AQA A-Level Chemistry: Reactions of ions in aqueous solution (7405) – 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.

These questions cover section 3.2.6, Reactions of ions in aqueous solution, of the AQA AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for AS and A-level exams from June 2016 onwards. The whole section is marked “A-level only”, so none of it is AS content. Inorganic chemistry is listed under A-level Paper 1, and Paper 3 can assess any content. Questions 7, 8 and 11 draw on required practical 11.

Before you start, you may want the study guide or the revision notes. Question 9 uses weak-acid pH from acids and bases. The AQA A-Level Chemistry hub and printable checklist cover the rest of the course, and the free diagnostics show where to focus.

Questions

1. For each of the four metal-aqua ions in this section, give the formula of the ion and the colour of its aqueous solution. [4]

2. Explain, in terms of charge/size ratio, why [Fe(H₂O)₆]³⁺(aq) is more acidic than [Fe(H₂O)₆]²⁺(aq) of the same concentration. Include an equation for the reaction that makes the iron(III) solution acidic. [4]

3. Sodium hydroxide solution is added drop by drop, and then in excess, to separate solutions of copper(II) sulfate and aluminium sulfate.

(a) For the copper(II) solution, describe what you see and write an equation for the reaction. [3] (b) For the aluminium solution, describe what you see with a few drops and with excess, and write an equation for each stage. [4]

4. This question is about aqueous ammonia.

(a) Write an equation for the reaction of [Fe(H₂O)₆]³⁺ with ammonia, and state the role of ammonia in this reaction. [2] (b) Aqueous ammonia is added drop by drop, then in excess, to [Cu(H₂O)₆]²⁺(aq). Describe what you see at each stage, and explain the change in excess ammonia. [4]

5. Sodium carbonate solution is added to separate solutions of iron(II) sulfate and iron(III) chloride. Describe what you see in each case. Write an equation for each reaction, and explain why the products differ. [5]

6. State what is meant by amphoteric. Write two equations to show that aluminium hydroxide is amphoteric. [3]

7. Three unlabelled solutions, A, B and C, each contain one of Fe²⁺, Fe³⁺ or Al³⁺ (as aqua ions). A student adds sodium carbonate solution to a sample of each:

Solution	Observation with Na₂CO₃(aq)
A	white precipitate and effervescence
B	green precipitate, no effervescence
C	brown precipitate and effervescence

Identify A, B and C, giving one piece of evidence for each. Suggest one further test that would confirm A, with its result. [6]

8. In required practical 11, a student tests an unknown solution with aqueous sodium hydroxide.

(a) Explain why the reagent should be added drop by drop rather than all at once. [1] (b) Explain why a fresh sample should be used for each reagent. [1] (c) State one hazard of a reagent used in this practical and a precaution to reduce the risk. [2]

9. A solution of [Fe(H₂O)₆]³⁺ has a concentration of 0.150 mol dm⁻³. Treat it as a weak monoprotic acid with Ka = 6.0 × 10⁻³ mol dm⁻³, and use the usual weak-acid approximation. Calculate the pH. [3]

10. 25.0 cm³ of 0.200 mol dm⁻³ [Fe(H₂O)₆]³⁺(aq) reacts with excess sodium carbonate solution.

(a) Write the equation for the reaction. [1] (b) Calculate the mass of [Fe(H₂O)₃(OH)₃] precipitate formed and the mass of carbon dioxide released. (Ar: Fe = 55.8, C = 12.0, O = 16.0, H = 1.0) [4]

11. A white crystalline solid, X, dissolves in water to give a colourless solution. Two salts are suspected: aluminium sulfate or magnesium sulfate.

(a) Describe a test using aqueous sodium hydroxide that would distinguish between them, giving the result for each salt. [3] (b) Give a second test, using a different reagent, that would distinguish them, with the result for each salt. [3] (c) Write an equation for the reaction in your test in (b) for the aluminium salt. [1] (d) State how you could confirm that X contains sulfate ions, and give the result. [1]

12. Explain why the copper(II) hydroxide precipitate dissolves in excess aqueous ammonia but not in excess sodium hydroxide, and why the aluminium hydroxide precipitate dissolves in excess sodium hydroxide but not in excess aqueous ammonia. [4]

Answers

1. [Fe(H₂O)₆]²⁺, pale green [1]; [Cu(H₂O)₆]²⁺, blue [1]; [Fe(H₂O)₆]³⁺, yellow-brown [1]; [Al(H₂O)₆]³⁺, colourless [1] [4] Examiner insight: Each mark needs formula and colour together; “clear” for the aluminium ion does not earn the mark, because the term required is “colourless”.

2. [Fe(H₂O)₆]³⁺ + H₂O ⇌ [Fe(H₂O)₅(OH)]²⁺ + H₃O⁺ [1]. Fe³⁺ has a higher charge and smaller radius, so a greater charge/size ratio [1]. It is more polarising and attracts electron density from the oxygen of the water ligands, weakening the O–H bonds [1]. So H⁺ is released more readily and the equilibrium lies further right (lower pH for iron(III)) [1] [4] Examiner insight: Charge alone is not enough for the ratio mark; the answer must name both charge and size, then link to the weakened O–H bond.

3. (a) Pale blue precipitate [1], no change in excess [1]; [Cu(H₂O)₆]²⁺ + 2OH⁻ → [Cu(H₂O)₄(OH)₂] + 2H₂O [1] (b) A few drops: white precipitate [1]. Excess: the precipitate dissolves to give a colourless solution [1]. [Al(H₂O)₆]³⁺ + 3OH⁻ → [Al(H₂O)₃(OH)₃] + 3H₂O [1]. [Al(H₂O)₃(OH)₃] + OH⁻ → [Al(H₂O)₂(OH)₄]⁻ + H₂O [1] (or giving [Al(OH)₄]⁻ with balanced water) [7] Examiner insight: An equation that gives the neutral hydroxide a charge, or leaves the aluminate ion uncharged, loses the equation mark even if the observations are right.

4. (a) [Fe(H₂O)₆]³⁺ + 3NH₃ → [Fe(H₂O)₃(OH)₃] + 3NH₄⁺ [1]; ammonia acts as a (Brønsted–Lowry) base, accepting H⁺ [1] (b) A few drops: pale blue precipitate [1]. Excess: the precipitate dissolves to give a deep blue solution [1]. NH₃ now acts as a ligand [1], replacing water and hydroxide ligands to form [Cu(NH₃)₄(H₂O)₂]²⁺ [1] [6] Examiner insight: Both stages are observations in their own right; an answer giving only “deep blue solution” misses the precipitate mark.

5. Iron(II): green precipitate [1]; [Fe(H₂O)₆]²⁺ + CO₃²⁻ → FeCO₃ + 6H₂O [1]. Iron(III): brown precipitate and effervescence [1]; 2[Fe(H₂O)₆]³⁺ + 3CO₃²⁻ → 2[Fe(H₂O)₃(OH)₃] + 3CO₂ + 3H₂O [1]. [Fe(H₂O)₆]³⁺ is more acidic (higher charge/size ratio), so it gives H⁺ to carbonate ions, releasing CO₂; [Fe(H₂O)₆]²⁺ is too weakly acidic, so the carbonate precipitates [1] [5] Examiner insight: Writing Fe₂(CO₃)₃ as a product is a chemical error and loses the equation mark, even if the rest is balanced.

6. Amphoteric: reacts with (dissolves in) both acids and bases [1]. [Al(H₂O)₃(OH)₃] + 3H⁺ → [Al(H₂O)₆]³⁺ [1]. [Al(H₂O)₃(OH)₃] + OH⁻ → [Al(H₂O)₂(OH)₄]⁻ + H₂O [1] [3] Examiner insight: Two equations are needed, one with acid and one with base; one equation cannot show amphoteric behaviour.

7. A is Al³⁺ [1]: white precipitate and effervescence, so a colourless 3+ ion [1]. B is Fe²⁺ [1]: green precipitate (FeCO₃) with no gas, so a 2+ ion [1]. C is Fe³⁺: brown precipitate and effervescence, so iron(III) [1]. Confirm A: add NaOH(aq) dropwise then in excess; a white precipitate forms and dissolves in excess [1] [6] Examiner insight: Identification marks usually need the supporting observation; a bare list “A = Al³⁺, B = Fe²⁺, C = Fe³⁺” with no evidence gains only the identification marks in this scheme.

8. (a) So that the precipitate can be seen forming before excess reagent is present (it may dissolve in excess) [1] (b) So that products of one reagent do not interfere with the observations for the next [1] (c) Sodium hydroxide is corrosive (or ammonia is an irritant) [1]; wear eye protection (or use ammonia in a well-ventilated room) [1] [4] Examiner insight: The precaution must match the hazard named; “wear a lab coat” for an irritant vapour does not answer the risk.

9. [H⁺]² = Ka × [HA] = 6.0 × 10⁻³ × 0.150 = 9.0 × 10⁻⁴ [1]. [H⁺] = 0.0300 mol dm⁻³ [1]. pH = −log 0.0300 = 1.52 [1] [3] Examiner insight: Give pH to two decimal places, and remember to square-root Ka × [HA]; taking −log of 9.0 × 10⁻⁴ itself gives 3.05, which is wrong.

10. (a) 2[Fe(H₂O)₆]³⁺ + 3CO₃²⁻ → 2[Fe(H₂O)₃(OH)₃] + 3CO₂ + 3H₂O [1] (b) n(Fe³⁺) = 0.200 × 25.0/1000 = 5.00 × 10⁻³ mol [1]. Mr of [Fe(H₂O)₃(OH)₃] = 55.8 + 3(18.0) + 3(17.0) = 160.8; mass = 0.804 g [1]. n(CO₂) = 5.00 × 10⁻³ × 3/2 = 7.50 × 10⁻³ mol [1]. Mass of CO₂ = 7.50 × 10⁻³ × 44.0 = 0.330 g [1] [5] Examiner insight: The 2 : 3 ratio from the equation is the usual point of loss; using 1 : 1 for CO₂ gives 0.220 g and loses that mark in this scheme.

11. (a) Add NaOH(aq) drop by drop, then in excess [1]. Both give a white precipitate at first [1]; only the aluminium precipitate dissolves in excess (magnesium hydroxide does not) [1] (b) Add sodium carbonate solution [1]. Aluminium sulfate: white precipitate with effervescence [1]. Magnesium sulfate: white precipitate with no effervescence [1] (c) 2[Al(H₂O)₆]³⁺ + 3CO₃²⁻ → 2[Al(H₂O)₃(OH)₃] + 3CO₂ + 3H₂O [1] (d) Add acidified barium chloride solution: a white precipitate shows sulfate [1] [8] Examiner insight: In a distinguishing test, a result is needed for both substances; stating only what happens with the aluminium salt loses a mark.

12. With excess NH₃, ammonia acts as a ligand and replaces water/hydroxide ligands on copper, forming the soluble [Cu(NH₃)₄(H₂O)₂]²⁺ [1]. Hydroxide ions do not form a soluble complex with copper here, so the precipitate remains in excess NaOH [1]. Aluminium hydroxide is amphoteric: excess OH⁻ removes a further H⁺ to give the soluble [Al(H₂O)₂(OH)₄]⁻ [1]. Ammonia is a weak base, so it does not remove this further H⁺, and aluminium hydroxide stays as a precipitate [1] [4] Examiner insight: Four separate points are needed for four marks; one sentence about “excess reagent” for both metals cannot gain them all.

Where marks are usually lost

• Missing the first observation (precipitate) when a reagent is added “drop by drop, then in excess”.
• Leaving out effervescence for 3+ ions with carbonate, or adding it for 2+ ions.
• Writing metal(III) carbonates as products.
• Charges wrong on neutral hydroxides, the aluminate ion or the copper–ammonia complex.
• Explaining acidity without both charge and size, or without the weakened O–H bond.
• Swapping which hydroxide dissolves in excess NH₃ (copper) and in excess NaOH (aluminium).
• Using 1 : 1 instead of 2 : 3 between M³⁺ and CO₂.
• Saying “clear” for “colourless”.

Next steps

• Recap with the revision notes.
• Re-read any weak section in the study guide.
• See the AQA A-Level Chemistry hub and the printable checklist.
• Try all free 10-minute diagnostics.
• Book a free trial class.

Official syllabus

AQA AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for AS and A-level exams June 2016 onwards (AQA). Section 3.2.6 Reactions of ions in aqueous solution (A-level only), including required practical 11.