Compounds, Formulae and Equations: Ionic Charges and Balanced Equations

This guide covers 2.1.2, Compounds, formulae and equations, from Module 2, Foundations in chemistry, of OCR A Level Chemistry A H432. It sits between Atomic structure and isotopes (2.1.1) and Amount of substance (2.1.3), and supplies the formula-writing and equation-construction skills that every mole calculation in 2.1.3 depends on.

Writing formulae of ionic compounds from ionic charges

You need to be able to write the formulae of ionic compounds from ionic charges. This has two required components:

• Predicting ionic charge from periodic table position. Elements in a given group tend to form ions of a predictable charge related to that group – for example, Group 1 metals typically form 1+ ions and Group 7 elements typically form 1− ions.
• Recall of named polyatomic ions. The specification names seven ions you must know by name and formula: nitrate (NO₃⁻), carbonate (CO₃²⁻), sulfate (SO₄²⁻), hydroxide (OH⁻), ammonium (NH₄⁺), zinc (Zn²⁺) and silver (Ag⁺). A specific convention applies here: unless an oxidation number is explicitly shown, “nitrate” and “sulfate” should always be assumed to mean NO₃⁻ and SO₄²⁻ respectively (other oxidation states of these ions, such as nitrate(III)/nitrite, are covered separately under Redox). Charges on any ion outside this named list will be provided to you in an exam rather than expected from memory.

Constructing balanced chemical and ionic equations

You need to be able to construct balanced chemical equations, including ionic equations, with state symbols, both for reactions you have studied and for unfamiliar reactions given appropriate information. This is explicitly linked to mathematical skill M0.2 in the specification, which the same code series attaches to stoichiometric, empirical-formula and atom-economy calculations elsewhere in the syllabus – reflecting that balancing an equation is fundamentally a ratio process of finding integer coefficients that make atom counts equal on both sides.

The required elements of a fully correct equation are: correct formulae for every reactant and product (drawing on the formula-writing skill above), balanced coefficients so that atoms of each element are equal on both sides, and correct state symbols – (s), (l), (g) or (aq) – for every species. For unfamiliar reactions, the specification expects you to apply general patterns (such as the ionic-charge rules above, or reaction-type patterns like acid-base neutralisation) rather than having memorised the specific reaction in advance.

The criss-cross method for combining charges

When writing a formula from two ions of different charge magnitude, the standard technique is to swap the numerical value of each ion’s charge to become the other ion’s subscript. For example, combining Zn²⁺ with OH⁻: the 2 from zinc’s charge becomes the subscript on hydroxide, and the 1 from hydroxide’s charge becomes the (implicit) subscript on zinc, giving Zn(OH)₂ – note the bracket around the polyatomic hydroxide ion, which is required whenever more than one of a polyatomic ion appears in a formula. This bracketing convention is a frequent source of lost marks: writing “ZnOH2” instead of “Zn(OH)₂” implies a different, incorrect structure. The same method applied to Ag⁺ and CO₃²⁻ gives Ag₂CO₃ (silver’s charge of 1 becomes carbonate’s subscript, and carbonate’s charge of 2 becomes silver’s subscript, though a subscript of 1 is never written).

Why this subtopic matters beyond Module 2

Formula-writing and equation-balancing are not a one-off skill confined to this subtopic. Every stoichiometry calculation in 2.1.3 Amount of substance depends on having a correctly balanced equation to read mole ratios from, and this same equation-writing skill recurs directly in 2.1.4 Acids (neutralisation equations) and 2.1.5 Redox (equations involving electron transfer and oxidation numbers) – both of which build directly on the formula and equation conventions established here.

A worked example

Consider the reaction between aqueous silver nitrate and aqueous sodium chloride. Using the named/predictable ions above: silver is Ag⁺, nitrate is NO₃⁻, sodium is Na⁺ (Group 1, so 1+), and chloride is Cl⁻ (Group 7, so 1−). The full balanced equation, with state symbols, is:

AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

The ionic equation removes the “spectator ions” (Na⁺ and NO₃⁻, which appear unchanged on both sides) to leave only the species that actually react:

Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

Working through both the full and ionic versions of the same reaction is a useful way to check you can move fluently between the two forms, since exam questions may ask for either.

Common formula-writing errors

Three mistakes recur often enough to be worth naming directly. Forgetting brackets around a polyatomic ion when more than one is present (writing CaOH2 instead of Ca(OH)₂) is the single most common error. Confusing a compound’s name oxidation state with the default assumed one – for instance, writing NO₂⁻ when a question simply says “nitrate” without specifying an oxidation number, when the unqualified term always means NO₃⁻ – is another. The third is applying the criss-cross method without first simplifying the resulting ratio to lowest terms: combining Mg²⁺ with O²⁻ using the raw criss-cross method gives Mg₂O₂, which must be simplified to the correct formula MgO.

How to approach it

Learn the seven named polyatomic ions (and their formulae) as a fixed list you can recall instantly, since the specification treats these as required knowledge rather than information that will be provided. Practise predicting simple ionic charges from periodic table group number until it becomes automatic, particularly for Groups 1, 2, 6 and 7 where the pattern is most consistent. When constructing equations for unfamiliar reactions, work systematically: write correct formulae first using the charge rules, then balance coefficients, then add state symbols last – attempting to balance an equation before the formulae are correct is one of the most common sources of error. For ionic equations specifically, get in the habit of identifying and cancelling spectator ions explicitly rather than trying to write the ionic equation directly from the reaction description.

Official syllabus

OCR, Cambridge OCR Level 3 Advanced GCE in Chemistry A (H432) specification, Version 3.1, May 2026: official specification PDF, section 2.1.2 “Compounds, formulae and equations”. Verified 2026-09-02.