Part 1: Acids, alkalis and salts

Separate acid strength from concentration and predict products using the reacting ions.

Acids release hydrogen ions in water

At this stage, an acid releases H⁺ ions in aqueous solution. An alkali is a soluble base that produces OH⁻ ions in aqueous solution. A base neutralises an acid; it need not dissolve in water. NaOH and KOH are alkalis; CuO is an insoluble base.

Hydrochloric acid HCl, nitric acid HNO₃ and sulfuric acid H₂SO₄ are familiar strong acids. Ethanoic acid CH₃COOH is weak: only a fraction of its molecules ionise. Strong means essentially complete ionisation in the relevant aqueous context; concentrated means a large amount of solute per volume. A dilute strong acid and a concentrated weak acid are both possible.

Aqueous ammonia is an alkali even though NH₃ contains no OH group: it reacts with water to form NH₄⁺ and OH⁻. A weak acid is not necessarily harmless; hazards also depend on concentration and exposure.

Separate strength, concentration and amount

Concentration says how much solute is present per unit volume. Strength describes how far an acid ionises in water. Amount says how many moles are present in the particular sample. A small volume of concentrated solution can contain fewer moles than a large volume of dilute solution; none of these words can replace another.

Compare 25.0 cm³ of 0.100 mol dm⁻³ HCl with the same volume and concentration of CH₃COOH. Each contains 0.00250 mol acid and each can neutralise 0.00250 mol OH⁻. Their initial hydrogen-ion concentrations differ because ethanoic acid is only partly ionised. As H⁺ is consumed during neutralisation, further weak-acid molecules ionise.

This is why “weak acid needs less alkali because it only partly reacts” is incorrect for complete neutralisation of equal amounts of a monoprotic acid. Keep quantitative pH and Ka calculations for Year 13; the AS task is to reason correctly about ionisation and reaction amounts.

Choose the salt from the acid and base

HCl makes chlorides, HNO₃ nitrates and H₂SO₄ sulfates. The base supplies the metal or ammonium ion. A carbonate also gives CO₂, seen as effervescence and confirmed by limewater. An oxide or hydroxide forms water; ammonia accepts H⁺ to form ammonium without making water in the overall NH₃ + acid equation.

Construct the salt before balancing the reaction

For magnesium hydroxide plus nitric acid, the salt combines Mg²⁺ and NO₃⁻, giving Mg(NO₃)₂. Start Mg(OH)₂ + HNO₃ → Mg(NO₃)₂ + H₂O. Two nitrate ions require 2HNO₃; the resulting four hydrogen atoms require 2H₂O. Final: Mg(OH)₂ + 2HNO₃ → Mg(NO₃)₂ + 2H₂O.

For ammonia plus sulfuric acid, complete neutralisation gives ammonium sulfate, (NH₄)₂SO₄: 2NH₃ + H₂SO₄ → (NH₄)₂SO₄. Do not add water automatically: ammonia has no hydroxide group to make water in this overall equation.

A carbonate has a different observation from an oxide or hydroxide because it generates CO₂. Write the salt formula first and include water and carbon dioxide, then balance. Neutralisation can still be exothermic when no gas or obvious colour change occurs; absence of bubbling is not evidence of no reaction.

Use the appropriate ionic equation

For a strong acid and hydroxide alkali, H⁺(aq) + OH⁻(aq) → H₂O(l). This is a reaction of ions; it is not a redox reaction because oxidation numbers do not change.

For a carbonate, CO₃²⁻ + 2H⁺ → CO₂ + H₂O. The factor of two is essential when calculating acid needed. Equations and mole ratios are more reliable than assuming equal acid and alkali volumes neutralise each other.

Quick checks

Original Finesse questions. Reveal the indicative worked solutions after attempting each question; these are not official OCR A mark allocations.

Sources