Acids, bases and salts: definitions and reactions

What you'll learn

This topic covers the fundamental chemistry of acids, bases and salts, including precise definitions, characteristic reactions and practical methods of salt preparation. Understanding these concepts is essential for CIE IGCSE Chemistry, appearing in both Paper 2 (structured questions) and Paper 4 (extended response), typically accounting for 8-12% of total marks across the papers.

Key terms and definitions

Acid — a substance that produces hydrogen ions (H⁺) when dissolved in water, or acts as a proton donor.

Base — a substance that neutralises an acid to form a salt and water only; metal oxides and metal hydroxides are bases.

Alkali — a soluble base that produces hydroxide ions (OH⁻) when dissolved in water.

Salt — an ionic compound formed when the hydrogen ion of an acid is replaced by a metal ion or ammonium ion.

Neutralisation — the reaction between an acid and a base to produce a salt and water, characterised by H⁺(aq) + OH⁻(aq) → H₂O(l).

pH — a measure of the concentration of hydrogen ions in a solution on a scale from 0 to 14, where pH 7 is neutral.

Indicator — a substance that shows different colours in acidic and alkaline solutions, used to determine pH.

Hydrated salt — a salt containing water of crystallisation chemically bonded in its crystal structure.

Core concepts

Properties and identification of acids and bases

Acids share several characteristic properties:

• Turn blue litmus paper red
• Have pH values less than 7
• React with metals, carbonates and bases
• Conduct electricity when dissolved in water due to mobile ions
• Taste sour (never taste laboratory chemicals)

Common laboratory acids include:

• Hydrochloric acid (HCl) — strong acid, completely ionised in solution
• Sulfuric acid (H₂SO₄) — strong dibasic acid
• Nitric acid (HNO₃) — strong acid, used in making fertilisers
• Ethanoic acid (CH₃COOH) — weak acid, partially ionised in solution

Bases and alkalis have distinct properties:

• Turn red litmus paper blue (alkalis)
• Have pH values greater than 7 (alkalis in solution)
• Feel soapy to touch (never handle concentrated alkalis)
• Neutralise acids

Common bases and alkalis:

• Sodium hydroxide (NaOH) — strong alkali, completely ionised
• Calcium hydroxide (Ca(OH)₂) — sparingly soluble base, forms limewater
• Ammonia solution (NH₃(aq)) — weak alkali
• Copper(II) oxide (CuO) — insoluble base
• Zinc oxide (ZnO) — amphoteric oxide (reacts with both acids and alkalis)

The pH scale and indicators

The pH scale ranges from 0 to 14:

• pH 0-6: acidic solutions (lower pH = more acidic)
• pH 7: neutral solutions (pure water, neutral salts)
• pH 8-14: alkaline solutions (higher pH = more alkaline)

Universal indicator shows a range of colours:

• Red/orange: pH 0-4 (strongly acidic)
• Yellow: pH 5-6 (weakly acidic)
• Green: pH 7 (neutral)
• Blue: pH 8-10 (weakly alkaline)
• Purple: pH 11-14 (strongly alkaline)

Litmus is a simpler indicator:

• Red in acids (pH < 7)
• Blue in alkalis (pH > 7)
• Purple in neutral solutions (pH 7)

Other indicators include phenolphthalein (colourless in acid, pink in alkali) and methyl orange (red in acid, yellow in alkali).

Reactions of acids

1. Acids with metals

Reactive metals (above hydrogen in the reactivity series) react with dilute acids:

acid + metal → salt + hydrogen

Examples:

• Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
• Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)
• 2Al(s) + 6HNO₃(aq) → 2Al(NO₃)₃(aq) + 3H₂(g)

The hydrogen gas produced can be tested with a lighted splint, which produces a 'squeaky pop' sound. Metals below hydrogen (copper, silver, gold) do not react with dilute acids.

2. Acids with metal oxides and hydroxides (bases)

This is a neutralisation reaction:

acid + base → salt + water

Examples:

• CuO(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l)
• 2HCl(aq) + Ca(OH)₂(aq) → CaCl₂(aq) + 2H₂O(l)
• 3H₂SO₄(aq) + Al₂O₃(s) → Al₂(SO₄)₃(aq) + 3H₂O(l)

The ionic equation for all acid-alkali neutralisations:

H⁺(aq) + OH⁻(aq) → H₂O(l)

3. Acids with carbonates

Carbonates and hydrogencarbonates react with acids to produce carbon dioxide:

acid + carbonate → salt + water + carbon dioxide

Examples:

• CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)
• Na₂CO₃(aq) + H₂SO₄(aq) → Na₂SO₄(aq) + H₂O(l) + CO₂(g)
• NaHCO₃(s) + HNO₃(aq) → NaNO₃(aq) + H₂O(l) + CO₂(g)

Carbon dioxide is tested using limewater (calcium hydroxide solution), which turns milky/cloudy white due to the formation of calcium carbonate precipitate.

4. Acids with ammonia

Ammonia solution (an alkali) neutralises acids without producing water in the typical sense:

acid + ammonia → ammonium salt

Examples:

• HCl(aq) + NH₃(aq) → NH₄Cl(aq)
• H₂SO₄(aq) + 2NH₃(aq) → (NH₄)₂SO₄(aq)
• HNO₃(aq) + NH₃(aq) → NH₄NO₃(aq)

Ammonium salts are important as fertilisers because they provide nitrogen for plant growth.

Preparing soluble salts

Method 1: Excess insoluble reactant (acid + metal/base/carbonate)

Used when the reactant is insoluble (most metal oxides, hydroxides, carbonates and some metals):

1. Add excess solid reactant to dilute acid in a beaker
2. Stir until no more reactant dissolves (acid is fully neutralised)
3. Filter to remove excess solid reactant
4. Evaporate the filtrate to reduce volume and concentrate the solution
5. Leave to crystallise (or cool to room temperature for faster crystallisation)
6. Filter to collect crystals, wash with distilled water
7. Dry the crystals between filter papers or in a warm oven

Example: Preparing copper(II) sulfate crystals:

• Warm dilute sulfuric acid in a beaker
• Add excess copper(II) oxide powder, stirring until no more dissolves
• Filter to remove excess black copper(II) oxide
• Evaporate the blue filtrate until crystals start to form at the surface
• Leave to crystallise, producing blue copper(II) sulfate crystals (CuSO₄·5H₂O)

Method 2: Titration (acid + alkali)

Used when both reactants are soluble (preparing salts from acids and alkalis):

1. Measure a fixed volume of alkali using a pipette into a conical flask
2. Add a few drops of indicator (phenolphthalein or methyl orange)
3. Add acid from a burette until the indicator changes colour (end point)
4. Record the volume of acid used
5. Repeat without indicator using the exact volumes determined
6. Evaporate the salt solution to crystallise the salt

Example: Preparing sodium chloride:

• Pipette 25.0 cm³ of sodium hydroxide solution into a flask
• Titrate with hydrochloric acid using phenolphthalein (pink → colourless)
• If 23.5 cm³ of acid is needed, repeat with 25.0 cm³ NaOH and exactly 23.5 cm³ HCl
• Evaporate to obtain white sodium chloride crystals

Preparing insoluble salts

Insoluble salts are prepared by precipitation reactions using two soluble compounds:

1. Mix solutions of two soluble salts that contain the required ions
2. The insoluble salt precipitates immediately
3. Filter to collect the precipitate
4. Wash with distilled water to remove soluble impurities
5. Dry the precipitate in a warm oven or between filter papers

Example: Preparing lead(II) iodide (yellow precipitate):

Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)

Ionic equation: Pb²⁺(aq) + 2I⁻(aq) → PbI₂(s)

Other common insoluble salts tested in CIE IGCSE:

• Silver chloride (white): AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)
• Barium sulfate (white): BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)
• Calcium carbonate (white): CaCl₂(aq) + Na₂CO₃(aq) → CaCO₃(s) + 2NaCl(aq)

Naming salts

The name of a salt has two parts:

1. First part: comes from the metal (or ammonium)
2. Second part: comes from the acid

Acid origins:

• Hydrochloric acid (HCl) → chloride salts
• Sulfuric acid (H₂SO₄) → sulfate salts
• Nitric acid (HNO₃) → nitrate salts
• Ethanoic acid (CH₃COOH) → ethanoate salts
• Carbonic acid (H₂CO₃) → carbonate salts
• Phosphoric acid (H₃PO₄) → phosphate salts

Examples:

• Sodium + hydrochloric acid → sodium chloride
• Copper + sulfuric acid → copper sulfate
• Calcium + nitric acid → calcium nitrate
• Zinc + ethanoic acid → zinc ethanoate

Worked examples

Example 1: Describing salt preparation

Question: Zinc chloride is a soluble salt. Describe how you would prepare a pure, dry sample of zinc chloride crystals starting from zinc oxide powder and dilute hydrochloric acid. (6 marks)

Answer:

1. Add excess zinc oxide powder to dilute hydrochloric acid in a beaker (1 mark)
2. Stir the mixture until no more zinc oxide reacts/dissolves (1 mark)
3. Filter the mixture to remove excess zinc oxide (1 mark)
4. Evaporate the filtrate to reduce the volume/until crystals start to form (1 mark)
5. Leave the solution to crystallise (1 mark)
6. Filter and dry the crystals between filter paper/in a warm oven (1 mark)

Examiner note: Each practical step must be clearly stated. Common missing points include removing excess reactant and the drying method.

Example 2: Writing balanced equations

Question: (a) Write a balanced chemical equation for the reaction between calcium carbonate and nitric acid. (3 marks) (b) Write the ionic equation for the neutralisation of any acid with any alkali. (2 marks)

Answer: (a) CaCO₃(s) + 2HNO₃(aq) → Ca(NO₃)₂(aq) + H₂O(l) + CO₂(g)

• Correct formulae (1 mark)
• Correct balancing (1 mark)
• State symbols (1 mark)

(b) H⁺(aq) + OH⁻(aq) → H₂O(l)

• Correct ions and product (1 mark)
• State symbols (1 mark)

Examiner note: State symbols are frequently required for full marks. Remember that neutralisation always involves H⁺ and OH⁻ ions combining to form water.

Example 3: Identifying unknown substances

Question: A student has three unlabelled solutions: dilute hydrochloric acid, sodium hydroxide solution, and sodium chloride solution. Describe tests the student could carry out to identify each solution. (4 marks)

Answer:

1. Test each solution with universal indicator/pH paper (1 mark)
2. The acid turns the indicator red/has pH less than 7 (1 mark)
3. The sodium hydroxide turns the indicator blue/purple/has pH greater than 7 (1 mark)
4. The sodium chloride is neutral/turns indicator green/has pH 7 (1 mark)

Alternative answer: Add each solution to a carbonate (e.g. calcium carbonate). The acid produces bubbles/effervescence/carbon dioxide gas; the other two solutions show no reaction.

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. Bases are substances that neutralise acids; alkalis are soluble bases that produce OH⁻ ions in solution. Copper oxide is a base but not an alkali because it's insoluble.

• Mistake: Writing "hydrogen chloride" instead of "hydrochloric acid" or confusing HCl(g) with HCl(aq). Correction: Hydrogen chloride is the gas (HCl); hydrochloric acid is the aqueous solution (HCl(aq)). Only the aqueous solution shows acidic properties because it contains H⁺ ions.

• Mistake: Stating that neutralisation produces "no products" or only produces salt. Correction: Neutralisation reactions between acids and bases always produce a salt AND water. When acids react with carbonates, three products form: salt, water AND carbon dioxide.

• Mistake: Filtering before the reaction is complete when preparing salts using excess solid. Correction: Add excess solid reactant and continue stirring until no more dissolves. Premature filtering means unreacted acid remains in the filtrate, preventing pure salt crystallisation.

• Mistake: Using indicator when preparing a salt by titration for crystallisation. Correction: Perform the titration WITH indicator to find the exact volumes needed, then repeat WITHOUT indicator using those exact volumes. Indicator would contaminate the salt crystals.

• Mistake: Writing incorrect salt names, especially with sulfate/sulfite or nitrate/nitrite confusion. Correction: Learn the acid-to-salt conversions: sulfuric acid → sulfate (not sulfite), nitric acid → nitrate (not nitrite). The salt name depends on the acid used, not the metal reactant.

Exam technique for "Acids, bases and salts: definitions and reactions"

• Command word "Describe": When describing salt preparation methods, structure your answer as a series of numbered steps. Include specific details like "excess", "filter", "evaporate", and "crystallise". Aim for 5-7 distinct practical steps for full marks (typically 5-6 marks allocated).

• Equation questions: State symbols are commonly required for full marks in balanced equations. If the question is worth 3 marks and you've written a correct, balanced equation, the third mark is almost certainly for state symbols: (s), (l), (g), (aq).

• "Explain" questions about pH/indicators: Always link your answer to H⁺ or OH⁻ ion concentration. For example, "The pH increases because H⁺ ions are neutralised by OH⁻ ions, reducing the hydrogen ion concentration."

• Practical questions: When asked about tests or observations, be specific. Don't write "the solution changes colour" — state what colour it changes FROM and TO. Don't write "gas is produced" — name the gas and describe how to test for it (e.g., "carbon dioxide produced, turns limewater milky").

Quick revision summary

Acids produce H⁺ ions in water, turn litmus red, and have pH < 7. Bases neutralise acids; alkalis are soluble bases producing OH⁻ ions with pH > 7. Neutralisation: H⁺ + OH⁻ → H₂O. Acids react with metals (→ salt + hydrogen), bases (→ salt + water), and carbonates (→ salt + water + CO₂). Prepare soluble salts using excess insoluble reactant or titration. Prepare insoluble salts by precipitation. Salt names: metal + acid origin (e.g., sulfuric → sulfate).