Redox Processes: Electron Transfer and Oxidation Number

· Redox = a reaction involving both oxidation and reduction.
· Oxidation = loss of electrons OR increase in oxidation number.
· Reduction = gain of electrons OR decrease in oxidation number.
· Mnemonic: OIL RIG = Oxidation Is Loss, Reduction Is Gain.
· In every redox reaction, electrons lost = electrons gained.
· A species cannot be oxidised unless another species is reduced at the same time.

This diagram links electron loss with oxidation and electron gain with reduction. It is useful for remembering that redox always involves two linked processes. Source

Oxidation Numbers / Oxidation States

· Oxidation number = the apparent charge an atom would have if bonding electrons were assigned to the more electronegative atom.
· Oxidation numbers are used to track electron transfer in redox reactions.
· Elements in their standard state have oxidation number 0: e.g. Na(s), Cl₂(g), O₂(g), H₂(g).
· A monatomic ion has oxidation number equal to its ionic charge: Na⁺ = +1, Cl⁻ = −1, Fe³⁺ = +3.
· The sum of oxidation numbers in a neutral compound = 0.
· The sum of oxidation numbers in a polyatomic ion = the charge on the ion.
· Common rules: Group 1 = +1, Group 2 = +2, Al = +3, F = −1.
· Oxygen is usually −2, except in peroxides or compounds with fluorine.
· Hydrogen is usually +1, except in metal hydrides where it is −1.

Oxidation-state diagrams show how oxidation numbers can be assigned to atoms in molecules or ions. They help students practise calculating oxidation numbers rather than guessing them. Source

Calculating Oxidation Numbers

· Step 1: assign known oxidation numbers first, e.g. O = −2, H = +1, Group 1 = +1.
· Step 2: let the unknown oxidation number be x.
· Step 3: write an equation using the total charge of the compound or ion.
· Step 4: solve for x.
· Example: in SO₄²⁻, oxygen = −2, so S + 4(−2) = −2; therefore S = +6.
· Example: in KMnO₄, K = +1, O = −2, so +1 + Mn + 4(−2) = 0; therefore Mn = +7.
· Always include the sign: write +6, not just 6.

Oxidation Number Changes

· If oxidation number increases, the species is oxidised.
· If oxidation number decreases, the species is reduced.
· Example: Zn → Zn²⁺ + 2e⁻; zinc changes from 0 to +2, so zinc is oxidised.
· Example: Cu²⁺ + 2e⁻ → Cu; copper changes from +2 to 0, so copper ions are reduced.
· Oxidation-number changes are often easier to spot than electron transfer in covalent reactions.

Oxidising and Reducing Agents

· An oxidising agent causes another species to be oxidised.
· The oxidising agent gains electrons and is itself reduced.
· A reducing agent causes another species to be reduced.
· The reducing agent loses electrons and is itself oxidised.
· Exam rule: identify the agent by what it does to the other species, then state what happens to the agent itself.
· Example: Cu²⁺ + Zn → Cu + Zn²⁺
· Zn is the reducing agent because it reduces Cu²⁺ to Cu; Zn is oxidised from 0 to +2.
· Cu²⁺ is the oxidising agent because it oxidises Zn to Zn²⁺; Cu²⁺ is reduced from +2 to 0.

This image is helpful for avoiding the common mistake of confusing the process with the agent. The reducing agent is oxidised, while the oxidising agent is reduced. Source

Balancing Redox Equations Using Oxidation Numbers

· Use oxidation-number changes to make electron loss = electron gain.
· Step 1: assign oxidation numbers to atoms that change.
· Step 2: identify which element is oxidised and which is reduced.
· Step 3: calculate the size of each oxidation-number change.
· Step 4: multiply species so the total increase equals the total decrease.
· Step 5: balance the remaining atoms and charges if required.
· Final check: atoms and total charge must be balanced on both sides.
· Never leave electrons in the final overall equation.

Disproportionation

· Disproportionation = a redox reaction where the same element is both oxidised and reduced.
· One species starts in an intermediate oxidation state and forms products where that element has both a higher and a lower oxidation number.
· Key test: look for the same element increasing and decreasing in oxidation number in the same reaction.
· Example pattern: X in one oxidation state → X in a higher oxidation state + X in a lower oxidation state.
· In exam answers, state clearly which species is oxidised and which species is reduced.

Roman Numerals in Compound Names

· A Roman numeral shows the oxidation number of an element in a compound or ion.
· Roman numerals are especially important for elements with variable oxidation states, e.g. transition metals.
· Iron(II) means iron has oxidation number +2.
· Iron(III) means iron has oxidation number +3.
· Copper(II) oxide contains Cu²⁺ and O²⁻, so the formula is CuO.
· Do not confuse a Roman numeral with the number of atoms in the formula.