Acids, Alkalis and Neutralisation

The pH Scale

The pH scale runs from 0 to 14 and measures the concentration of H⁺ ions in a solution:

• pH 0–6: Acid (more H⁺ ions than OH⁻ ions)
• pH 7: Neutral (equal H⁺ and OH⁻ ions) — pure water at 25°C
• pH 8–14: Alkali (more OH⁻ ions than H⁺ ions)

The pH scale is logarithmic: a change of 1 pH unit represents a 10× change in H⁺ ion concentration. So pH 3 is 10× more acidic than pH 4, and 100× more acidic than pH 5.

Indicators: Litmus (red in acid, blue in alkali); universal indicator (full colour range from red to purple); phenolphthalein (colourless in acid, pink in alkali).

Strong vs Weak Acids

Strong acids fully dissociate (ionise) in water → maximum H⁺ ions for a given concentration:

• Hydrochloric acid: HCl → H⁺ + Cl⁻
• Sulfuric acid: H₂SO₄ → 2H⁺ + SO₄²⁻
• Nitric acid: HNO₃ → H⁺ + NO₃⁻

Weak acids only partially dissociate:

• Ethanoic acid (vinegar), citric acid, carbonic acid
• A 1 mol/dm³ solution of ethanoic acid has a much higher pH than 1 mol/dm³ HCl — fewer H⁺ ions

Important: Strength (strong/weak) refers to degree of dissociation, NOT concentration. You can have a dilute strong acid or a concentrated weak acid.

Neutralisation

Neutralisation: The reaction between an acid and a base (or alkali) to produce a salt and water.

General equation: Acid + Base → Salt + Water

Examples:

• HCl + NaOH → NaCl + H₂O (hydrochloric acid + sodium hydroxide → sodium chloride + water)
• H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O
• HNO₃ + NaOH → NaNO₃ + H₂O

Ionic equation for neutralisation: H⁺(aq) + OH⁻(aq) → H₂O(l) This is always the same, regardless of which specific acid or alkali is used.

Making Salts

From acids + alkalis (titration):

1. Add indicator to the alkali in a conical flask
2. Add acid from a burette until the end point (indicator changes colour) — neutralisation point
3. Repeat without indicator; evaporate water to obtain salt crystals

Naming salts:

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

From acids + metal oxides/carbonates:

• Acid + metal oxide → salt + water
• Acid + metal carbonate → salt + water + carbon dioxide
• H₂SO₄ + CuO → CuSO₄ + H₂O (copper sulfate)
• HCl + CaCO₃ → CaCl₂ + H₂O + CO₂

pH and H⁺ ions — the Core Concept

Acids donate H⁺ ions (protons) to solutions. Bases accept H⁺ ions (or donate OH⁻ ions). Neutralisation is fundamentally about H⁺ and OH⁻ ions combining to form water.

Exam tip: Know the formula for the ionic equation of neutralisation; know how to name salts from the acid used; understand that pH is logarithmic (each unit = 10× change in [H⁺]).