H⁺ and H₂O on both sides.
4. Check atoms and total charge.

Small reminder, copper metal with silver ions:

Cu → Cu²⁺ + 2e⁻
Ag⁺ + e⁻ → Ag       (× 2)
2Ag⁺ + Cu → 2Ag + Cu²⁺

Final checks on any redox equation

• Atoms of every element balance.
• Total charge on the left equals total charge on the right.
• Electrons appear only in half-equations, never in the overall equation.
• Electrons lost in the oxidation equal electrons gained in the reduction.
• H⁺ and H₂O appear on one side only after cancelling.
• Add state symbols only if the question asks for them.

Command words for this section

• State / Give: the answer alone, such as “+6” or “Fe³⁺”. No working needed, but a sign is.
• Deduce: use the rules or the equation to reach the answer; show the step if it is not obvious.
• Show that: every step must be written, because the answer is given.
• Write a half-equation: species, H⁺, H₂O and e⁻ all balanced.
• Explain: link the change to electrons, for example “Fe²⁺ loses an electron, so its oxidation state rises from +2 to +3”.
• Identify: name or give the formula of a single species.

Must-know distinctions

Do not confuse	With
Oxidising agent (reduced)	Species oxidised (the reducing agent)
Oxidation state (+3)	Ionic charge (3+)
Half-equation (contains e⁻)	Overall equation (no e⁻)
Electrons on the left = reduction	Electrons on the right = oxidation
Spectator ion (no change)	Reacting ion (oxidation state changes)
H in H₂O, HCl (+1)	H in metal hydrides (−1)
O in oxides (−2)	O in peroxides (−1)

Quick self-test

Cover the answers, then try each one.

1. State the oxidation state of manganese in K₂MnO₄.
2. State the oxidation state of sulfur in H₂S.
3. State the oxidation state of chlorine in Cl₂O₇.
4. State the oxidation state of oxygen in Na₂O₂, and give a reason.
5. State the oxidation state of nitrogen in NH₂OH.
6. Iron(II) chloride is converted into iron(III) chloride. Has iron been oxidised or reduced? Explain in terms of electrons.
7. Write a half-equation for aluminium forming aluminium ions.
8. Aluminium reacts with copper(II) ions: 2Al + 3Cu²⁺ → 2Al³⁺ + 3Cu. Identify the reducing agent.
9. Write a half-equation for chlorate(I) ions, ClO⁻, being reduced to chloride ions in acidic solution.
10. Write a half-equation for hydrogen peroxide being oxidised to oxygen.
11. Hydrogen peroxide is reduced to water in acid: H₂O₂ + 2H⁺ + 2e⁻ → 2H₂O. Combine this with the oxidation of Fe²⁺ to Fe³⁺ to give the overall equation.
12. In the reaction 2Fe³⁺ + 2I⁻ → 2Fe²⁺ + I₂, identify the oxidising agent and write the half-equation for the oxidation.

Answers

1. +6. 2(+1) + x + 4(−2) = 0, so x = +6.
2. −2. 2(+1) + x = 0.
3. +7. 2x + 7(−2) = 0, so 2x = +14.
4. −1. Na₂O₂ is a peroxide; Na is +1, so 2(+1) + 2x = 0 gives x = −1.
5. −1. x + 2(+1) + (−2) + (+1) = 0, so x = −1.
6. Oxidised. Iron goes from +2 to +3, so each Fe²⁺ ion loses one electron.
7. Al → Al³⁺ + 3e⁻
8. Al (aluminium atoms donate electrons; Al goes from 0 to +3).
9. ClO⁻ + 2H⁺ + 2e⁻ → Cl⁻ + H₂O (Cl goes from +1 to −1, two electrons gained).
10. H₂O₂ → O₂ + 2H⁺ + 2e⁻ (O goes from −1 to 0; two O atoms, two electrons lost).
11. Multiply Fe²⁺ → Fe³⁺ + e⁻ by 2, then add: H₂O₂ + 2Fe²⁺ + 2H⁺ → 2Fe³⁺ + 2H₂O. Charge: left +6, right +6.
12. Fe³⁺ is the oxidising agent (Fe goes from +3 to +2). Oxidation: 2I⁻ → I₂ + 2e⁻.

Where marks are usually lost

• Describing the oxidising agent as “the substance oxidised”. It accepts electrons, so it is reduced.
• Defining oxidation only as “gain of oxygen”. The specification definition is electron loss.
• Giving an oxidation state without a sign (writing 6 instead of +6), or writing it as 6+.
• Missing a named exception, such as treating H in a metal hydride as +1.
• Giving the total for all atoms of the element instead of the value per atom, e.g. giving +14 for Cl in Cl₂O₇ instead of +7.
• In half-equations, balancing charge with extra H⁺ instead of electrons.
• Leaving electrons, or H⁺ and H₂O on both sides, in an overall equation.
• Choosing multipliers that make the electron numbers unequal, so the overall equation has the wrong ratio.
• Naming the element rather than the species, e.g. “chlorine is reduced” when the question asks which ion is the oxidising agent.

Related pages

• Study guide for this section
• Practice questions for this section
• Bonding: ions and ionic charges
• Atomic structure revision notes
• Mole revision notes, for redox titration calculations
• Free diagnostics

Official syllabus

AQA AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for AS and A-level exams June 2016 onwards, section 3.1.7 Oxidation, reduction and redox equations. Published by AQA.