Half-Equations and Redox: Explore the concepts of oxidation and reduction, and learn how to write and combine half-equations to form overall redox equations.
Half-Equations and Redox is part of Oxidation, reduction and redox equations in A Level AQA Chemistry.
5 min
Reading time
68
Lessons in this topic
A Level AQA
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Reviewed by Revise Right
Content review
Half-Equations and Redox: Explore the concepts of oxidation and reduction, and learn how to write and combine half-equations to form overall redox equations.
Use this lesson to stay inside Oxidation, reduction and redox equations while connecting the detail back to the wider Physical chemistry topic.
Read the core explanations from Half-Equations and Redox before testing yourself from memory.
Picture a reaction between iron and copper sulphate solution. Iron dissolves and copper metal deposits on the surface. Something is being transferred between the two species — and that something is electrons.
Redox reactions are, at their heart, electron-transfer events. One species loses electrons; another gains them. These two processes always happen together — you cannot have one without the other.
The memory tool that holds this together is OIL RIG: Oxidation Is Loss, Reduction Is Gain — of electrons. Burn that into your memory now, because every half-equation you write depends on it.
Now that you know what oxidation and reduction mean, the definitions of oxidising and reducing agents follow directly. An oxidising agent is an electron acceptor — it accepts electrons from another species, and in doing so, it causes that species to be oxidised.
A reducing agent is an electron donor — it donates electrons to another species, causing that species to be reduced. The reducing agent is itself oxidised in the process.
The relationship is reciprocal and worth stating clearly: the oxidising agent gets reduced; the reducing agent gets oxidised. This is where many students trip up in exam questions, so let's address it head-on.
With the definitions of oxidising and reducing agents secure, you are ready to write the equations that show electron transfer explicitly. These are called half-equations.
A half-equation shows either the oxidation process or the reduction process in isolation, with electrons written explicitly as reactants or products. This makes the electron transfer visible in a way that an overall equation does not.
The method for writing a half-equation follows a clear sequence. Work through each stage deliberately — rushing leads to unbalanced charges, which costs marks.
Worked example: Writing a Half-Equation: Oxidation of Iron(II) to Iron(III)
Write the half-equation for the oxidation of Fe²⁺(aq) to Fe³⁺(aq).
Answer: $\text{Fe}^{2+}(aq) \rightarrow \text{Fe}^{3+}(aq) + e^-$
Now apply the same logic to a reduction half-equation. The only difference is that electrons appear on the left-hand side — they are consumed by the species being reduced.
Worked example: Writing a Half-Equation: Reduction of Manganate(VII) in Acid
Write the half-equation for the reduction of MnO₄⁻(aq) to Mn²⁺(aq) in acidic solution.
Answer: $\text{MnO}_4^{-}(aq) + 8\text{H}^+(aq) + 5e^- \rightarrow \text{Mn}^{2+}(aq) + 4\text{H}_2\text{O}(l)$
Once you have both half-equations, combining them is a matter of ensuring the electrons cancel completely. The number of electrons lost in the oxidation half-equation must equal the number gained in the reduction half-equation.
If the electron counts do not match, multiply one or both half-equations by appropriate integers to equalise them. Then add the two equations together and cancel any species that appear on both sides.
Worked example: Combining Half-Equations: Fe²⁺ and MnO₄⁻ in Acid
Combine the oxidation of Fe²⁺ to Fe³⁺ with the reduction of MnO₄⁻ to Mn²⁺ in acidic solution to give the overall redox equation.
Answer: $5\text{Fe}^{2+}(aq) + \text{MnO}_4^{-}(aq) + 8\text{H}^+(aq) \rightarrow 5\text{Fe}^{3+}(aq) + \text{Mn}^{2+}(aq) + 4\text{H}_2\text{O}(l)$
Notice that electrons do not appear in the final overall equation. If you can still see electrons after combining, something has gone wrong — either the electron counts were not equalised, or the cancellation was incomplete.
The manganate(VII) example is the most common in AQA papers, but dichromate(VI) ions, $\text{Cr}_2\text{O}_7^{2-}$, appear regularly too. The method is identical — balance atoms, then charges, then combine.
Worked example: Half-Equation: Reduction of Dichromate(VI) in Acid
Write the half-equation for the reduction of Cr₂O₇²⁻(aq) to Cr³⁺(aq) in acidic solution.
Answer: $\text{Cr}_2\text{O}_7^{2-}(aq) + 14\text{H}^+(aq) + 6e^- \rightarrow 2\text{Cr}^{3+}(aq) + 7\text{H}_2\text{O}(l)$
This half-equation involves 6 electrons. If you were combining it with the Fe²⁺ oxidation half-equation (which produces 1 electron per iron), you would multiply the iron half-equation by 6 to equalise the electrons before adding.
Both atom balance and charge balance are required for a fully correct half-equation; neither condition on its own is sufficient.
When combining half-equations, the most common error is forgetting to multiply all coefficients in a half-equation when scaling it up. If you multiply the iron half-equation by 5, every species in that equation — including the Fe²⁺ and Fe³⁺ — must be multiplied by 5.
| Stage | What to check | Common error |
|---|---|---|
| Writing a half-equation | Electrons on the correct side: products for oxidation, reactants for reduction | Placing electrons on the wrong side — loses the mark for electron placement |
| Balancing atoms | Use H₂O for oxygen, H⁺ for hydrogen (in acid) | Forgetting to add H⁺ after adding H₂O, leaving hydrogen unbalanced |
| Balancing charge | Count total charge on each side; add electrons to equalise | Miscounting charge when multiple ions are present |
| Combining half-equations | Multiply to equalise electron counts before adding | Multiplying only some species in a half-equation, not all |
| Final overall equation | No electrons remaining; charges and atoms both balance | Leaving electrons in the final equation — automatic error |
AQA questions on this topic often ask you to 'identify the oxidising agent' or 'state which species is reduced'. These are one-mark answers, but they require precise language. Name the specific species (e.g. 'MnO₄⁻') rather than a vague description.
Half-equations and the ability to combine them sit at the heart of AQA redox chemistry. Every topic that follows — electrode potentials, electrolysis, and transition metal chemistry — builds directly on the electron-transfer language you have practised here.
Published by Revise Right. Use the official specification reference to check the requirements for your course. Content updated: 2026-09-05.
Move to the related lessons or back to the topic page when you need the wider sequence before exam-style practice.
Redox reactions are, at their heart, electron-transfer events. One species loses electrons; another gains them. These two processes always happen together — you cannot have one without the other.
By the end of this lesson, you should be able to Define oxidation and reduction in terms of electron transfer; Identify oxidising agents and reducing agents in redox reactions; Write half-equations for oxidation and reduction processes; Combine half-equations to construct a balanced overall redox equation.
Read each section, then close the notes and explain the idea from memory. Check your explanation against the page and finish with an exam question from Physical chemistry.