Acids, Bases and Salts

What you'll learn

This topic examines the behaviour and reactions of acids, bases and alkalis, including pH measurement, neutralisation reactions and methods for preparing salts. You'll learn how to identify substances using pH, write equations for acid-base reactions, and understand the principles behind salt formation. This content directly addresses key areas of the WJEC specification and accounts for approximately 10-15% of your examination marks.

Key terms and definitions

Acid — a substance that releases hydrogen ions (H⁺) when dissolved in water, forming an aqueous solution with pH below 7.

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

Alkali — a soluble base that releases hydroxide ions (OH⁻) when dissolved in water, forming a solution with pH above 7.

pH scale — a numerical scale from 0 to 14 that measures the acidity or alkalinity of a solution; 7 is neutral, below 7 is acidic, above 7 is alkaline.

Neutralisation — the chemical reaction between an acid and a base that produces a salt and water, with pH moving towards 7.

Salt — an ionic compound formed when the hydrogen ions in an acid are replaced by metal ions or ammonium ions.

Indicator — a substance that changes colour depending on the pH of the solution it is added to.

Titration — a practical technique used to determine the exact volume of one solution needed to completely react with a known volume of another solution.

Core concepts

The pH scale and indicators

The pH scale runs from 0 to 14 and indicates the concentration of hydrogen ions in solution. Lower pH values correspond to higher concentrations of H⁺ ions.

Universal indicator shows a range of colours:

• Red/orange: pH 0-3 (strong acid)
• Yellow: pH 4-6 (weak acid)
• Green: pH 7 (neutral)
• Blue: pH 8-10 (weak alkali)
• Purple: pH 11-14 (strong alkali)

Litmus is a simpler indicator:

• Red in acidic solutions (pH < 7)
• Purple in neutral solutions (pH 7)
• Blue in alkaline solutions (pH > 7)

Phenolphthalein is another common indicator:

• Colourless in acidic and neutral solutions
• Pink in alkaline solutions

pH can be measured precisely using:

• pH meters (electronic devices)
• pH probes
• Universal indicator solution or paper
• Specific indicators like litmus or phenolphthalein

Common acids, bases and alkalis

Common acids you must know:

Hydrochloric acid: HCl

• Found in gastric juice in the stomach
• Used in cleaning metals and masonry

Sulfuric acid: H₂SO₄

• Used in car batteries
• Industrial manufacture of fertilizers and detergents

Nitric acid: HNO₃

• Used in fertilizer production
• Manufacturing explosives

Ethanoic acid: CH₃COOH

• Weak acid found in vinegar
• Used in food preservation

Common bases and alkalis:

Metal oxides are bases:

• Copper(II) oxide: CuO
• Zinc oxide: ZnO
• Magnesium oxide: MgO

Metal hydroxides can be bases or alkalis:

• Sodium hydroxide: NaOH (alkali - soluble)
• Calcium hydroxide: Ca(OH)₂ (alkali - slightly soluble)
• Copper(II) hydroxide: Cu(OH)₂ (base - insoluble)

Metal carbonates are also bases:

• Calcium carbonate: CaCO₃ (limestone, chalk, marble)
• Sodium carbonate: Na₂CO₃
• Copper(II) carbonate: CuCO₃

Neutralisation reactions

Neutralisation always involves H⁺ ions from the acid reacting with OH⁻ ions from the alkali or base to form water:

H⁺(aq) + OH⁻(aq) → H₂O(l)

Acid + Metal oxide → Salt + Water

Example: Hydrochloric acid + Copper(II) oxide

2HCl(aq) + CuO(s) → CuCl₂(aq) + H₂O(l)

Acid + Metal hydroxide → Salt + Water

Example: Sulfuric acid + Sodium hydroxide

H₂SO₄(aq) + 2NaOH(aq) → Na₂SO₄(aq) + 2H₂O(l)

Acid + Metal carbonate → Salt + Water + Carbon dioxide

Example: Nitric acid + Calcium carbonate

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

Note: Reactions with carbonates produce carbon dioxide gas, which can be tested with limewater (turns milky/cloudy).

Reactions of acids with metals

Acids react with metals above hydrogen in the reactivity series to produce a salt and hydrogen gas.

Acid + Metal → Salt + Hydrogen

Example: Hydrochloric acid + Magnesium

2HCl(aq) + Mg(s) → MgCl₂(aq) + H₂(g)

Example: Sulfuric acid + Zinc

H₂SO₄(aq) + Zn(s) → ZnSO₄(aq) + H₂(g)

Testing for hydrogen gas: Use a lighted splint near the mouth of the test tube. Hydrogen burns with a squeaky pop.

Metals below hydrogen in the reactivity series (e.g., copper, silver, gold) do not react with dilute acids.

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

From the acid:

• Hydrochloric acid (HCl) → chloride salts
• Sulfuric acid (H₂SO₄) → sulfate salts
• Nitric acid (HNO₃) → nitrate salts
• Ethanoic acid (CH₃COOH) → ethanoate salts

Examples:

• Sodium + Hydrochloric acid → Sodium chloride
• Copper(II) oxide + Sulfuric acid → Copper(II) sulfate
• Zinc + Nitric acid → Zinc nitrate
• Calcium carbonate + Ethanoic acid → Calcium ethanoate

Making soluble salts

Method 1: Acid + Insoluble base/metal/carbonate

This method uses excess solid reactant to ensure all the acid reacts.

Steps:

1. Add excess insoluble base/metal/carbonate to warm dilute acid
2. Stir until no more solid dissolves and fizzing stops
3. Filter to remove excess unreacted solid
4. Heat the filtrate to evaporate some water (crystallisation)
5. Leave to cool and crystallise
6. Filter, wash with distilled water and dry the crystals

This produces pure, dry salt crystals.

Method 2: Titration (Acid + Alkali)

Use when both reactants are soluble (e.g., acid + alkali).

Steps:

1. Use a pipette to measure a fixed volume of alkali into a conical flask
2. Add a few drops of indicator
3. Add acid from a burette until the indicator changes colour (endpoint)
4. Record the volume of acid needed
5. Repeat without indicator, using the exact volumes found
6. Evaporate the solution to crystallise the salt

Making insoluble salts: Precipitation

Mix solutions of two soluble salts to produce an insoluble salt as a precipitate.

Example: Making lead(II) iodide (insoluble)

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

Steps:

1. Mix the two solutions
2. Filter to collect the precipitate
3. Wash with distilled water
4. Dry in a warm oven or between filter papers

Worked examples

Example 1: Writing balanced equations

Question: Write a balanced symbol equation for the reaction between nitric acid and magnesium oxide. (3 marks)

Answer: 2HNO₃(aq) + MgO(s) → Mg(NO₃)₂(aq) + H₂O(l)

Mark scheme:

• Correct products: Mg(NO₃)₂ and H₂O (1 mark)
• Correct balancing (2 marks)
• State symbols correct (awarded if all present)

Example 2: Naming salts and describing preparation

Question: A student wants to prepare a sample of copper(II) sulfate crystals from copper(II) oxide and dilute sulfuric acid.

a) Name the type of reaction taking place. (1 mark)

b) Describe how the student could obtain pure, dry crystals of copper(II) sulfate. (4 marks)

Answer:

a) Neutralisation (1 mark)

b)

• Add excess copper(II) oxide to warm dilute sulfuric acid and stir (1 mark)
• Filter to remove unreacted copper(II) oxide (1 mark)
• Heat/warm the filtrate to evaporate some water OR pour into an evaporating basin and heat gently (1 mark)
• Leave to cool and crystallise, then filter and dry the crystals (1 mark)

Example 3: Ionic equations

Question:

a) Write the ionic equation for the neutralisation of any acid by any alkali. (2 marks)

b) Explain why this equation represents all neutralisation reactions between acids and alkalis. (2 marks)

Answer:

a) H⁺(aq) + OH⁻(aq) → H₂O(l) (2 marks - 1 for reactants, 1 for product)

b)

• In all neutralisation reactions, hydrogen ions from the acid react with hydroxide ions from the alkali (1 mark)
• The other ions (spectator ions) remain in solution and do not participate in the reaction (1 mark)

Common mistakes and how to avoid them

• Confusing bases and alkalis: Remember that all alkalis are bases, but not all bases are alkalis. Alkalis must be soluble. Sodium hydroxide is an alkali; copper(II) oxide is a base but not an alkali because it's insoluble.

• Forgetting carbon dioxide in carbonate reactions: When acids react with carbonates, always write three products: salt, water AND carbon dioxide. Missing CO₂ loses marks.

• Incorrect salt names: The metal or ammonium comes first, then the part from the acid. Sulfuric acid makes sulfates (not sulfurides). Hydrochloric acid makes chlorides (not hydrochlorides).

• Poor practical technique descriptions: In salt preparation questions, be specific. Don't write "heat it" — specify what you're heating (the filtrate/solution). Don't write "filter it" without saying what you're removing (excess solid/unreacted base).

• Mixing up reactivity series: Only metals above hydrogen react with dilute acids to produce hydrogen gas. Copper will not react with dilute hydrochloric acid.

• Incomplete balancing: Always check both sides of your equation for each element. In 2HCl + CuO → CuCl₂ + H₂O, count: H (2=2), Cl (2=2), Cu (1=1), O (1=1).

Exam technique for "Acids, Bases and Salts"

• Command word awareness: "Describe" requires a method with clear steps in sequence. "Explain" requires reasoning using scientific knowledge. For a 4-mark "describe" question on salt preparation, give four distinct steps.

• Equation questions: State symbols are often required for full marks at GCSE level — check the question carefully. Use (s) for solids, (l) for liquids, (aq) for aqueous solutions, (g) for gases.

• Practical questions: Include safety considerations where appropriate (e.g., wearing goggles, using a fume cupboard for concentrated acids). Mention specific apparatus (burette, pipette, conical flask for titrations).

• Show your working: In calculations involving neutralisation or titration, write the equation first, then show mole calculations clearly. Examiners award method marks even if your final answer is incorrect.

Quick revision summary

Acids release H⁺ ions in solution and have pH below 7. Bases react with acids; alkalis are soluble bases releasing OH⁻ ions with pH above 7. Neutralisation produces salt and water: H⁺ + OH⁻ → H₂O. Acids react with metals (above hydrogen) to give salt + hydrogen, with carbonates to give salt + water + CO₂, and with bases/alkalis to give salt + water. Soluble salts are made by reacting acids with excess insoluble bases or by titration. Insoluble salts form by precipitation.