Revise Electron transfer reactions for IB Chemistry (R3.2)

At the centre of every redox reaction are two processes that always occur together, because electrons lost by one species are the electrons gained by another. The mnemonic OIL RIG captures both: Oxidation Is Loss of electrons, Reduction Is Gain of electrons. Historically 'oxidation' meant gaining oxygen and 'reduction' meant losing it, but the modern, general definition is written purely in terms of electron transfer — and equivalently in terms of oxidation state, which rises on oxidation and falls on reduction.

In an ionic reaction it is obvious which species gains and loses electrons, but most redox reactions involve covalent bonds where electrons are shared, not fully transferred. To track electrons in every case, chemists assign each atom an oxidation state (or oxidation number): the hypothetical charge it would carry if all its bonds to different elements were 100% ionic. A change in an atom's oxidation state between reactants and products is the definitive signal that a redox process has taken place. The states are assigned by a short hierarchy of rules — when two rules seem to conflict, the one higher in the list wins.

Naming the agents is where marks are most often thrown away, so fix the logic firmly. The oxidising agent is the species that brings about oxidation in something else — it does so by accepting the electrons, which means the oxidising agent is itself REDUCED. The reducing agent brings about reduction in something else by donating electrons, so the reducing agent is itself OXIDISED. The 'agent' is the cause; the change named applies to its partner, not to itself.

Because a redox reaction is two processes at once, the cleanest way to build and balance it is to write each half separately as a half-equation, with the electrons shown explicitly. A reduction half-equation gains electrons on the left-hand side; an oxidation half-equation loses electrons on the right-hand side. Each half must be balanced for both atoms AND charge — the electrons are what balance the charge. Once you have both halves, you scale them so the number of electrons is identical, then add them so the electrons cancel and disappear from the overall equation. In acidic solution you balance oxygen with H₂O and hydrogen with H⁺ before adding electrons.

Metals differ in how readily they give up electrons: the more reactive a metal, the more easily it is oxidised. Ranking metals by this tendency gives the reactivity (activity) series. Its practical payoff is displacement: a more reactive metal will reduce the ion of a less reactive metal, pushing it out of solution and taking its place. The classic demonstration is a strip of zinc in copper(II) sulfate solution — the blue colour fades as Zn²⁺ forms and a reddish deposit of copper metal appears on the zinc.

Usually the species that is oxidised and the species that is reduced are different. In a disproportionation reaction they are the same element in the same starting species: one portion is oxidised while another is reduced, so a single element ends up in two different oxidation states in the products. The standard example is chlorine reacting with cold, dilute sodium hydroxide: Cl₂ + 2NaOH → NaCl + NaClO + H₂O. Chlorine starts at oxidation state 0; in NaCl it has fallen to −1 (reduction) and in NaClO it has risen to +1 (oxidation). The same element is both the oxidising and the reducing agent.

(HL only.) At Higher Level the qualitative reactivity series is made quantitative. Each reduction half-reaction is assigned a standard electrode potential, E°, measured in volts against the standard hydrogen electrode (defined as exactly 0.00 V) under standard conditions: 298 K, 100 kPa gas pressure and 1 mol dm⁻³ solutions. Ranking half-reactions by E° gives the electrochemical series. A more positive E° means the species is more readily reduced, i.e. a stronger oxidising agent; a more negative E° means the species is more readily oxidised, i.e. a stronger reducing agent. These values are tabulated in the data booklet — you do not memorise them.

E°{cell} = E°{cathode} - E°_{anode}

(HL only.) For example, a Daniell cell pairs zinc, $E° = -0.76$ V, with copper, $E° = +0.34$ V. Copper has the more positive potential, so it is reduced at the cathode and zinc is oxidised at the anode. The cell potential is $E°_{cell} = (+0.34) - (-0.76) = +1.10$ V. The positive result confirms the familiar zinc-displaces-copper reaction is spontaneous — the same conclusion the qualitative reactivity series gave, now with a number attached.

Redox questions are marked analytically: each mark is tied to a specific step — a method mark (M) for the correct approach or an answer mark (A) for the correct result — and error-carried-forward (ECF) means an early slip does not have to cost you every mark after it, PROVIDED your method is written down. Study how each mark below is earned by a specific line of working.

Redox underpins much of the chemistry ahead. Reactivity trends explain why some metals corrode and others are used to protect them; electrochemistry turns these half-equations into working cells and batteries; and quantitative redox titrations combine half-equations with the mole calculations from S1.4 to measure unknown concentrations. Master the core habit here — assign oxidation states, split into half-equations, balance atoms then charge, then recombine so the electrons cancel — and every electrochemical topic becomes a variation on a method you already own.