Redox — OCR A-Level Chemistry

An oxidation number (oxidation state) is a bookkeeping value that tracks electron control in a species. For an uncombined element each atom has oxidation number 0, as in Fe, O₂ and S₈. For a monatomic ion the oxidation number equals the ion charge, so Na⁺ is +1 and O²⁻ is −2. In compounds, hydrogen is usually +1 and oxygen usually −2, with exceptions such as metal hydrides (H = −1) and peroxides (O = −1). Fluorine is always −1. The sum of oxidation numbers equals the overall charge: zero for a neutral compound, the ion charge for a polyatomic ion. In H₂SO₄, hydrogen contributes 2 × (+1) = +2 and oxygen 4 × (−2) = −8, so sulfur must be +6 to give a total of 0. Write oxidation numbers sign-first, for example +6, never as a charge superscript.

Assigning oxidation numbers relies on a short set of rules applied in order. An atom in an uncombined element has oxidation number 0, so each atom in O₂, Cl₂ or Fe is 0. A monatomic ion takes the value of its charge, so Mg²⁺ is +2 and Cl⁻ is −1. In compounds, fluorine is always −1; hydrogen is usually +1 except in metal hydrides where it is −1; oxygen is usually −2 except in peroxides where it is −1 and in compounds with fluorine. The sum of all oxidation numbers equals the overall charge: 0 for a neutral compound, the stated charge for a polyatomic ion. To find an unknown, multiply each known value by its subscript, subtract from the overall charge and divide by the number of unknown atoms. In Cr₂O₇²⁻, oxygen gives 7 × (−2) = −14, so chromium must total +12, giving +6 per Cr. Write values sign-first, for example +6.

Oxidation numbers let you build a formula from the charges that combine to give a neutral compound or a stated ion. Assign the known oxidation number to each element using its group and the rules (oxygen usually −2, hydrogen usually +1, group 1 metals +1, group 2 metals +2), then choose the ratio of atoms whose total oxidation numbers sum to zero for a neutral species or to the ion charge for a polyatomic ion. For example, iron(III) oxide needs Fe at +3 and O at −2, so the smallest whole-number ratio is two Fe to three O, giving Fe₂O₃. For an ion such as sulfate, S at +6 and four O at −2 give a total of −2, matching SO₄²⁻. Always reduce the ratio to its simplest whole numbers and check the total charge balances.

Some transition elements form more than one compound or ion, so the name alone is ambiguous. A Roman numeral in brackets after the element name states the oxidation number of that element in that particular species. Iron can be +2 or +3, so iron(II) chloride contains Fe²⁺ and iron(III) chloride contains Fe³⁺; copper(I) oxide contains Cu⁺ while copper(II) oxide contains Cu²⁺. The numeral gives the magnitude of the oxidation number, not the number of atoms and not a charge written as a superscript. Read the numeral, assign that oxidation number to the element, then deduce the formula or the ion charge from the other element present.

Redox reactions are electron-transfer processes, and OCR expects you to recognise them from oxidation numbers, from half-equations and from observations. Oxidation is an increase in oxidation number, reduction is a decrease, and the two always occur together because the electrons lost by one species are gained by another. For example, in Zn + Cu²⁺ → Zn²⁺ + Cu, zinc's oxidation number rises from 0 to 2⁺ (oxidised) while copper's falls from 2⁺ to 0 (reduced). You must also handle disproportionation, where one element is simultaneously oxidised and reduced, as in chlorine with cold dilute alkali. Practise assigning oxidation numbers using the rules, then check that the changes balance in a full equation.

This criterion asks you to describe oxidation and reduction using two complementary frameworks: electron transfer and oxidation number change. In electron-transfer terms, oxidation is loss of electrons and reduction is gain of electrons, summarised by OIL RIG. In oxidation-number terms, oxidation is an increase and reduction a decrease. You should be able to move between the two descriptions for the same reaction. For example, in Fe²⁺ → Fe³⁺ + e⁻, iron loses an electron (oxidation) and its oxidation number rises from 2⁺ to 3⁺. Practise converting a half-equation into an oxidation-number statement and vice versa, and check that electrons appear on the correct side.

Redox in OCR A Level Chemistry is defined by electron transfer. Oxidation is loss of electrons and reduction is gain of electrons, remembered as OIL RIG. In the reaction between zinc and copper(II) ions, zinc atoms each lose two electrons to form Zn²⁺ while Cu²⁺ ions each gain two electrons to form copper atoms, so zinc is oxidised and copper(II) ions are reduced. A multiple-choice question may ask which species is oxidised, which is the reducing agent, or how many electrons move. The reducing agent is the species that donates electrons and is itself oxidised; the oxidising agent accepts electrons and is itself reduced. Count electrons from the charge change on each ion, remembering that a species losing electrons becomes more positive.

Oxidation number tracks the apparent charge on an atom and gives a second way to recognise redox. An increase in oxidation number is oxidation and a decrease is reduction. Assign oxidation numbers using the rules: an uncombined element has oxidation number 0; a monatomic ion has its charge; oxygen is usually −2 and hydrogen usually +1 in compounds; the sum of oxidation numbers equals the overall charge. In manganate(VII) the manganese has oxidation number +7, and it falls to +2 in Mn²⁺, so manganese is reduced. In a multiple-choice question, work out the oxidation number of the element before and after, then decide whether it rose or fell.

A metal above hydrogen in the reactivity series reacts with a dilute acid to give a salt and hydrogen gas. The metal atoms lose electrons (oxidation) and the hydrogen ions in the acid gain electrons (reduction), so the reaction is redox. For example, magnesium with dilute hydrochloric acid gives magnesium chloride and hydrogen: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g). Write full balanced equations, including state symbols, and check that atoms and charge balance. The salt contains the metal cation and the acid anion, so the acid chosen fixes the salt formed. Copper, silver and gold sit below hydrogen and do not react with dilute acids in this way.

To interpret a redox equation, assign oxidation numbers to each element and track how they change. An increase in oxidation number means oxidation (loss of electrons); a decrease means reduction (gain of electrons). In Mg + 2HCl → MgCl₂ + H₂, magnesium goes from 0 to 2⁺ (oxidised) and hydrogen in HCl goes from 1⁺ to 0 (reduced). Apply the same method to unfamiliar reactions: work out oxidation numbers from the rules, spot which element changes, and predict the direction of electron transfer. Remember that oxidation numbers are a bookkeeping tool, and that the element oxidised is the reducing agent while the element reduced is the oxidising agent.