Chemical changes: pH scale, neutralisation and salt production

What you'll learn

This revision guide covers the essential chemistry of acids, bases and salts required for AQA GCSE Chemistry. You'll master the pH scale, understand neutralisation reactions and learn how to predict and name the salts produced. These concepts appear frequently in both Paper 1 and Paper 2, particularly in structured questions worth 4-6 marks.

Key terms and definitions

Acid — A substance that produces hydrogen ions (H⁺) when dissolved in water; has a pH less than 7.

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

Alkali — A soluble base that produces hydroxide ions (OH⁻) when dissolved in water; has a pH greater than 7.

pH scale — A numerical scale from 0 to 14 that measures the acidity or alkalinity of a solution; 7 is neutral.

Neutralisation — A 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 whether it is in an acidic or alkaline solution.

Titration — An accurate technique used to find the exact volume of alkali needed to neutralise a known volume of acid (or vice versa).

Core concepts

The pH scale and indicators

The pH scale ranges from 0 to 14 and indicates the concentration of hydrogen ions in a solution:

• pH 0-6: Acidic solutions (the lower the pH, the more acidic)
• pH 7: Neutral solutions (pure water, neutral salt solutions)
• pH 8-14: Alkaline solutions (the higher the pH, the 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 paper provides a simpler test:

• Red litmus turns blue in alkalis (pH > 7)
• Blue litmus turns red in acids (pH < 7)
• No colour change in neutral solutions

Phenolphthalein is colourless in acids and neutral solutions, but turns pink in alkalis.

The pH scale is logarithmic — each decrease of 1 on the pH scale represents a ten-fold increase in hydrogen ion concentration. For example, pH 3 has ten times more H⁺ ions than pH 4.

Common acids and alkalis

Common laboratory acids:

• Hydrochloric acid (HCl)
• Sulfuric acid (H₂SO₄)
• Nitric acid (HNO₃)
• Ethanoic acid (CH₃COOH) — a weak acid found in vinegar

Common alkalis:

• Sodium hydroxide (NaOH)
• Potassium hydroxide (KOH)
• Calcium hydroxide (Ca(OH)₂) — also called slaked lime
• Ammonia solution (NH₃)

Acids in solution produce H⁺ ions. The reaction for hydrochloric acid is: HCl(aq) → H⁺(aq) + Cl⁻(aq)

Alkalis in solution produce OH⁻ ions. The reaction for sodium hydroxide is: NaOH(aq) → Na⁺(aq) + OH⁻(aq)

Neutralisation reactions

Neutralisation occurs when an acid reacts with a base. The general word equation is:

acid + base → salt + water

The ionic equation for all neutralisation reactions is: H⁺(aq) + OH⁻(aq) → H₂O(l)

This shows that neutralisation involves hydrogen ions from the acid combining with hydroxide ions from the alkali to form water.

Four key types of neutralisation reactions:

1. Acid + metal oxide

   • Example: sulfuric acid + copper oxide → copper sulfate + water
   • H₂SO₄(aq) + CuO(s) → CuSO₄(aq) + H₂O(l)

2. Acid + metal hydroxide (alkali)

   • Example: nitric acid + sodium hydroxide → sodium nitrate + water
   • HNO₃(aq) + NaOH(aq) → NaNO₃(aq) + H₂O(l)

3. Acid + metal carbonate

   • Example: hydrochloric acid + calcium carbonate → calcium chloride + water + carbon dioxide
   • 2HCl(aq) + CaCO₃(s) → CaCl₂(aq) + H₂O(l) + CO₂(g)
   • Note: This produces carbon dioxide gas in addition to salt and water

4. Acid + ammonia

   • Example: sulfuric acid + ammonia → ammonium sulfate
   • H₂SO₄(aq) + 2NH₃(aq) → (NH₄)₂SO₄(aq)
   • Note: This doesn't produce water because ammonia is not a metal oxide or hydroxide

Naming salts

The name of a salt has two parts:

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

From the acid:

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

Examples:

• Sodium hydroxide + hydrochloric acid → sodium chloride
• Magnesium oxide + sulfuric acid → magnesium sulfate
• Potassium hydroxide + nitric acid → potassium nitrate
• Copper carbonate + ethanoic acid → copper ethanoate

Making soluble salts

Method 1: Acid + insoluble base (metal oxide, metal hydroxide, or metal carbonate)

This is used when the base is insoluble and the salt is soluble.

Procedure:

1. Add excess solid base to warm dilute acid in a beaker
2. Stir until no more base reacts (some solid remains)
3. Filter to remove excess unreacted base
4. Heat the filtrate to evaporate some water and concentrate the solution
5. Leave to crystallise
6. Filter and dry the crystals

Example: Making copper sulfate crystals from copper oxide and sulfuric acid.

Method 2: Titration (acid + alkali)

This method is used when both the acid and alkali are soluble — you cannot use excess and filter.

Procedure:

1. Use a pipette to measure a fixed volume of alkali into a conical flask
2. Add a few drops of indicator (e.g., phenolphthalein or methyl orange)
3. Fill a burette with acid
4. Add acid slowly from the burette until the indicator changes colour (end point)
5. Record the volume of acid needed
6. Repeat without indicator using the exact volumes found
7. Evaporate water to crystallise the salt

Example: Making sodium chloride from sodium hydroxide solution and hydrochloric acid.

Making insoluble salts — precipitation

Some salts are insoluble in water. These are made by precipitation — mixing two soluble salt solutions to form an insoluble salt.

Solubility rules you need to know:

• All sodium, potassium and ammonium salts are soluble
• All nitrates are soluble
• Most chlorides are soluble (except silver chloride and lead chloride)
• Most sulfates are soluble (except barium sulfate, calcium sulfate and lead sulfate)
• Most carbonates are insoluble (except sodium, potassium and ammonium carbonates)
• Most hydroxides are insoluble (except sodium, potassium and calcium hydroxides)

Procedure for making an insoluble salt:

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

Example: To make lead iodide (insoluble):

• Mix lead nitrate solution with potassium iodide solution
• Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)
• Yellow precipitate of lead iodide forms

Worked examples

Example 1: Predicting salt names

Question: A student reacts magnesium carbonate with dilute hydrochloric acid. Name the salt produced and write a balanced symbol equation for the reaction. [3 marks]

Answer:

Salt name: magnesium chloride [1 mark]

Balanced equation: MgCO₃(s) + 2HCl(aq) → MgCl₂(aq) + H₂O(l) + CO₂(g) [2 marks]

Mark scheme notes: 1 mark for correct formula and state symbols, 1 mark for balancing

Example 2: Practical method

Question: Describe how to make pure, dry crystals of copper sulfate from copper oxide and dilute sulfuric acid. [6 marks]

Answer:

1. Add excess copper oxide to warm dilute sulfuric acid in a beaker [1 mark]
2. Stir until no more copper oxide reacts / some solid remains [1 mark]
3. Filter the mixture to remove excess copper oxide [1 mark]
4. Heat the filtrate gently / evaporate some water to make a saturated solution [1 mark]
5. Leave the solution to crystallise / cool slowly [1 mark]
6. Filter and dry the crystals / pat dry with filter paper [1 mark]

Mark scheme notes: Must include filtering twice (steps 3 and 6) and crystallisation step

Example 3: pH changes during neutralisation

Question: A student adds sodium hydroxide solution to 25 cm³ of dilute hydrochloric acid. Sketch a graph showing how the pH changes as sodium hydroxide is added. Mark the point where neutralisation is complete. [3 marks]

Answer:

Graph should show:

• pH starting around 1-2 (acidic) [1 mark]
• pH rising to 7 and then beyond to 12-13 (alkaline) [1 mark]
• Neutralisation point marked at pH 7 [1 mark]

The curve should be steepest near pH 7 where neutralisation occurs

Common mistakes and how to avoid them

• Confusing bases and alkalis: Remember that all alkalis are bases, but not all bases are alkalis. Only soluble bases are alkalis. Copper oxide is a base but not an alkali (it's insoluble).

• Getting salt names wrong: Always check which acid was used. Students often write "sodium sulfate" when hydrochloric acid was used — this should be sodium chloride. The acid determines the second part of the salt's name.

• Forgetting carbon dioxide in carbonate reactions: When an acid reacts with a carbonate, three products form: salt, water AND carbon dioxide. Don't write just salt and water.

• Incorrect practical method descriptions: When making soluble salts from insoluble bases, you must add excess solid and then filter. Students often forget to mention filtering out the excess base.

• Mixing up indicator colours: Red litmus turns blue in alkali, blue litmus turns red in acid. Remember: "red to blue = alkali true."

• Writing incorrect formulae for salts: Ammonium salts contain NH₄⁺, not NH₃. For example, ammonium chloride is NH₄Cl, not NH₃Cl.

Exam technique for "Chemical changes: pH scale, neutralisation and salt production"

• Command word "Describe a method": Include at least 4-5 distinct steps in logical order. Use numbered points. Always mention filtering, evaporating/crystallising and drying when making salts. Expect 4-6 marks for these questions.

• Predicting products: For neutralisation questions, identify the salt name (2 parts: metal + acid ending) and mention water. If a carbonate is involved, add carbon dioxide as a third product. Symbol equations are usually worth 2 marks: 1 for formulae, 1 for balancing.

• 6-mark questions on practical methods: These require detailed step-by-step procedures. State what you would see (observations), mention safety precautions if appropriate, and describe how to obtain pure, dry crystals. Structure answers in numbered points.

• pH and indicators: Know the approximate pH values (not just "acidic" or "alkaline"). Strong acids are pH 0-2, weak acids pH 3-6, neutral is pH 7, weak alkalis pH 8-10, strong alkalis pH 11-14. Universal indicator questions often ask for colours and approximate pH values.

Quick revision summary

The pH scale measures acidity (0-6), neutral (7), or alkalinity (8-14). Acids produce H⁺ ions, alkalis produce OH⁻ ions. Neutralisation reactions between acids and bases produce salts plus water. Salt names come from the metal/ammonium (first part) and the acid (second part). Make soluble salts using excess insoluble base then filtering, or by titration for soluble reactants. Make insoluble salts by precipitation. Metal carbonates react with acids to produce salt, water and carbon dioxide.