What you'll learn
This topic covers the fundamental chemistry of acids and bases, including their properties, reactions, and practical applications. Understanding acids, bases, and their neutralisation reactions is essential for Paper 2 and Paper 4, accounting for approximately 10-15% of exam marks. You must know how to predict products, write balanced equations, and explain the pH scale.
Key terms and definitions
Acid — a substance that produces hydrogen ions (H⁺) when dissolved in water; has pH less than 7.
Base — a substance that neutralises acids; metal oxides and metal hydroxides are common bases.
Alkali — a soluble base that produces hydroxide ions (OH⁻) when dissolved in water; has pH greater than 7.
Neutralisation — the reaction between an acid and a base to produce a salt and water only; H⁺(aq) + OH⁻(aq) → H₂O(l).
Salt — an ionic compound formed when the hydrogen ion of an acid is replaced by a metal ion or ammonium ion.
Indicator — a substance that changes colour depending on whether it is in acidic or alkaline solution.
Strong acid — an acid that is completely ionised in aqueous solution, releasing all its hydrogen ions.
Weak acid — an acid that is only partially ionised in aqueous solution, releasing only some of its hydrogen ions.
Core concepts
Properties and identification of acids and bases
Acids share characteristic properties that allow identification:
- Taste sour (never taste chemicals in the laboratory)
- Turn blue litmus paper red
- Have pH values less than 7
- React with metals above hydrogen in the reactivity series to produce hydrogen gas
- React with carbonates to produce carbon dioxide gas
- Conduct electricity when dissolved in water due to mobile ions
Common laboratory acids include:
- Hydrochloric acid (HCl) — strong acid
- Sulfuric acid (H₂SO₄) — strong acid
- Nitric acid (HNO₃) — strong acid
- Ethanoic acid (CH₃COOH) — weak acid found in vinegar
- Citric acid — weak acid found in citrus fruits
Bases and alkalis have contrasting properties:
- Feel soapy or slippery (never touch corrosive chemicals)
- Turn red litmus paper blue
- Have pH values greater than 7
- React with acids in neutralisation reactions
Common bases and alkalis:
- Sodium hydroxide (NaOH) — strong alkali
- Potassium hydroxide (KOH) — strong alkali
- Calcium hydroxide (Ca(OH)₂) — weak alkali, also called slaked lime
- Ammonia solution (NH₃(aq)) — weak alkali
- Copper(II) oxide (CuO) — insoluble base
- Zinc oxide (ZnO) — insoluble base
The pH scale and indicators
The pH scale measures the acidity or alkalinity of a solution, ranging from 0 to 14:
- pH 0-6: acidic solutions (lower pH = stronger acid)
- pH 7: neutral solution (pure water)
- pH 8-14: alkaline solutions (higher pH = stronger alkali)
Each unit change on the pH scale represents a tenfold change in hydrogen ion concentration. Moving from pH 3 to pH 2 means the hydrogen ion concentration increases by a factor of 10.
Universal indicator shows a range of colours:
- Red/orange: pH 1-4 (strong to weak acids)
- Yellow: pH 5-6 (weak acid)
- Green: pH 7 (neutral)
- Blue: pH 8-11 (weak to strong alkali)
- Purple: pH 12-14 (very strong alkali)
Litmus is a simpler indicator:
- Red in acid
- Purple in neutral
- Blue in alkali
Methyl orange turns:
- Red in acid (pH < 4.4)
- Yellow in alkali (pH > 6.2)
Phenolphthalein turns:
- Colourless in acid
- Pink/magenta in alkali (pH > 8.2)
Strong acids vs weak acids
The distinction between strong and weak acids relates to ionisation in water:
Strong acids completely dissociate:
- HCl(aq) → H⁺(aq) + Cl⁻(aq)
- H₂SO₄(aq) → 2H⁺(aq) + SO₄²⁻(aq)
- HNO₃(aq) → H⁺(aq) + NO₃⁻(aq)
All acid molecules release their hydrogen ions, producing high concentrations of H⁺ ions and lower pH values (typically pH 0-2 at normal concentrations).
Weak acids partially dissociate:
- CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq)
Only a small proportion of acid molecules release hydrogen ions. An equilibrium exists between molecules and ions. This produces lower concentrations of H⁺ ions and higher pH values (typically pH 3-6 at similar concentrations to strong acids).
The same distinction applies to bases: sodium hydroxide is a strong alkali (completely dissociates), while ammonia is a weak alkali (partially dissociates).
Neutralisation reactions
Neutralisation is the reaction between H⁺ ions from acids and OH⁻ ions from alkalis:
H⁺(aq) + OH⁻(aq) → H₂O(l)
This ionic equation applies to all neutralisation reactions involving alkalis. The reaction is exothermic, releasing heat energy.
Acid + Metal oxide → Salt + Water
Example: sulfuric acid + copper(II) oxide → copper(II) sulfate + water
H₂SO₄(aq) + CuO(s) → CuSO₄(aq) + H₂O(l)
Acid + Metal hydroxide → Salt + Water
Example: hydrochloric acid + sodium hydroxide → sodium chloride + water
HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)
Acid + Metal carbonate → Salt + Water + Carbon dioxide
Example: nitric acid + calcium carbonate → calcium nitrate + water + carbon dioxide
2HNO₃(aq) + CaCO₃(s) → Ca(NO₃)₂(aq) + H₂O(l) + CO₂(g)
The carbon dioxide produced can be tested with limewater, which turns milky/cloudy.
Acid + Ammonia → Ammonium salt
Example: hydrochloric acid + ammonia → ammonium chloride
HCl(aq) + NH₃(aq) → NH₄Cl(aq)
Ammonium salts are important fertilisers.
Preparing salts
CIE IGCSE Chemistry requires knowledge of two main methods for salt preparation:
Method 1: Insoluble base + Acid (for soluble salts)
- Add excess insoluble base (metal oxide or metal carbonate) to warm dilute acid in a beaker
- Stir until no more base reacts (excess base remains)
- Filter to remove unreacted base
- Heat the filtrate to evaporate some water
- Leave to crystallise
- Filter and dry the crystals between filter paper
This method produces pure, dry salt crystals. The excess base ensures all acid is neutralised.
Method 2: Alkali + Acid (for soluble salts using titration)
- Use an indicator to find the exact volume of alkali needed to neutralise a measured volume of acid
- Repeat without indicator, using the same volumes
- Evaporate the solution to crystallisation point
- Filter and dry the crystals
This method is necessary when both reactants are soluble, as you cannot use excess and filter.
Naming salts: The salt name comes from:
- First part: metal (or ammonium) from the base
- Second part: acid used (hydrochloric → chloride, sulfuric → sulfate, nitric → nitrate)
Examples:
- Sodium hydroxide + sulfuric acid → sodium sulfate
- Zinc oxide + hydrochloric acid → zinc chloride
- Calcium carbonate + nitric acid → calcium nitrate
Applications of neutralisation
Agriculture: Farmers add slaked lime (calcium hydroxide) or powdered limestone (calcium carbonate) to acidic soils to neutralise excess acidity and optimise pH for crop growth.
Indigestion remedies: Antacid tablets contain bases such as calcium carbonate or magnesium hydroxide that neutralise excess hydrochloric acid in the stomach.
Insect stings: Bee stings are acidic and can be treated with baking soda (sodium hydrogencarbonate, a weak alkali). Wasp stings are alkaline and can be treated with vinegar (ethanoic acid).
Industrial processes: Neutralisation is used to treat acidic factory waste before discharge, preventing environmental damage to rivers and lakes.
Worked examples
Example 1: A student adds magnesium ribbon to dilute hydrochloric acid. A gas is produced.
(a) Name the gas produced. [1]
(b) Describe a test for this gas. [2]
(c) Write a balanced chemical equation for the reaction. [2]
Solution:
(a) Hydrogen [1]
(b) Use a lighted splint [1]; the gas burns with a squeaky pop [1]
(c) Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g) [2] (1 mark for correct formulae, 1 mark for balancing)
Example 2: A student prepares copper(II) sulfate crystals by adding excess copper(II) oxide to warm dilute sulfuric acid.
(a) How can the student tell when the reaction is complete? [1]
(b) Why is excess copper(II) oxide used? [1]
(c) Describe how the student obtains pure, dry copper(II) sulfate crystals from the reaction mixture. [4]
Solution:
(a) Black copper(II) oxide remains in the beaker/solid stops disappearing [1]
(b) To ensure all the acid is neutralised/used up [1]
(c) Filter to remove excess copper(II) oxide [1]; heat/warm the filtrate/solution to evaporate some water [1]; leave to cool/crystallise [1]; filter and dry the crystals between filter paper/in a warm oven [1]
Example 3: Ethanoic acid has a pH of 4.5 at a concentration of 0.1 mol/dm³. Hydrochloric acid at the same concentration has a pH of 1.0.
(a) State which is the stronger acid. [1]
(b) Explain the difference in pH values. [2]
Solution:
(a) Hydrochloric acid [1]
(b) Hydrochloric acid is a strong acid/completely ionised [1]; it produces a higher concentration of hydrogen ions/H⁺ ions than ethanoic acid, which is a weak acid/partially ionised [1]
Common mistakes and how to avoid them
Mistake: Confusing bases and alkalis, using the terms interchangeably. Correction: All alkalis are bases, but not all bases are alkalis. A base is any substance that neutralises acids; an alkali is a soluble base that produces OH⁻ ions in water. Copper(II) oxide is a base but not an alkali because it is insoluble.
Mistake: Writing H₂ instead of H⁺ when describing acids. Correction: Acids produce hydrogen ions (H⁺), not hydrogen gas (H₂). Hydrogen gas is only produced when acids react with metals.
Mistake: Stating that weak acids are dilute and strong acids are concentrated. Correction: Strength refers to ionisation (how much the acid dissociates); concentration refers to the amount of acid dissolved in a given volume. You can have dilute solutions of strong acids and concentrated solutions of weak acids.
Mistake: Forgetting to include water as a product in neutralisation equations. Correction: All neutralisation reactions involving acids and bases/alkalis produce water. The general equation is acid + base → salt + water.
Mistake: Incorrectly naming salts, such as calling the product of sodium hydroxide and sulfuric acid "sodium sulfide" instead of sodium sulfate. Correction: Learn the three main acids: hydrochloric acid makes chlorides, sulfuric acid makes sulfates, nitric acid makes nitrates.
Mistake: In salt preparation, not using excess solid base or not filtering the mixture. Correction: Excess insoluble base ensures all acid is used, preventing contamination of crystals with acid. Filtering removes the unreacted excess base. Without these steps, the salt will not be pure.
Exam technique for Acids and Bases
Command word "Describe": When describing a test (for hydrogen, carbon dioxide, or a neutralisation reaction), you must state what you do AND what you observe. For example, "add limewater" scores 0 marks; "add limewater, which turns milky/cloudy" scores full marks.
Writing equations: Check whether the question asks for a word equation, symbol equation, or ionic equation. For neutralisation, the ionic equation H⁺(aq) + OH⁻(aq) → H₂O(l) is frequently required. Always include state symbols when the question specifies them, typically worth 1 mark.
Explaining pH differences: When comparing strong and weak acids, use the terms "ionisation," "dissociation," or "concentration of hydrogen ions." Simply saying "strong acid is more acidic" earns no marks. Examiners look for explanation of the difference in ionisation.
Salt preparation methods: Six-mark questions on preparing salts require a logical sequence with key points: excess reactant, filtering, evaporating, crystallising, drying. Each step usually earns 1 mark. The question may ask why each step is necessary, so understand the purpose of each stage.
Quick revision summary
Acids produce H⁺ ions in water (pH < 7); bases neutralise acids; alkalis are soluble bases producing OH⁻ ions (pH > 7). Strong acids completely ionise; weak acids partially ionise. Neutralisation produces salt and water: H⁺(aq) + OH⁻(aq) → H₂O(l). Acid + metal oxide/hydroxide → salt + water. Acid + carbonate → salt + water + CO₂. Salt preparation uses excess insoluble base, filtration, evaporation, and crystallisation for pure crystals.