OxfordAQA International GCSE Chemistry: Identification of ions (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.

These questions are for OxfordAQA International GCSE Chemistry (9202) and follow the OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards. They cover section 3.4.3 Identification of ions (statements a to f) and the flame test required practical. There are no tiers, so you should be able to attempt all of it. Relative atomic masses are given in the questions that need them.

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Questions

1. State the flame test colour given by a compound of:

(a) barium [1] (b) lithium [1] (c) potassium [1]

2. Describe how to carry out a flame test to identify the metal ion in a small sample of an unknown solid. [4]

3. A student flame-tests a sample labelled “potassium chloride” and sees only a strong yellow flame. She had used the same wire loop, without cleaning it, for a sodium salt a moment earlier.

(a) Suggest why the flame was yellow. [1] (b) Give two changes to her method that would make the result reliable. [2]

4. Sodium hydroxide solution is added drop by drop, then in excess, to three solutions.

Solution	Observation
J	blue precipitate, which stays in excess
K	white precipitate, which stays in excess
L	green precipitate, which stays in excess

(a) Identify the metal ion in J and in L. [2] (b) K could contain either of two metal ions. Name both, and describe a test that would decide between them. [3]

5. Scale scraped from the inside of a kettle is thought to contain a carbonate.

(a) Describe a test for carbonate ions, giving the result you would expect. [3] (b) Name the white solid that forms in the test. [1]

6. Separate solutions of lithium iodide, barium chloride and calcium bromide are each acidified with dilute nitric acid. Silver nitrate solution is then added to each.

(a) Give the colour of the precipitate in each tube. [3] (b) Explain why an acid is added before the silver nitrate. [1] (c) Explain why hydrochloric acid would not be a suitable acid for this test. [1]

7. A packet of bath salts is labelled as magnesium sulfate.

(a) Describe a test to show that sulfate ions are present, and give the result. [2] (b) State why dilute hydrochloric acid is added in this test. [1] (c) Explain why dilute sulfuric acid could not be used in its place. [1]

8. A solution containing 1.74 g of potassium sulfate, K₂SO₄, is acidified with dilute hydrochloric acid. Excess barium chloride solution is added.

K₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s) + 2KCl(aq)

Calculate the mass of barium sulfate precipitate formed. Relative atomic masses: K = 39, S = 32, O = 16, Ba = 137. [3]

9. A solution contains 0.0250 mol of a sodium halide. Excess silver nitrate solution, acidified with dilute nitric acid, is added. The precipitate is filtered, washed with distilled water and dried. Its mass is 4.70 g.

(a) Calculate the relative formula mass of the silver halide. [2] (b) Identify the halide ion. Relative atomic masses: Ag = 108, Cl = 35.5, Br = 80, I = 127. [1] (c) State the colour of the precipitate. [1] (d) Explain why the precipitate is washed before it is dried and weighed. [1]

10. Two solids, X and Y, are tested. Fresh portions of solution are used for each test.

Test	X	Y
Flame test	green	no characteristic colour
Sodium hydroxide solution, then excess	no precipitate	white precipitate, dissolves in excess
Dilute nitric acid, then silver nitrate solution	cream precipitate	white precipitate

(a) Name solid X and write its formula. [3] (b) Name solid Y and write its formula. [3] (c) Write a balanced equation, with state symbols, for the reaction when a little sodium hydroxide solution is added to a solution of Y. [2]

11. A technician has four unlabelled bottles. Each holds one of these white solids: lithium carbonate, calcium chloride, magnesium sulfate, sodium iodide.

(a) Describe a series of tests that would identify the contents of each bottle. Give the expected result for each solid. [6] (b) 0.555 g of the lithium carbonate reacts completely with excess dilute hydrochloric acid:

Li₂CO₃(s) + 2HCl(aq) → 2LiCl(aq) + H₂O(l) + CO₂(g)

Relative atomic masses: Li = 7, C = 12, O = 16. Take the molar gas volume at room temperature and pressure as 24 dm³. Calculate, in cm³, the volume of gas released. [3]

12. A sample of well water may contain chloride ions and sulfate ions. A student adds dilute hydrochloric acid and barium chloride solution to a portion of the water and sees a white precipitate. She then adds dilute nitric acid and silver nitrate solution to the same test tube, sees a further white precipitate, and concludes that chloride ions are present in the water.

(a) Explain why her conclusion about chloride ions is not valid. [2] (b) Describe how she should test the water for chloride ions instead. [2]

Answers

1. (a) Green [1] (b) Crimson [1] (c) Lilac [1] Examiner insight: Use the exact colour words; “red” describes calcium, so write crimson for lithium to make the answer unambiguous.

2. Clean a nichrome wire loop: dip it in dilute hydrochloric acid and heat it, repeating until it no longer colours the flame [1]. Dip the loop in the acid, then into the solid so some sticks to it [1]. Hold the loop in the edge of a blue (non-luminous) Bunsen flame [1]. Observe the flame colour and match it to a known colour, for example crimson for lithium [1]. Examiner insight: The cleaning step is the point most often missing; say how you check the wire is clean (no colour in the flame), not just “clean it”.

3. (a) Sodium compound left on the uncleaned wire gave a yellow flame, which hid the pale lilac colour of potassium [1]. (b) Clean the wire in acid and heat it until it gives no colour before each test [1]. Use a separate loop for each sample, or repeat the test to check that the colour is the same [1]. Examiner insight: “Be more careful” is too vague to credit; each improvement should name a specific action that removes the sodium contamination.

4. (a) J: copper(II), Cu²⁺ [1]. L: iron(II), Fe²⁺ [1]. (b) Calcium or magnesium ions [1]. Carry out a flame test on the solid [1]: a red flame shows calcium; no characteristic colour shows magnesium [1]. Examiner insight: Aluminium is ruled out because its precipitate would dissolve; more sodium hydroxide cannot separate the other two.

5. (a) Add dilute hydrochloric acid to the scale [1]. Bubbles of gas (fizzing) are seen [1]. Pass the gas through limewater: it turns cloudy white, showing carbon dioxide [1]. (b) Calcium carbonate [1]. Examiner insight: The limewater result is what proves the gas is carbon dioxide; “it fizzes” on its own does not identify a carbonate.

6. (a) Lithium iodide: yellow [1]. Barium chloride: white [1]. Calcium bromide: cream [1]. (b) Carbonate ions would react with silver ions too and give a misleading precipitate; the acid destroys them first [1]. (c) Hydrochloric acid brings in its own chloride ions, so white silver chloride would appear whatever the sample contained [1]. Examiner insight: The precipitate colour depends only on the halide ion; the metal in each compound is a distractor and does not change the answer.

7. (a) Dissolve some bath salts, add dilute hydrochloric acid, followed by barium chloride solution [1]. Result: a white precipitate [1]. (b) It reacts with any carbonate ions, which could otherwise form white barium carbonate and be mistaken for a positive result [1]. (c) Sulfuric acid contains sulfate ions, so a white precipitate would form even if the bath salts had none [1]. Examiner insight: Name both reagents in the right order; a description that gives barium chloride alone leaves out the acid that makes the result trustworthy.

8. Mr of K₂SO₄ = 2(39) + 32 + 4(16) = 174, so moles = 1.74 ÷ 174 = 0.0100 mol [1]. The mole ratio is 1 : 1, so 0.0100 mol of BaSO₄ forms; Mr of BaSO₄ = 137 + 32 + 64 = 233 [1]. Mass = 0.0100 × 233 = 2.33 g [1]. Examiner insight: Write the 1 : 1 ratio from the equation; set-out working is far easier to credit than a bare number.

9. (a) Mr = mass ÷ moles [1] = 4.70 ÷ 0.0250 = 188 [1]. (b) 188 − 108 = 80, which is bromine, so the ion is bromide, Br⁻ [1]. (c) Cream [1]. (d) To remove soluble substances (such as sodium nitrate and excess silver nitrate) that would add to the mass [1]. Examiner insight: State that one mole of sodium halide gives one mole of silver halide; that is why the precipitate is 0.0250 mol.

10. (a) Green flame: barium ion, Ba²⁺ [1]. Cream precipitate: bromide ion, Br⁻ [1]. Barium bromide, BaBr₂ [1]. (b) White hydroxide that dissolves in excess: aluminium ion, Al³⁺ [1]. White silver halide: chloride ion, Cl⁻ [1]. Aluminium chloride, AlCl₃ [1]. (c) AlCl₃(aq) + 3NaOH(aq) → Al(OH)₃(s) + 3NaCl(aq): correct formulae [1]; balanced with correct state symbols [1]. Examiner insight: Give the evidence for each ion as well as the name; the formula must balance the charges, so BaBr and AlCl are not acceptable.

11. (a) One valid plan (other correct plans gain credit): Flame test on each solid. Crimson flame: lithium carbonate [1]. Red flame: calcium chloride [1]. Yellow flame: sodium iodide [1]. No characteristic colour: magnesium sulfate [1]. Confirm the negative ions on fresh portions: lithium carbonate fizzes with dilute acid and the gas turns limewater cloudy white [1]; magnesium sulfate solution, acidified with dilute hydrochloric acid, turns cloudy white when barium chloride is added, or sodium iodide gives a yellow precipitate with dilute nitric acid and silver nitrate [1]. (b) Mr of Li₂CO₃ = 2(7) + 12 + 3(16) = 74; moles = 0.555 ÷ 74 = 0.00750 mol [1]. Moles of CO₂ = 0.00750 mol (1 : 1 ratio) [1]. Volume = 0.00750 × 24 000 = 180 cm³ [1]. Examiner insight: In a planning answer, tie every test to a named bottle and its expected result; a bare list of tests identifies nothing.

12. (a) The barium chloride and hydrochloric acid she added both contain chloride ions [1]. These chloride ions would form a white silver chloride precipitate even if the water had none [1]. (b) Take a fresh portion of the water [1]. Add dilute nitric acid, then silver nitrate solution; a white precipitate shows chloride ions [1]. Examiner insight: “Use a new sample” is only half the answer; a full answer also gives the correct reagents, in order, and the expected observation.

Where marks are usually lost

• Writing “red” for lithium or “crimson” for calcium.
• Leaving out how the flame test wire is cleaned and checked.
• Calling every white hydroxide precipitate “aluminium” without the excess sodium hydroxide result.
• Using hydrochloric acid with silver nitrate, or sulfuric acid with barium chloride.
• Running a halide test on a portion that already contains chloride from earlier reagents.
• Formulae with unbalanced charges when naming the identified compound.

Next steps

• Recap fast with the identification of ions revision notes.
• Re-learn anything you missed in the identification of ions study guide.
• Practise more moles questions with the quantitative chemistry practice set.
• Browse every unit on the OxfordAQA IGCSE Chemistry hub and mark your progress on the printable checklist.
• Spot weak areas quickly with a free 10-minute diagnostic.
• 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. Scope of this set: section 3.4.3 Identification of ions, with the flame test required practical from section 6.2.