OxfordAQA International GCSE Chemistry: The properties of acids and bases, and preparation of salts (9202) – Practice Questions

These are original questions written for Marlbridge, for revision and practice on this content. They are not reproduced past-paper questions, and they do not replicate the exam’s exact structure, question count or mark tariffs – examination boards hold copyright in their own papers. Use these alongside the official past papers from your board or school.

Written for OxfordAQA International GCSE Chemistry (9202), this set follows the OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards. It covers section 3.5.1 The properties of acids and bases (statements a to g) and section 3.5.2 Preparation of salts (statements a to c). Every question here is for all candidates, as the course has no tiers.

Links: acids, bases and salts study guide · acids, bases and salts revision notes · course hub · checklist · free diagnostics

Questions

1. This question is about bases and alkalis.

(a) State the difference between a base and an alkali. [2] (b) Which one of these is an alkali: magnesium oxide, potassium hydroxide, copper(II) hydroxide? [1]

2. Name the salt formed in each reaction.

(a) iron(II) oxide and hydrochloric acid [1] (b) aluminium hydroxide and sulfuric acid [1] (c) ammonia and hydrochloric acid [1]

3. Four colourless solutions, A to D, are tested with universal indicator: A yellow, B purple, C red, D green.

(a) Describe how the student should use universal indicator to find the approximate pH of each solution. [2] (b) State which solution contains the highest concentration of hydrogen ions. [1] (c) State which solution is neutral and give its pH. [1]

4. This question is about the ions in acids and alkalis.

(a) Name the ion that makes a solution alkaline. [1] (b) Write the ionic equation for neutralisation. Include state symbols. [2] (c) Explain why sodium hydroxide with nitric acid and potassium hydroxide with hydrochloric acid have the same ionic equation. [1]

5. Dilute sulfuric acid is added from a burette to 25.0 cm³ of potassium hydroxide solution. The pH is recorded after each addition.

Volume of acid added / cm³	0	4	8	12	16	20	24
pH	13.0	12.8	12.6	12.2	7.0	1.9	1.6

(a) State the volume of acid that exactly neutralised the alkali. [1] (b) Describe how the pH changes as the acid is added and explain the change in terms of ions. [3] (c) Write a balanced equation for the reaction. [2] (d) Name the salt in the flask at 16 cm³. [1]

6. This question is about ammonia and limewater.

(a) Explain why a solution of ammonia in water turns universal indicator blue. [2] (b) Name the product of ammonia with hydrochloric acid and write a balanced equation. [2] (c) A gas is bubbled through limewater and the limewater turns cloudy. Name the solid that forms and write a balanced equation for its formation. [3]

7. A student wants to make crystals of iron(II) sulfate from iron filings.

(a) Name the acid she should use. [1] (b) Explain why this method would be unsuitable for making copper sulfate from copper, and also for making sodium sulfate from sodium. [2] (c) Write a balanced equation for the reaction of iron with the acid. [1] (d) Explain why the iron is added in excess, and state how she can tell that the reaction has finished. [2]

8. Magnesium nitrate can be made from magnesium oxide, an insoluble base.

(a) Write a balanced equation for the reaction of magnesium oxide with nitric acid. [2] (b) A student adds 3.00 g of magnesium oxide to warm nitric acid. After filtering and drying, 0.48 g of solid is left. Calculate the mass of magnesium oxide that reacted, and explain why it matters that some solid was left. [2] (c) Describe how the student gets dry crystals of magnesium nitrate from the filtrate. [3]

9. To make ammonium sulfate, a student puts 20.0 cm³ of ammonia solution and indicator in a conical flask and runs in dilute sulfuric acid from a burette.

(a) Explain why an indicator is used instead of adding an excess of one reactant and filtering. [1] (b) The volumes of acid needed in three trials were 18.10 cm³, 18.20 cm³ and 18.55 cm³. Using only the results that agree within 0.10 cm³, calculate the mean volume. [2] (c) Describe how the student then makes a solution of ammonium sulfate that contains no indicator. [2] (d) Write a balanced equation for the reaction. [1]

10. This question is about insoluble salts.

(a) Name two solutions that could be mixed to make silver chloride. [2] (b) Name two solutions that could be mixed to make barium carbonate, and write a balanced equation for the reaction. [3] (c) Explain why the precipitate is washed with distilled water after filtering. [1]

11. A technician has these chemicals: zinc hydroxide (insoluble), copper metal, copper(II) sulfate solution, sodium carbonate solution, and dilute hydrochloric, nitric and sulfuric acids.

(a) Describe how to make pure, dry crystals of zinc sulfate. Include an equation. [5] (b) Describe how to make a pure, dry sample of copper(II) carbonate, an insoluble salt. [3] (c) Explain why the copper metal cannot be used to make copper(II) chloride, and name a type of copper compound that could be used instead. [2]

12. Acidic waste water from a metal-finishing works contains dissolved copper(II) ions. Calcium hydroxide is added before release.

(a) Give two reasons for adding calcium hydroxide. [2] (b) Name the solid containing copper that forms. [1] (c) Suggest how this solid is removed from the water. [1] (d) Give one example of precipitation being used to remove unwanted ions when treating water for drinking, naming the precipitate. [2]

Answers

1. (a) A base is a metal oxide or hydroxide [1]; an alkali is a base (hydroxide) that is soluble in water [1]. (b) Potassium hydroxide [1]. Examiner insight: “An alkali neutralises acids” does not separate it from a base; solubility earns the second mark.

2. (a) Iron(II) chloride [1] (b) Aluminium sulfate [1] (c) Ammonium chloride [1] Examiner insight: “Ammonia chloride” is not the name of a salt; the ion is ammonium, and the spelling change is usually needed for the mark.

3. (a) Add a few drops of universal indicator to a sample (or use indicator paper) [1]; compare the colour with the pH colour chart [1]. (b) C (red) [1]. (c) D, pH 7 [1]. Examiner insight: A method that never mentions the colour chart is incomplete; the colour alone is not the pH.

4. (a) Hydroxide ion, OH⁻ [1]. (b) H⁺(aq) + OH⁻(aq) → H₂O(l): correct ions and product [1]; correct state symbols [1]. (c) In both, the only reaction is between H⁺ and OH⁻ ions; the other ions do not change [1]. Examiner insight: When state symbols are asked for, an equation without them usually earns only one of the two marks.

5. (a) 16 cm³ [1]. (b) The pH starts at 13 and falls slowly [1], then drops sharply through 7 at 16 cm³ to below 2 [1]. H⁺ from the acid reacts with OH⁻; once all OH⁻ is used up, extra H⁺ makes the solution acidic [1]. (c) 2KOH + H₂SO₄ → K₂SO₄ + 2H₂O: correct formulae [1]; balanced [1]. (d) Potassium sulfate [1]. Examiner insight: Quote values from the table (13, 7, 1.6) before giving the ionic reason; a general “it gets more acidic” scores little.

6. (a) Ammonia reacts with water to form hydroxide ions [1]: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, so the solution is alkaline [1]. (b) Ammonium chloride [1]; NH₃ + HCl → NH₄Cl [1]. (c) Calcium carbonate [1]. Ca(OH)₂ + CO₂ → CaCO₃ + H₂O: correct formulae [1]; balanced [1]. Examiner insight: Adding water as a product in part (b) makes the equation wrong; ammonia joins with the acid directly.

7. (a) Sulfuric acid [1]. (b) Copper is not reactive enough to react with dilute sulfuric acid [1]; sodium is too reactive and would react dangerously [1]. (c) Fe + H₂SO₄ → FeSO₄ + H₂ [1] (d) So that all the acid is used up and the solution contains only the salt [1]; bubbling (fizzing) stops and iron is left over [1]. Examiner insight: Part (b) needs one point about copper and one about sodium; “it doesn’t work” earns nothing.

8. (a) MgO + 2HNO₃ → Mg(NO₃)₂ + H₂O: correct formulae [1]; balanced [1]. (b) 3.00 − 0.48 = 2.52 g [1]. Left-over solid shows that all the acid has reacted, so no acid contaminates the salt [1]. (c) Heat the filtrate to evaporate some water until crystals form on a cold glass rod [1]; leave to cool [1]; filter off the crystals and dry between filter papers [1]. Examiner insight: “Evaporate all the water” or “heat until dry” describes something different from crystallisation and usually loses the first mark.

9. (a) Both reactants are solutions, so there is no excess solid to filter off [1]. (b) Use 18.10 and 18.20 [1]; mean = (18.10 + 18.20) ÷ 2 = 18.15 cm³ [1]. (c) Repeat with 20.0 cm³ of ammonia solution and 18.15 cm³ of acid [1], but without adding indicator [1]. (d) 2NH₃ + H₂SO₄ → (NH₄)₂SO₄ [1] Examiner insight: Including the 18.55 cm³ trial in the mean is the usual slip; give two decimal places, like the readings.

10. (a) Silver nitrate solution [1] and sodium chloride solution (or any soluble chloride) [1]. (b) Barium chloride (or barium nitrate) solution and sodium carbonate solution [1]. BaCl₂ + Na₂CO₃ → BaCO₃ + 2NaCl: correct formulae [1]; balanced [1]. (c) To remove the soluble salt (such as sodium chloride) left on the solid [1]. Examiner insight: Naming “silver” or “chlorine” in part (a) does not score; each reagent must be a named solution of a soluble compound.

11. (a) Add zinc hydroxide to warm dilute sulfuric acid, stirring, until some is left undissolved [1]. Filter off the excess solid [1]. Heat the filtrate until it is saturated [1], then cool so crystals form, filter and dry them [1]. Zn(OH)₂ + H₂SO₄ → ZnSO₄ + 2H₂O [1]. (b) Mix copper(II) sulfate solution and sodium carbonate solution [1]; filter off the precipitate [1]; wash with distilled water and dry [1]. (c) Copper is below hydrogen in the reactivity series, so it does not react with dilute hydrochloric acid [1]; use an insoluble base such as copper(II) oxide or copper(II) hydroxide [1]. Examiner insight: The excess, the filtering and the crystallising are separate creditworthy steps; leaving one out costs a mark.

12. (a) It neutralises the acid in the waste water [1]; it supplies hydroxide ions that precipitate the copper(II) ions [1]. (b) Copper(II) hydroxide [1]. (c) Let it settle and filter (or decant) the water off [1]. (d) Removing calcium ions from hard water [1] by precipitating them as calcium carbonate [1]. Examiner insight: In part (a), two separate points are needed for two marks; repeating “it neutralises” in different words scores once.

Where marks are usually lost

• Calling an insoluble oxide or hydroxide an alkali.
• Adding water to the products of ammonia and an acid.
• Forgetting state symbols in H⁺(aq) + OH⁻(aq) → H₂O(l) when they are asked for.
• Using universal indicator without the colour chart.
• Missing the “filter off the excess” step, or keeping the residue when the salt is in the filtrate.
• Naming an element or an insoluble solid as a reagent for an insoluble salt.
• Including a non-concordant trial in a mean volume.

Next steps

• Recap fast with the acids, bases and salts revision notes.
• Go back over weak spots in the acids, bases and salts study guide.
• Browse every unit from the OxfordAQA IGCSE Chemistry hub and record progress on the printable checklist.
• Spot weak topics quickly using our free 10-minute diagnostics.
• Book a free trial class.

Official syllabus

OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards, published by OxfordAQA. This practice set covers section 3.5.1 The properties of acids and bases and section 3.5.2 Preparation of salts.