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IB Chemistry Redox Reactions: A Clear Guide

IB Chemistry Redox Reactions: A Clear Guide

11 min readOliver Kidd (Co-founder)26 Oct 2026

A redox reaction is any reaction in which electrons move from one species to another. One species loses electrons (oxidation) while another gains them (reduction). In IB Chemistry you need to spot these reactions, track oxidation states, name the oxidising and reducing agents, write and combine half-equations, and explain what happens at the electrodes of a cell. This guide takes you through each skill in that order, with worked examples you can try yourself.

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TL;DR

  • Oxidation is loss of electrons and reduction is gain of electrons. They always happen together.
  • A reaction is redox when an oxidation state changes. An increase means oxidation, a decrease means reduction.
  • The reducing agent is the species that is oxidised. The oxidising agent is the species that is reduced.
  • To balance a redox equation, balance each half-equation for atoms and charge, equalise the electrons, add, then cancel.
  • In a cell, oxidation happens at the anode and reduction at the cathode. In a galvanic cell, a spontaneous reaction produces electrical energy.

Table of Contents

What oxidation and reduction mean

Oxidation is the loss of electrons. Reduction is the gain of electrons. Because an electron lost by one species must be gained by another, the two processes cannot be separated. That is why we call the whole reaction oxidation-reduction, or redox.

Zinc gives a clean example. In Zn → Zn²⁺ + 2e⁻, zinc loses two electrons, so it is oxidised. Whatever accepts those electrons is reduced.

Two electrons moving from zinc to an acceptor

A helpful idea: oxidation and reduction are relative. Whether a species is oxidised or reduced depends on what it reacts with. Iron(II) is oxidised by acidified manganate(VII), yet it can be reduced by a stronger reducing agent. Always ask, “compared with what?”

A memory aid many students use is OIL RIG: Oxidation Is Loss, Reduction Is Gain. Say it, but write the full words in an exam answer.

Using oxidation states to identify redox

Electron transfer is easy to see in ionic reactions. It is harder to see in covalent ones. Oxidation states solve this. An oxidation number is the charge an atom would have if the compound were ionic, and the numbers add up to the charge on the molecule or ion.

The test is simple. If oxidation numbers change, the reaction is redox: an increase means oxidation and a decrease means reduction. The same source adds a useful caution: oxidation numbers are bookkeeping, not always real atomic charges, especially in covalent compounds.

Rules to apply in order

  1. Atoms in an element have an oxidation state of 0.
  2. A simple ion has an oxidation state equal to its charge.
  3. Fluorine is −1. Oxygen is usually −2 and hydrogen usually +1.
  4. The sum for a neutral compound is 0, and for an ion it equals the ion’s charge.

A quick worked example

Take 2FeCl₂ + Cl₂ → 2FeCl₃. Iron goes from +2 to +3, so it is oxidised. Chlorine in Cl₂ goes from 0 to −1, so it is reduced. Redox, confirmed.

Now compare H₂O + SO₃ → H₂SO₄. Sulfur is +6 in both SO₃ and H₂SO₄, and nothing else changes. This is not redox. Questions that ask “which of these is a redox reaction?” are almost always testing this exact check.

You can practise the full set of electron-transfer skills with IB Chemistry R3.2 electron transfer questions, which sit under Reactivity 3.2 in the current syllabus.

Oxidising and reducing agents

Students mix these up more than any other term. The logic is this: the oxidised species gives up electrons, so it is the reducing agent. The reduced species takes electrons, so it is the oxidising agent. The agent is named for what it does to the other species, not for what happens to itself.

Use oxidation states as your check. The reducing agent’s oxidation number goes up. The oxidising agent’s goes down.

Species Change in oxidation state Process Role
Fe²⁺ → Fe³⁺ +2 to +3 Oxidised Reducing agent
Cl₂ → 2Cl⁻ 0 to −1 Reduced Oxidising agent
MnO₄⁻ → Mn²⁺ +7 to +2 Reduced Oxidising agent

The Royal Society of Chemistry puts the relationship plainly: oxidising agents accept electrons and are reduced, while reducing agents donate electrons and are oxidised. In an exam answer, always name the species and give the reason: “Fe²⁺ is the reducing agent because it loses an electron and its oxidation state rises from +2 to +3.”

Writing and combining half-equations

Half-equations are where most marks are won or lost. Write separate equations for the electron-losing and electron-gaining processes, balance atoms and charge, then place electrons on the product side for oxidation and the reactant side for reduction.

The method in acidic solution

  1. Write the skeleton half-equation for each process.
  2. Balance every atom except oxygen and hydrogen.
  3. Balance oxygen by adding H₂O.
  4. Balance hydrogen by adding H⁺.
  5. Balance charge by adding electrons.
  6. Multiply the half-equations so electrons lost equal electrons gained, add them, and cancel the electrons.

Worked example: manganate(VII) and iron(II)

Reduction: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O

Oxidation: Fe²⁺ → Fe³⁺ + e⁻

The reduction takes five electrons, so multiply the oxidation by 5. Add the two and the electrons cancel:

MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O

Check the charge. On the left: −1 + 8 + 10 = +17. On the right: +2 + 15 = +17. It balances.

The pattern to remember is that oxidation half-equations have electrons on the right and reduction half-equations have them on the left. If you ever find them on the wrong side, stop and recheck.

What redox looks like in electrochemical cells

Redox is not only a paper exercise. It powers batteries and drives electrolysis, and both appear in exam questions.

The key phrase to learn is “oxidation at the anode, reduction at the cathode”. This holds for every type of cell. What changes is the direction of the energy flow.

  • Galvanic (voltaic) cell: a spontaneous redox reaction produces electrical energy. The two half-reactions are kept in separate half-cells, and electrons travel through an external wire from the anode to the cathode.
  • Electrolytic cell: electrical energy forces a non-spontaneous reaction to happen. Oxidation still occurs at the anode and reduction at the cathode.

A galvanic cell example: in a zinc and copper cell, zinc is the more reactive metal, so it is oxidised at the negative electrode: Zn → Zn²⁺ + 2e⁻. Copper(II) ions are reduced at the positive electrode: Cu²⁺ + 2e⁻ → Cu. Electrons flow through the wire from zinc to copper, and ions move through the salt bridge to keep each half-cell neutral.

When a question asks you to “write the half-equation for the oxidation reaction”, find the electrode where the more reactive species is losing electrons. Then write it with electrons as a product.

Common mistakes that cost marks

  • Swapping the agents. The reducing agent is oxidised. Say it out loud before you write it down.
  • Forgetting that oxidation states are per atom. In Cr₂O₇²⁻, each chromium is +6, not +12.
  • Slipping when multiplying. OpenStax flags the multiplication step as a common source of error, so multiply every term in the half-equation, including the electrons.
  • Balancing atoms but not charge. Check every atom and the total charge before you move on.
  • Leaving electrons in the final equation. They must cancel completely.
  • Mixing up anode and cathode charges. The sign depends on the cell type, but oxidation is always at the anode.

How we teach redox at Tiber Tutor

Redox rewards a particular kind of confidence. Students who truly understand it stop memorising rules and begin to see electrons moving. Here is how we encourage that, as examiners and educators.

How we teach redox at Tiber Tutor — overview diagram

Start with oxidation states, because they give you a fast, reliable test. Then move to half-equations, because they show the electrons. Then connect both to cells, because that is where questions combine ideas. Redox links forward to energy and equilibrium too, so it pays to revisit it, for instance alongside enthalpy and ionic bonding.

We also advise students to practise on questions written in the style of the real papers, then mark their own answers against the examiner mark scheme. That is where you learn how marks are really awarded, such as the mark for naming the species and the mark for the oxidation-state change. Small habits like those turn a 5 into a 6 or 7.

— Oliver

Practise redox with Tiber Tutor

Once you understand the method, you need volume and feedback. Tiber Tutor brings exam-style questions, animated videos, notes, flashcards, cram sheets and topic tests together in one place. They are all linked to the syllabus, so a weak spot in redox leads you straight to the right material. Progress tracking and performance analytics show which sub-skills (half-equations, agents, cells) need work.

The content is written by IB examiners, and it is priced to be accessible: the Per-Subject Plan is $9 per month and the All-Access Plan is $19 per month. For the price, that is a strong offer for examiner-written material with this much depth, and it starts with a free 7-day trial. See the plans and pricing, try the IB Chemistry tests, or sit a full IB Chemistry mock exam when you feel ready.

FAQ

Is redox the same as electron transfer?

In most IB cases, yes. Oxidation is loss of electrons and reduction is gain. In covalent reactions electrons are shared rather than fully transferred, so we use oxidation states to track the change instead.

How do I know which species is the oxidising agent?

Find the species whose oxidation state goes down. It has gained electrons, so it has been reduced, and that makes it the oxidising agent. The species whose oxidation state goes up is the reducing agent.

Where do I put the electrons in a half-equation?

Put them on the right for oxidation and on the left for reduction. Then check that the total charge is equal on both sides of the half-equation.

Which electrode is the anode in a galvanic cell?

The anode is where oxidation happens, so it is the negative electrode in a galvanic cell. Electrons leave from it and travel through the external wire to the cathode, where reduction occurs.

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