OxfordAQA International GCSE Chemistry: Group properties (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 3.7.1 Group properties (statements a to f); no required practical is attached to this section. The qualification is untiered, so every question suits every candidate. Use relative atomic masses Li = 7, K = 39, Mg = 24, Cl = 35.5 and Br = 80, and a molar gas volume of 24 dm³ at room temperature and pressure.

Links: group properties study guide · group properties revision notes · course hub · checklist · free diagnostics

Questions

1. This question is about Group 1.

(a) Explain why the Group 1 elements are called the alkali metals. [1] (b) State the charge on the ion that every Group 1 element forms. [1]

2. Lithium burns in oxygen to form lithium oxide.

(a) Write the formula of lithium oxide. [1] (b) Describe the appearance of lithium oxide, and of the solution formed when a lithium compound such as lithium chloride dissolves in water. [2] (c) Describe what happens to the electrons when a lithium atom forms a lithium ion. [2]

3. A teacher drops a small piece of sodium onto water in a trough, then repeats the demonstration with potassium.

(a) Give two observations for sodium. [2] (b) Give one observation that is seen with potassium but not with sodium. [1] (c) Write a balanced equation for potassium reacting with water. [2] (d) Describe a test to show the gas produced is hydrogen. [1]

4. Halogens react with metals to form metal halides.

(a) Write the formulae of: (i) iron(II) chloride, (ii) barium bromide, (iii) sodium iodide. [3] (b) Write a balanced equation for aluminium reacting with bromine. [2]

5. Use these values for four halogens.

Halogen	Melting point / °C	Boiling point / °C
Fluorine	−220	−188
Chlorine	−101	−34
Bromine	−7	59
Iodine	114	184

(a) Give the state of bromine at 70 °C and of iodine at 25 °C. [2] (b) Name the halogen that is a liquid at −50 °C. [1] (c) Describe the trend in melting point down Group 7. [1] (d) Explain this trend in terms of the molecules. [2]

6. A student has three colourless solutions, P, Q and R. One is potassium chloride, one is potassium bromide and one is potassium iodide. She adds chlorine water to a sample of each, then bromine water to fresh samples.

• P: no change with chlorine water; no reaction with bromine water.
• Q: turns brown with chlorine water and with bromine water.
• R: turns yellow-orange with chlorine water; no reaction with bromine water.

(a) Identify P, Q and R. [3] (b) Explain why P does not react with bromine water. [1] (c) Write the ionic equation for the reaction of Q with bromine water. [2]

7. Explain why caesium reacts more vigorously with water than sodium does. [4]

8. Fluorine is at the top of Group 7 and iodine is below bromine.

(a) Predict whether iodine reacts with sodium fluoride solution. Give a reason. [1] (b) Explain, in terms of electrons and energy levels, why fluorine is more reactive than iodine. [4]

9. 0.78 g of potassium reacts completely with water. The solution is made up to 250 cm³ with water.

(a) Calculate the amount, in moles, of potassium used. [1] (b) Calculate the volume of hydrogen produced at room temperature and pressure. Give your answer in cm³. [2] (c) Calculate the concentration of potassium hydroxide in the final solution, in mol/dm³. [2] (d) State the colour of universal indicator in this solution. [1]

10. Chlorine gas is bubbled into a solution containing 0.92 g of magnesium bromide, MgBr₂.

(a) Explain why a reaction takes place. [1] (b) Write a balanced equation for the reaction. [2] (c) State what you would see. [1] (d) Calculate the maximum mass of bromine that can form. [3] (e) Calculate the minimum volume of chlorine needed at room temperature and pressure, in cm³. [1]

11. 0.70 g of lithium reacts completely with chlorine.

(a) Describe how the ions in lithium chloride form from lithium and chlorine atoms. Give the charge on each ion. [3] (b) Give the balanced symbol equation for this reaction. [1] (c) Calculate the mass of lithium chloride formed. [2] (d) The lithium chloride is dissolved in water. Describe the solution. [1]

12. Going down Group 1 the elements become more reactive, but going down Group 7 they become less reactive. Explain both trends, and why the same change in atomic structure leads to opposite trends. [6]

Answers

1. (a) Their hydroxides dissolve in water to give alkaline solutions [1]. (b) +1 (a 1+ ion) [1]. Examiner insight: “They are reactive metals” does not answer (a); the name comes from the alkaline solutions formed by their hydroxides.

2. (a) Li₂O [1] (b) White solid [1]; colourless solution [1]. (c) The atom loses its one outer electron [1], giving an ion with a 1+ charge (Li⁺) [1]. Examiner insight: Say the electron is lost or transferred; “lithium shares an electron” describes covalent bonding and scores nothing here.

3. (a) Any two: floats; melts into a ball; moves around on the surface; fizzes or bubbles; gets smaller and disappears [1] [1]. (b) The hydrogen catches fire with a lilac flame (or reacts more violently) [1]. (c) All four formulae right [1]; balancing right: 2K + 2H₂O → 2KOH + H₂ [1]. (d) Hold a lighted splint near the gas; it gives a pop [1]. Examiner insight: “Bubbles of hydrogen” is a conclusion, not an observation; describe what you see, such as fizzing.

4. (a) (i) FeCl₂ [1] (ii) BaBr₂ [1] (iii) NaI [1] (b) Al, Br₂ and AlBr₃ written correctly [1]; numbers 2, 3 and 2 in place: 2Al + 3Br₂ → 2AlBr₃ [1]. Examiner insight: Check every halide is 1−; formulae such as BaBr or AlBr show the charge has not been balanced.

5. (a) Bromine at 70 °C: gas [1]. Iodine at 25 °C: solid [1]. (b) Chlorine (−50 °C lies between its melting and boiling points) [1]. (c) Melting point increases down the group [1]. (d) The molecules get larger down the group [1], so the forces between molecules are stronger and need more energy to overcome [1]. Examiner insight: For (a), compare the temperature with both the melting point and the boiling point; 70 °C is above bromine’s boiling point, so it is a gas, not a liquid.

6. (a) P: potassium chloride [1]. Q: potassium iodide [1]. R: potassium bromide [1]. (b) Chlorine sits above bromine in Group 7 and is more reactive, so bromine cannot push chloride out of the salt [1]. (c) Correct species Br₂ + I⁻ → Br⁻ + I₂ [1]; balanced Br₂ + 2I⁻ → 2Br⁻ + I₂ [1]. Examiner insight: Start from the solution that reacts with both reagents; it must contain the least reactive halide, which is iodide.

7. Caesium has more occupied energy levels than sodium, so its outer electron is in a higher energy level [1]. That electron sits further out, with more inner electrons screening it from the nuclear charge [1]. So the nucleus pulls on it less strongly [1]. So caesium loses its outer electron more easily and reacts more vigorously [1]. Examiner insight: Give every step of the chain; leaving out “weaker attraction” breaks the link between size and reactivity.

8. (a) No, because iodine is less reactive than fluorine [1]. (b) Halogens react by gaining one electron [1]. In fluorine the outer energy level is closer to the nucleus (lower energy level) [1] and less shielded [1], so the incoming electron is attracted more strongly and gained more easily [1]. Examiner insight: An answer about losing electrons applies to metals and gains no credit for a halogen explanation.

9. (a) 0.78 ÷ 39 = 0.020 mol [1] (b) Moles of H₂ = 0.020 ÷ 2 = 0.010 mol [1]; volume = 0.010 × 24 000 = 240 cm³ [1]. (c) Moles of KOH = 0.020 mol; 250 cm³ = 0.250 dm³ [1]; 0.020 ÷ 0.250 = 0.080 mol/dm³ [1]. (d) Blue or purple [1]. Examiner insight: Convert 250 cm³ to 0.250 dm³ before dividing; dividing by 250 gives an answer one thousand times too small.

10. (a) Chlorine is more reactive than bromine, so it displaces bromine from its salt [1]. (b) Correct reactant formulae Cl₂ and MgBr₂ [1]; correct product formulae MgCl₂ and Br₂, giving Cl₂ + MgBr₂ → MgCl₂ + Br₂ (already balanced) [1]. (c) The colourless solution turns yellow-orange [1]. (d) Mr MgBr₂ = 24 + (2 × 80) = 184 [1]; moles = 0.92 ÷ 184 = 0.0050 mol, and 1 mol MgBr₂ gives 1 mol Br₂ [1]; mass = 0.0050 × 160 = 0.80 g [1]. (e) 0.0050 × 24 000 = 120 cm³ [1]. Examiner insight: Use Mr = 160 for Br₂, not 80; bromine is formed as diatomic molecules, and that slip halves the answer.

11. (a) Each lithium atom loses one electron [1] and each chlorine atom gains one electron [1], forming Li⁺ and Cl⁻ ions [1]. (b) 2Li + Cl₂ → 2LiCl [1] (c) Moles Li = 0.70 ÷ 7 = 0.10 mol, so 0.10 mol LiCl [1]; mass = 0.10 × 42.5 = 4.25 g [1]. (d) Colourless [1]. Examiner insight: For (d) do not describe the solution as alkaline; lithium chloride is a salt, and only the hydroxide gives an alkaline solution.

12. Indicative points, one mark each, up to 6:

• Down both groups the outer electrons are in higher energy levels (more shells) [1].
• They are further from the nucleus and more shielded, so attraction to the nucleus is weaker [1].
• Group 1 atoms react by losing their one outer electron [1].
• Weaker attraction means this electron is lost more easily, so Group 1 reactivity increases [1].
• Group 7 atoms react by gaining one electron into the outer energy level [1].
• Weaker attraction means an electron is gained less easily, so Group 7 reactivity decreases; the trends are opposite because one group loses and the other gains electrons [1]. Examiner insight: A strong answer treats the two groups separately and states whether each loses or gains electrons; one merged sentence rarely covers six distinct points.

Where marks are usually lost

• Writing Cl, Br or I as single atoms in equations, which makes balancing impossible.
• Calling the gas from Group 1 and water oxygen, or saying the metal dissolves.
• Explaining Group 7 reactivity in terms of losing electrons.
• Saying “bigger atoms” without distance, shielding and weaker attraction.
• Using an atomic mass for a diatomic element (Br = 80 instead of Br₂ = 160).
• Forgetting the cm³ to dm³ conversion in concentration and gas-volume answers.
• Writing an equation for a displacement that does not happen.
• Describing a Group 1 salt solution as alkaline; only the hydroxides are.

Next steps

• Recap fast with the group properties revision notes.
• Re-learn anything you missed in the group properties study guide.
• Practise more mole questions with the quantitative chemistry practice set.
• See the whole course on the OxfordAQA IGCSE Chemistry hub and tick off topics on the printable checklist.
• Try one of the 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 3.7.1 Group properties.