Acids, bases and salts

What you'll learn

This revision guide covers the essential content on acids, bases and salts for the CIE IGCSE Co-ordinated Science (Double Award) specification. You'll learn how to identify acids and bases using indicators, understand neutralisation reactions, prepare common salts using different methods, and test for key ions in solution. These concepts are fundamental to understanding chemical reactions and appear frequently in both Paper 1 (multiple choice) and Paper 2 (structured questions).

Key terms and definitions

Acid — a substance that produces hydrogen ions (H⁺) when dissolved in water, with a pH below 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, with a pH above 7

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

Neutralisation — the reaction between an acid and a base to produce a salt and water only

Indicator — a substance that changes colour depending on whether it is in acidic or alkaline solution

Precipitation — the formation of an insoluble solid when two solutions are mixed together

Titration — a technique used to accurately measure the volume of one solution that reacts exactly with a known volume of another solution

Core concepts

The pH scale and indicators

The pH scale measures acidity or alkalinity and runs 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)

Common laboratory acids include:

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

Common laboratory alkalis include:

• Sodium hydroxide (NaOH)
• Potassium hydroxide (KOH)
• Calcium hydroxide (Ca(OH)₂)
• Aqueous ammonia (NH₃)

Universal indicator shows different colours across the pH range:

• Red/orange: strongly acidic (pH 1-3)
• Yellow: weakly acidic (pH 4-6)
• Green: neutral (pH 7)
• Blue: weakly alkaline (pH 8-10)
• Purple: strongly alkaline (pH 11-14)

Litmus paper provides a simpler test:

• Red litmus turns blue in alkali
• Blue litmus turns red in acid
• No colour change indicates neutral solution

Neutralisation reactions

Neutralisation occurs when acids react with bases. The general equation is:

acid + base → salt + water

The ionic equation for neutralisation shows the essential reaction:

H⁺(aq) + OH⁻(aq) → H₂O(l)

Specific examples of neutralisation reactions:

Acid + metal oxide:

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

Acid + metal hydroxide:

• hydrochloric acid + sodium hydroxide → sodium chloride + water
• HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

Acid + metal carbonate:

This reaction produces a salt, water and carbon dioxide:

• 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 using limewater, which turns milky/cloudy white.

Acid + ammonia:

Ammonia solution (aqueous ammonia) is an alkali that neutralises acids:

• hydrochloric acid + ammonia → ammonium chloride
• HCl(aq) + NH₃(aq) → NH₄Cl(aq)

Preparing soluble salts

There are three main methods for preparing soluble salts:

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

This is suitable when the base is insoluble:

1. Add excess base to warm dilute acid (ensures all acid reacts)
2. Stir until no more base dissolves
3. Filter to remove excess unreacted base
4. Heat filtrate to evaporate some water (concentrating the solution)
5. Leave to crystallise
6. Filter to collect crystals and dry between filter paper

Example: Preparing copper(II) sulfate crystals from copper(II) oxide and sulfuric acid.

Method 2: Acid + alkali (soluble base)

This method requires titration to find the exact volumes needed:

1. Perform titration using indicator to find volumes that exactly neutralise
2. Repeat without indicator using exact volumes found
3. Evaporate water to concentrate solution
4. Leave to crystallise
5. Filter and dry crystals

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

Method 3: Acid + reactive metal

Only suitable for metals above hydrogen in the reactivity series (but not too reactive):

1. Add metal pieces to dilute acid
2. Allow to react until fizzing stops
3. Filter to remove excess metal
4. Crystallise as before

Example: Preparing zinc sulfate from zinc and dilute sulfuric acid.

Do not use this method with very reactive metals (sodium, potassium, calcium) as reactions are too vigorous and dangerous.

Preparing insoluble salts

Insoluble salts are prepared by precipitation:

1. Mix two suitable soluble salt solutions
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

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

• lead(II) nitrate + potassium iodide → lead(II) iodide + potassium nitrate
• Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)

Common insoluble salts include:

• Most carbonates (except sodium, potassium and ammonium)
• Most hydroxides (except sodium, potassium and calcium — slightly soluble)
• Lead(II) chloride, lead(II) iodide
• Silver chloride, silver bromide, silver iodide
• Barium sulfate

Testing for ions

Testing for metal ions (cations):

Add sodium hydroxide solution to the unknown solution. Observe the precipitate colour:

• Calcium (Ca²⁺): white precipitate
• Copper(II) (Cu²⁺): blue precipitate
• Iron(II) (Fe²⁺): green precipitate
• Iron(III) (Fe³⁺): brown/rust-coloured precipitate
• Zinc (Zn²⁺): white precipitate (dissolves in excess sodium hydroxide)
• Aluminium (Al³⁺): white precipitate (dissolves in excess sodium hydroxide)

Flame tests provide an alternative for some metal ions:

• Calcium: red/orange flame
• Copper: blue-green flame
• Sodium: yellow/orange flame
• Potassium: lilac flame

Testing for ammonium ions (NH₄⁺):

Add sodium hydroxide solution and warm gently. Ammonia gas is released:

• Test with damp red litmus paper — turns blue
• Pungent smell

Testing for carbonate ions (CO₃²⁻):

Add dilute acid. Carbon dioxide gas is produced:

• Bubbles/fizzing observed
• Test gas with limewater — turns milky/cloudy

Testing for chloride, bromide and iodide ions (halides):

Add dilute nitric acid, then silver nitrate solution:

• Chloride (Cl⁻): white precipitate of silver chloride
• Bromide (Br⁻): cream precipitate of silver bromide
• Iodide (I⁻): yellow precipitate of silver iodide

The nitric acid removes interfering ions like carbonates.

Testing for sulfate ions (SO₄²⁻):

Add dilute hydrochloric acid, then barium chloride solution:

• White precipitate of barium sulfate forms

The hydrochloric acid removes interfering ions like carbonates.

Everyday applications of acids, bases and salts

Neutralisation in everyday life:

• Indigestion remedies contain bases (antacids) like magnesium hydroxide or calcium carbonate to neutralise excess stomach acid (hydrochloric acid)
• Farmers spread lime (calcium hydroxide or calcium carbonate) on acidic soils to neutralise acid and improve growing conditions
• Toothpaste contains mild alkalis to neutralise acids produced by bacteria
• Treating insect stings: bee stings (acidic) treated with bicarbonate of soda (alkaline); wasp stings (alkaline) treated with vinegar (acidic)

Important salts and their uses:

• Sodium chloride: food preservative, making other chemicals, de-icing roads
• Ammonium nitrate: fertiliser (provides nitrogen for plant growth)
• Calcium sulfate (gypsum): making plaster of Paris for medical casts
• Silver chloride: photographic film (light-sensitive)

Worked examples

Example 1: Naming salts and writing equations

Question: Complete the word equation and write the balanced symbol equation for the reaction between nitric acid and magnesium carbonate. Name the salt produced. [4 marks]

Answer:

• Word equation: nitric acid + magnesium carbonate → magnesium nitrate + water + carbon dioxide [1]
• The salt is magnesium nitrate [1]
• Symbol equation: 2HNO₃(aq) + MgCO₃(s) → Mg(NO₃)₂(aq) + H₂O(l) + CO₂(g) [2]

Examiner note: Remember that acid + carbonate reactions always produce three products: salt, water and carbon dioxide. The salt name takes the metal from the base (magnesium) and the ending from the acid (nitrate from nitric acid).

Example 2: Identifying ions through tests

Question: A student has an unknown solution. Describe the tests and results that would confirm the solution contains copper(II) sulfate. [4 marks]

Answer:

Test for copper(II) ions:

• Add sodium hydroxide solution to the unknown solution [1]
• Blue precipitate forms, confirming copper(II) ions present [1]

Test for sulfate ions:

• Add dilute hydrochloric acid, then add barium chloride solution [1]
• White precipitate forms, confirming sulfate ions present [1]

Examiner note: You must describe both the test procedure and the positive result. State reagents clearly and use correct colour descriptions.

Example 3: Preparing a salt

Question: Describe how to prepare a dry sample of zinc sulfate crystals from zinc oxide and dilute sulfuric acid. Your method should produce pure crystals. [6 marks]

Answer:

1. Add excess zinc oxide powder to warm dilute sulfuric acid in a beaker [1]
2. Stir until no more zinc oxide dissolves/reacts [1]
3. Filter the mixture to remove excess (unreacted) zinc oxide [1]
4. Heat the filtrate to evaporate some water/concentrate the solution [1]
5. Leave the solution to cool and crystallise [1]
6. Filter to collect crystals and dry them between filter paper/in a warm oven [1]

Examiner note: Use excess base to ensure all acid is used up. Don't forget the filtering step — examiners commonly award marks for this. State "filtrate" rather than "solution" after filtering.

Common mistakes and how to avoid them

• Confusing acids and alkalis: Remember that all alkalis are bases, but not all bases are alkalis. A base is only an alkali if it dissolves in water. Metal oxides like copper(II) oxide are bases but not alkalis because they're insoluble.

• Incomplete salt names: The salt name has two parts — the metal (or ammonium) from the base, and the ending from the acid. Hydrochloric acid makes chlorides, sulfuric acid makes sulfates, nitric acid makes nitrates. Don't write just "zinc salt" — specify "zinc chloride" or "zinc sulfate".

• Wrong order in ion tests: When testing for halide ions or sulfate ions, always add the acid first (nitric acid for halides, hydrochloric acid for sulfates) to remove interfering ions, then add the second reagent (silver nitrate or barium chloride).

• Forgetting state symbols: In symbol equations, include (s) for solids, (l) for liquids, (aq) for aqueous solutions, and (g) for gases. Marks are often awarded specifically for correct state symbols.

• Using wrong indicators for titrations: Use methyl orange or phenolphthalein for titrations, not universal indicator. Universal indicator changes through many colours and makes it difficult to identify the precise end-point.

• Heating to dryness when crystallising: Don't heat the solution until all water evaporates. Heat to concentrate it, then allow crystallisation to occur slowly by cooling. Heating to dryness produces powder, not crystals, and may decompose the salt.

Exam technique for "Acids, bases and salts"

• Command word precision: "Describe" requires you to state what you do and what you observe. "Explain" requires you to give reasons using scientific knowledge. For a "describe" question about preparing a salt, list the steps in order — don't explain why you do them unless asked.

• Equation marks: Symbol equations typically award 1 mark for correct formulae and 1 mark for balancing. Write formulae correctly first (check charges on ions), then balance. Show all working — don't rub out incorrect attempts completely as you may still gain partial marks.

• Test questions structure: When identifying an unknown substance through tests, examiners expect: (1) the reagent/procedure, (2) the observation/result, (3) the conclusion. Each may be worth a separate mark, so include all three parts.

• Practical questions: For salt preparation methods, marks are typically awarded for: using excess reagent, filtering, concentrating by evaporation, crystallisation, and drying. Each step may be worth 1 mark, so don't combine steps or miss any out.

Quick revision summary

Acids have pH below 7 and produce H⁺ ions in water. Bases neutralise acids; soluble bases (alkalis) produce OH⁻ ions and have pH above 7. Neutralisation produces salts and water. Prepare soluble salts using acid + insoluble base (or metal), using titration for soluble bases. Prepare insoluble salts by precipitation. Test for cations using sodium hydroxide solution (or flame tests); test for anions using specific reagents: silver nitrate for halides, barium chloride for sulfates, acid for carbonates.