Writing formulas of ionic compounds

· Ionic compounds must have overall charge = 0.
· Write the cation first, then the anion.
· Use ion charges to work out the smallest whole-number ratio of ions.
· Do not write charges in the final formula: Mg²⁺ + 2Cl⁻ → MgCl₂.
· Use brackets when more than one polyatomic ion is needed: Ca²⁺ + 2NO₃⁻ → Ca(NO₃)₂.
· Do not use brackets if only one polyatomic ion is present: Na⁺ + NO₃⁻ → NaNO₃, not Na(NO₃).
· Predict common charges from Periodic Table position: Group 1 = +1, Group 2 = +2, Group 13 = +3, Group 15 = –3, Group 16 = –2, Group 17 = –1.
· For transition metals, the Roman numeral gives the oxidation number/charge: iron(III) = Fe³⁺, copper(I) = Cu⁺.
· Example: iron(III) sulfate = Fe³⁺ and SO₄²⁻ → balance charges 2 × Fe³⁺ and 3 × SO₄²⁻ → Fe₂(SO₄)₃.

This diagram shows how ionic charges determine the formula of an ionic compound. It is especially useful for remembering when brackets are needed around polyatomic ions. Source

Ions to memorise

· Nitrate = NO₃⁻
· Carbonate = CO₃²⁻
· Sulfate = SO₄²⁻
· Hydroxide = OH⁻
· Ammonium = NH₄⁺
· Zinc ion = Zn²⁺
· Silver ion = Ag⁺
· Hydrogen carbonate = HCO₃⁻
· Phosphate = PO₄³⁻
· Exam tip: learn these names, formulas and charges exactly; they are assumed knowledge for formula writing.

Balanced chemical equations

· A balanced equation has the same number of each type of atom on both sides.
· Balance using coefficients in front of formulas, never by changing subscripts inside formulas.
· Example: Mg + 2HCl → MgCl₂ + H₂.
· Coefficients give the mole ratio of reacting substances.
· Always check that atoms and, for ionic equations, overall charge are balanced.
· Formulae must be correct before balancing; a wrong formula usually makes the whole equation wrong.
· Include state symbols when required: (s) solid, (l) liquid, (g) gas, (aq) aqueous solution.

Ionic equations and spectator ions

· An ionic equation shows only the particles that actually react.
· Spectator ions are ions unchanged during the reaction; they appear on both sides and must be removed.
· To write an ionic equation:
· Start with the balanced full equation.
· Split aqueous ionic compounds into ions.
· Keep solids, liquids, gases and covalent molecules unchanged.
· Cancel identical spectator ions from both sides.
· Example: AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq).
· Full ionic: Ag⁺(aq) + NO₃⁻(aq) + Na⁺(aq) + Cl⁻(aq) → AgCl(s) + Na⁺(aq) + NO₃⁻(aq).
· Net ionic: Ag⁺(aq) + Cl⁻(aq) → AgCl(s).
· Final ionic equations must be balanced for atoms and charge.

Empirical and molecular formula

· Empirical formula = simplest whole-number ratio of atoms of each element in a compound.
· Molecular formula = actual number of atoms of each element in one molecule.
· A molecular formula is a whole-number multiple of the empirical formula.
· Example: C₆H₁₂O₆ has empirical formula CH₂O.
· To calculate an empirical formula from masses or percentages:
· Convert each mass/percentage to moles.
· Divide all mole values by the smallest mole value.
· Multiply if needed to get whole numbers.
· Use the whole-number ratio as subscripts.
· To calculate a molecular formula:
· Find empirical formula mass.
· Calculate multiplier: Mr ÷ empirical formula mass.
· Multiply every subscript in the empirical formula by this multiplier.
· Exam tip: if given percentages, assume 100 g so percentages become masses.

This diagram summarises how empirical formula mass links to molecular formula. It is useful for remembering the multiplier method used in exam calculations. Source

Hydrated, anhydrous and water of crystallisation

· Anhydrous = contains no water of crystallisation.
· Hydrated = contains water of crystallisation in a fixed ratio.
· Water of crystallisation = water molecules chemically included in a crystal lattice.
· Hydrated salts are written with a dot: CuSO₄·5H₂O.
· The dot means “combined with”; it does not mean multiplication in the chemical reaction sense.
· CuSO₄·5H₂O means 1 mol of CuSO₄ is combined with 5 mol of H₂O.
· Heating a hydrated salt can remove water: CuSO₄·5H₂O(s) → CuSO₄(s) + 5H₂O(g).
· To calculate water of crystallisation:
· Find mass of water lost = mass of hydrated salt – mass of anhydrous salt.
· Convert water and anhydrous salt masses to moles.
· Divide by the smallest mole value to get the ratio salt : water.
· Write the formula as salt·xH₂O.

Common exam pitfalls

· Do not change subscripts when balancing equations.
· Do not include spectator ions in the final ionic equation.
· Do not forget state symbols when the question asks for them.
· Do not write charges in the final formula of an ionic compound.
· Do not forget brackets around polyatomic ions when more than one is needed.
· Do not round mole ratios too early in empirical formula calculations.
· Do not confuse empirical formula with molecular formula.
· Do not treat water of crystallisation as ordinary liquid water in a formula; it is part of the crystal structure.