OxfordAQA A-Level Chemistry: Bonding (9620) – 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.

This set tests section 3.1.3 Bonding (3.1.3.1 to 3.1.3.7) from the OxfordAQA International AS and A-level Chemistry (9620) specification, Version 5.3, for International AS and A-level exams May/June 2020 onwards. All of it is International AS content, so no question here is International A-level only. Bonding sits in Unit 1, and Unit 2 questions may draw on it.

Learn the content first from the ionic, covalent and metallic bonding guide, the shapes, polarity and intermolecular forces guide and the bonding revision notes. See also the course hub, the printable checklist and a free 10-minute diagnostic. Approximate melting and boiling points are quoted in the questions where needed.

Questions

1. Write the formula of (a) aluminium sulfate, (b) ammonium carbonate and (c) magnesium nitride. [3]

2. Ammonia reacts with boron trifluoride, BF₃, to form H₃N→BF₃.

(a) State what is meant by a co-ordinate (dative covalent) bond. [1] (b) Explain why NH₃ can form this bond and why BF₃ can accept it. [2] (c) Describe how the bond between N and B is drawn in a displayed formula. [1] (d) A nitrogen molecule contains a triple bond. State how many electrons are shared. [1]

3. Magnesium melts at about 650 °C, far above sodium at about 98 °C.

(a) Describe the bonding in magnesium metal. [2] (b) Explain why magnesium has the higher melting point. [3]

4. Four solids, W to Z, were tested.

Solid	Melting point / °C	Conducts as solid?	Conducts when molten?	Other
W	1540	Yes	Yes	Can be hammered into sheets
X	above 3000	No	Does not melt readily	Insoluble in all solvents tried
Y	54	No	No	Soluble in hexane
Z	780	No	Yes	Dissolves in water

For each solid, give the type of structure and explain your choice using the data. [8]

5. (a) Explain why melting ice absorbs energy but freezing water releases energy. [2] (b) Sodium chloride melts at about 801 °C and iodine at about 114 °C. Explain the difference in terms of what is overcome on melting. [3]

6. Draw a labelled diagram of one layer of the sodium chloride structure containing nine ions, and name the force that holds the ions together. [4]

7. Give the shape and the bond angle (or a statement of the angle) for each species. [8]

(a) BCl₃ (b) NH₄⁺ (c) ClF₃ (d) XeF₂

8. (a) Use electron-pair repulsion theory to deduce the shape of the ICl₄⁻ ion. [3] (b) Explain why the H–O–H angle in water (104.5°) is smaller than the H–N–H angle in ammonia (107°). [3]

9. (a) Define electronegativity. [1] (b) Show that the P–Cl bond in PCl₃ is polar by adding partial charges. [1] (c) Explain why PCl₃ has a permanent dipole but SF₆, which has six polar S–F bonds, does not. [3]

10. Approximate boiling points: HF 20 °C, HCl −85 °C, HBr −67 °C, HI −35 °C.

(a) Explain the increase from HCl to HI. [2] (b) Explain why HF does not fit this trend. [3]

11. Ethanol and methoxymethane both have the formula C₂H₆O. Ethanol boils at about 78 °C and methoxymethane at about −24 °C.

(a) Name the types of intermolecular force in each liquid. [2] (b) Explain the difference in boiling point. [2] (c) Show a hydrogen bond between two ethanol molecules, including lone pair and partial charges. [2]

12. The density of liquid water at 0 °C is 1.000 g cm⁻³ and that of ice is 0.917 g cm⁻³.

(a) Calculate the percentage increase in volume when 275 g of water freezes. [3] (b) Explain why ice is less dense than liquid water. [4]

Answers

1. (a) Al³⁺ and SO₄²⁻ need two and three respectively to balance: Al₂(SO₄)₃ [1]. (b) NH₄⁺ and CO₃²⁻: (NH₄)₂CO₃ [1]. (c) Mg²⁺ and N³⁻: Mg₃N₂ [1] [3] Examiner insight: Brackets matter. A compound ion needed more than once written without brackets changes the formula and is likely to lose the credit.

2. (a) A shared pair of electrons where both electrons come from one atom. [1] (b) N in NH₃ has a lone pair to donate [1]. B in BF₃ has only six electrons in its outer shell, so it has room to accept a pair [1]. (c) An arrow pointing from N to B. [1] (d) Six electrons (three shared pairs). [1] [5] Examiner insight: Say which atom donates and which accepts, and name the lone pair. An arrow drawn from B to N shows the bond the wrong way round.

3. (a) A lattice of Mg²⁺ ions [1] attracted to a sea of delocalised electrons [1]. (b) Mg²⁺ has a higher charge than Na⁺ and each atom gives two electrons, so there are more delocalised electrons [1]. The attraction between ions and electrons is stronger [1], so more energy is needed to overcome it [1]. [5] Examiner insight: A comparison needs both substances. Mentioning only the charge on magnesium is less secure than also linking it to the number of delocalised electrons.

4. W is metallic [1]: it conducts as a solid because delocalised electrons move, and is malleable [1]. X is macromolecular (giant covalent) [1]: very high melting point from covalent bonds throughout, with no mobile charges [1]. Y is molecular [1]: low melting point as only weak intermolecular forces are overcome, and no charged particles [1]. Z is ionic [1]: conducts only when molten because ions are then free to move, and dissolves in water [1] [8] Examiner insight: A type with no reason from the data earns little. Link each property to the particles that carry charge or the forces that must be broken.

5. (a) Melting needs energy to overcome the hydrogen bonds between molecules [1]. When water freezes these bonds form and energy is released [1]. (b) Melting sodium chloride overcomes strong electrostatic attraction between oppositely charged ions [1]. Melting iodine overcomes only weak induced dipole–dipole forces between I₂ molecules, and the I–I covalent bonds are not broken [1]. Much less energy is needed, so iodine melts at a lower temperature [1] [5] Examiner insight: For a molecular solid, an answer that says covalent bonds break on melting is a common way to lose credit. Name the intermolecular force.

6. A 3 × 3 square grid [1], with ions alternating so each Na⁺ has Cl⁻ neighbours [1], labelled Na⁺ and Cl⁻ with five of one and four of the other [1]:

Na+ Cl- Na+
Cl- Na+ Cl-
Na+ Cl- Na+

The force is electrostatic attraction between oppositely charged ions. [1] [4] Examiner insight: Check the count and the alternation. Two like ions touching along a row is a common slip in a sketch.

7. (a) Trigonal planar [1], 120° [1]. (b) Tetrahedral [1], 109.5° [1]. (c) T-shaped (five pairs, two lone) [1], angle slightly less than 90° [1]. (d) Linear (five pairs, three lone) [1], 180° [1] [8] Examiner insight: A shape name alone is only half the task. Give the angle, and for lone-pair species mention that the pairs are not all bonding.

8. (a) Electrons: 7 + 4 + 1 = 12, so six pairs [1]; four bonding and two lone [1]; the lone pairs sit opposite each other, giving square planar, 90° [1]. (b) Both have four electron pairs in a tetrahedral arrangement [1]. Water has two lone pairs and ammonia has one [1]. Lone pair–lone pair and lone pair–bond pair repulsion are greater than bond pair–bond pair repulsion, so water’s bonds are squeezed closer [1] [6] Examiner insight: Put the repulsion order in your own words. Saying “lone pairs repel more” without comparing them to bonding pairs is thin.

9. (a) How strongly an atom pulls the shared pair of a covalent bond towards itself. [1] (b) P^δ+–Cl^δ−. [1] (c) PCl₃ is pyramidal, so its three dipoles do not cancel [1]. SF₆ is octahedral and symmetrical [1], so its six equal S–F dipoles cancel [1] [5] Examiner insight: A “polar bonds, therefore polar molecule” answer is incomplete. State the shape and the cancellation.

10. (a) The molecules have more electrons, so stronger induced dipole–dipole forces [1], needing more energy to overcome [1]. (b) HF has a hydrogen atom bonded to fluorine, which has a lone pair [1], so it forms hydrogen bonds between molecules [1]. These are stronger than the forces in HCl, so more energy is needed and the boiling point is higher than the trend predicts [1] [5] Examiner insight: Say what is overcome (intermolecular forces), not the H–F bond.

11. (a) Ethanol: hydrogen bonding, permanent dipole–dipole and induced dipole–dipole forces [1]. Methoxymethane: permanent dipole–dipole and induced dipole–dipole forces only [1]. (b) Ethanol has an O–H group, so its molecules hydrogen bond [1]. Methoxymethane has no hydrogen on oxygen, so less energy is needed to separate the molecules [1]. (c) A dotted line from the δ+ hydrogen of one ethanol molecule to a lone pair on the δ− oxygen of the other [1], with the O–H···O atoms in a straight line and the δ+ and δ− shown [1] [6] Examiner insight: A hydrogen bond drawn to the oxygen atom but not to a lone pair is easy to mark down. Show the lone pair.

12. (a) Liquid volume = 275 ÷ 1.000 = 275 cm³ [1]. Ice volume = 275 ÷ 0.917 = 299.9 cm³ [1]. Increase = 24.9 ÷ 275 × 100 = 9.1 % [1]. (b) In ice each molecule forms hydrogen bonds to four others [1] in an open tetrahedral lattice with gaps [1]. These hydrogen bonds hold molecules further apart than in liquid water [1]. On melting some bonds break and the molecules pack more closely [1] [7] Examiner insight: Give the working to 3 significant figures at least, and avoid saying the bonds “expand”. The lattice holds molecules apart.

Where marks are usually lost

• Writing “bonds break” for a molecular substance when the answer needs intermolecular forces.
• Stating a bond angle without the lone pair reasoning, or giving 109.5° for a molecule with lone pairs.
• Counting only the atoms around the centre and forgetting the charge of an ion when finding electron pairs.
• Saying a molecule is non-polar because the bonds are non-polar, instead of explaining cancellation.
• Missing the lone pair and the N, O or F link in hydrogen bonding answers.
• Describing a conductor without naming the particles that move.
• Using “weaker forces” without saying which forces are weaker, or comparing only one substance.
• Drawing dative bonds as plain lines or lattices with like ions side by side.

Next steps

• Revisit the bonding revision notes.
• Reread the shapes, polarity and forces guide and the bonding and structure guide.
• Check the course hub and the printable checklist.
• Take one of the free 10-minute diagnostics.
• Book a free trial class.

Frequently asked questions

Is bonding AS content in OxfordAQA 9620 Chemistry? Yes. Section 3.1.3 Bonding is International AS content. It sits in Unit 1, and Unit 2 questions may also draw on bonding principles.

What are the most common mistakes in bonding and shapes questions? Frequent slips are writing “bonds break” for a molecular substance, stating a bond angle without the lone-pair reasoning, and forgetting the charge of an ion when counting electron pairs.

How do you count electron pairs around a central atom in an ion? Count the central atom’s outer electrons and the electrons used in bonding, then add one electron for each negative charge or remove one for each positive charge before halving. Forgetting the charge of an ion is a common error.

How should you compare the melting points of two substances? Name the forces or bonds in both substances, say which are stronger, and link this to the energy needed to overcome them. Writing “weaker forces” without saying which forces, or describing only one substance, loses marks.

Official syllabus

OxfordAQA, International AS and A-level Chemistry (9620) specification, Version 5.3, for International AS and A-level exams May/June 2020 onwards, section 3.1.3 Bonding (3.1.3.1 to 3.1.3.7), published by OxfordAQA.