Acids, bases and salts

pH and the pH scale

The pH scale measures how acidic or alkaline a solution is:

• pH 0–6: acid (more acidic closer to 0)
• pH 7: neutral
• pH 8–14: alkali (more alkaline closer to 14)

Measured using: universal indicator, pH probe, or indicator solutions.

Acids produce H⁺ ions in water. Alkalis produce OH⁻ ions in water.

Common acids: hydrochloric acid (HCl), sulfuric acid (H₂SO₄), nitric acid (HNO₃). Common alkalis: sodium hydroxide (NaOH), potassium hydroxide (KOH), calcium hydroxide (Ca(OH)₂), ammonia (NH₃(aq)).

Neutralisation

Acid + Base → Salt + Water

A base is any substance that neutralises an acid. An alkali is a base that dissolves in water.

Ionic equation for any neutralisation: H⁺(aq) + OH⁻(aq) → H₂O(l)

Examples:

• HCl + NaOH → NaCl + H₂O
• H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O
• HCl + CaCO₃ → CaCl₂ + H₂O + CO₂ (carbonates also produce CO₂)

Naming salts

Salt name = metal + acid suffix:

• Hydrochloric acid → chloride
• Sulfuric acid → sulfate
• Nitric acid → nitrate

Examples: sodium chloride (NaCl), copper(II) sulfate (CuSO₄), calcium nitrate (Ca(NO₃)₂).

Methods of salt preparation

1. Acid + insoluble base/oxide (excess solid method): Add excess solid until no more dissolves → filter → evaporate filtrate. e.g. H₂SO₄ + CuO → CuSO₄ + H₂O

2. Acid + carbonate: Acid + carbonate → salt + water + CO₂ e.g. 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂

3. Precipitation (insoluble salts): Mix two solutions whose ions form an insoluble product → filter, wash and dry precipitate. e.g. BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)

4. Titration (acid + alkali): Add acid to alkali from a burette using an indicator to find endpoint; repeat without indicator; evaporate the neutral solution.

Titration calculations (Higher tier)

Concentration = moles / volume (dm³) Moles = concentration × volume (dm³) 1 dm³ = 1000 cm³, so divide cm³ by 1000 to get dm³.

From the balanced equation, use mole ratios to find moles of the other reactant.