Redox is a transfer of oxygen, hydrogen or electrons between substances. Oxidation and reduction always occur together in the same reaction, at the same time, which is why the whole process is called redox (reduction + oxidation).
Redox chemistry underpins electrolysis, simple cells, rusting and combustion. It is examined in every O-Level Science (Chemistry) paper, either directly or hidden inside another topic. The topic builds from the four paired definitions to oxidation states, oxidising and reducing agents, the two colour tests, electrolysis, simple cells and the hydrogen fuel cell.
What these notes cover
- •The four definitions of oxidation and reduction (OILRIG)
- •Oxidation state rules
- •Half-equations and ionic equations
- •Oxidising and reducing agents
- •Test for a reducing agent: acidified potassium manganate(VII)
- •Test for an oxidising agent: aqueous potassium iodide
- •Electrolysis and the electrode rules
- •Electrolysis of molten lead(II) bromide
- •Simple cells and the zinc–copper cell
- •The hydrogen fuel cell
- •Redox in combustion, rusting, bleaching and batteries
- •Common mistakes and exam technique
Key definitions
- Oxidation
- Gain of oxygen, loss of hydrogen, loss of electrons, or an increase in oxidation state.
- Reduction
- Loss of oxygen, gain of hydrogen, gain of electrons, or a decrease in oxidation state. It is the exact reverse of oxidation.
- Oxidation state
- A signed number (also called oxidation number) assigned to each atom in a species, used to track electron transfer on paper. Always write the sign before the number, e.g. +2, never 2.
- Oxidising agent
- An electron acceptor that causes the other substance to be oxidised. It is itself reduced, so its oxidation state decreases.
- Reducing agent
- An electron donor that causes the other substance to be reduced. It is itself oxidised, so its oxidation state increases.
Key formulas
Oxidation state sum rule
sum of oxidation states of all atoms = overall charge of the species
The overall charge is 0 for a neutral molecule, or the charge on an ion. A free element is 0, O is usually −2 and H is usually +1; use the sum rule to find the oxidation state of the one unknown atom in a formula.
Oxidation half-equation
M → Mⁿ⁺ + ne⁻
Electrons lost are written on the right, e.g. Zn(s) → Zn²⁺(aq) + 2e⁻. The electrons lost must equal the electrons gained in the other half-equation.
Reduction half-equation
Xⁿ⁺ + ne⁻ → X
Electrons gained are written on the left, e.g. Cu²⁺(aq) + 2e⁻ → Cu(s). In electrolysis, cations are reduced in this way at the cathode (−).
Electrolysis: reaction at the anode (+)
2X⁻ → X₂ + 2e⁻
Anions migrate to the anode and are oxidised, e.g. 2Br⁻(l) → Br₂(g) + 2e⁻. AN OX, RED CAT: oxidation at the anode, reduction at the cathode.
Test for a reducing agent
MnO₄⁻(aq) + 8H⁺(aq) + 5e⁻ → Mn²⁺(aq) + 4H₂O(l)
Acidified potassium manganate(VII) turns from purple to colourless as Mn goes from +7 to +2. K⁺ and SO₄²⁻ are spectator ions and are left out.
Test for an oxidising agent
2I⁻(aq) → I₂(aq) + 2e⁻
Aqueous potassium iodide turns from colourless to brown as iodine forms. The oxidation state of I goes from −1 to 0.
Worked example
Chlorine water is added to aqueous potassium iodide: Cl₂(aq) + 2KI(aq) → 2KCl(aq) + I₂(aq). Identify what is oxidised and what is reduced, name the oxidising and reducing agents, and state the observation.
- 1Assign oxidation states: Cl is 0 in Cl₂ and −1 in KCl; I is −1 in KI and 0 in I₂; K stays +1 throughout, so it is a spectator.
- 2The oxidation state of iodine increases from −1 to 0, so I⁻ (in KI) is oxidised.
- 3The oxidation state of chlorine decreases from 0 to −1, so Cl₂ is reduced.
- 4Name the agents by the opposite of what happens to each: Cl₂ is reduced, so chlorine is the oxidising agent; KI is oxidised, so potassium iodide (the I⁻ ion) is the reducing agent.
- 5Observation: the colourless solution turns brown as iodine is formed.
Watch out: If your answer says the oxidising agent is oxidised, you have swapped them — the agent always does the opposite to itself. Quote both colours: colourless to brown.
Common mistakes
Saying the oxidising agent is oxidised
Wrong: An oxidising agent gets oxidised in the reaction.
Right: The oxidising agent accepts electrons and is itself reduced. The reducing agent donates electrons and is itself oxidised.
Swapping the two colour tests
Wrong: Potassium manganate(VII) turning colourless shows an oxidising agent is present.
Right: Acidified potassium manganate(VII) turns from purple to colourless when a reducing agent is present; aqueous potassium iodide turns from colourless to brown (I₂ formed) when an oxidising agent is present.
Electrons flowing through the electrolyte
Wrong: Electrons travel from one electrode to the other through the solution to complete the circuit.
Right: Electrons flow only through the external wire. Inside the electrolyte the current is carried by ions: cations move towards the cathode and anions move towards the anode.
Questions students ask
What does OILRIG stand for in chemistry?
OILRIG stands for Oxidation Is Loss of electrons, Reduction Is Gain of electrons. It is the definition to fall back on when the others fail, because it works even when a reaction involves no oxygen and no hydrogen at all.
Why must the electrolyte be molten or aqueous in electrolysis?
In a solid ionic compound the ions are held in fixed positions in the giant ionic lattice, so they cannot move and the solid cannot conduct electricity. When the compound is molten or dissolved in water, the lattice is broken down and the ions are free to move, so they can carry charge and electrolysis becomes possible.
Which electrode is negative in a simple cell?
The more reactive metal is the negative electrode: it loses electrons, is oxidised and dissolves. The electrons travel through the external wire to the less reactive metal, which is the positive electrode. The further apart the two metals are in the reactivity series, the higher the voltage.
What do I need to know about the hydrogen fuel cell for Combined Chemistry?
Hydrogen, obtained from water or hydrocarbons, is a potential fuel that reacts with oxygen in a fuel cell to produce electricity directly, and water is the only product. The overall equation is 2H₂(g) + O₂(g) → 2H₂O(l). Electrode materials, the electrolyte inside and the half-equations at each electrode are not required.
In the full PDF
- •Labelled diagrams of the zinc–copper simple cell and an electrolysis cell to sketch from memory
- •Further worked examples: oxidation states in MnO₄⁻ and SO₄²⁻, zinc in copper(II) sulfate, the manganate(VII) test with iron(II) ions, molten lead(II) bromide, and Mg/Cu versus Fe/Cu cells
- •Redox in everyday contexts: combustion, rusting, oxidising and reducing bleaches, and batteries
- •The remaining misconceptions, including electrode polarity in electrolysis versus simple cells and the sign of the oxidation-state change
- •A seven-point exam technique checklist
- •A one-page summary of the whole topic