AS Chemistry: Ionic, Covalent and Metallic Bonding — 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 available free from your board.

Related: Ionic, Covalent and Metallic Bonding revision notes

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Section A

1. Define ionic bonding and metallic bonding. [4]

2. Define electronegativity and state its trend across a period, with a reason. [3]

3. What is a dative covalent bond? Give one example. [2]

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Section B

4. Compare diamond and graphite.

(a) State the number of covalent bonds each carbon forms in each. [2]

(b) Explain why graphite conducts electricity but diamond does not. [3]

(c) Explain why graphite is a good lubricant. [2]

(d) Both have very high melting points. Explain why. [2]

5. Sodium chloride and magnesium oxide are both giant ionic lattices.

(a) Explain why MgO has a much higher melting point than NaCl. [3]

(b) Explain why solid NaCl does not conduct electricity but molten NaCl does. [2]

(c) Pauling electronegativity values are: Na 0.9, Al 1.6, Cl 3.2. Use them to predict the type of bonding in sodium chloride and in aluminium chloride, explaining your reasoning. [3]

6. Iodine, I₂, has a melting point of 114 °C, but visibly sublimes at temperatures below this because it has a noticeable vapour pressure as a solid.

(a) Name the forces broken when iodine sublimes. [1]

(b) A student writes: “The strong covalent bonds in iodine are broken.” Explain why this is incorrect. [2]

7. Describe, using electron configurations, how calcium fluoride (CaF₂) forms an ionic bond. [3]

8. Two elements have a large Pauling electronegativity difference between them. Explain why this favours ionic rather than covalent bonding. [2]

9. Explain why metals are malleable, in terms of the structure of metallic bonding. [3]

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Answers

1. Ionic: the electrostatic attraction [1] between oppositely charged ions [1]. Metallic: the electrostatic attraction [1] between positive metal ions and a sea of delocalised electrons [1].

2. The ability of an atom to attract the bonding electrons in a covalent bond [1]. It increases across a period [1] because nuclear charge increases while shielding stays roughly constant and atomic radius decreases [1]. Fluorine, at the top right of the Periodic Table, is the most electronegative element overall.

3. A covalent bond in which both electrons of the shared pair come from the same atom [1]. Example: NH₄⁺ (or H₃O⁺, or any metal–ligand bond) [1]. Once formed, a dative bond is identical in every way to an ordinary covalent bond.

4. (a) Diamond: four [1]. Graphite: three [1]. Both are examples of a giant covalent structure, distinct from the giant ionic and metallic lattices discussed elsewhere in this topic.

(b) In graphite each carbon uses only three of its four outer electrons in bonding [1], leaving one delocalised electron per atom [1] which is free to move along the layers, carrying charge [1]. In diamond all four are used in covalent bonds, so none are free.

(c) The layers are held together only by weak instantaneous dipole–induced dipole forces [1], so they can slide over one another easily [1].

(d) Both are giant covalent structures [1] in which many strong covalent bonds must be broken to melt them [1].

5. (a) Mg²⁺ and O²⁻ carry double the charge of Na⁺ and Cl⁻ [1] and the ions are smaller [1], so the electrostatic attraction between ions is much stronger and more energy is needed to overcome it [1].

(b) In the solid the ions are held in fixed positions in the lattice and cannot move [1]. When molten, the ions are free to move and carry charge [1].

(c) Differences: Na–Cl 2.3, Al–Cl 1.6 [1]. The large difference for NaCl means chlorine takes the electron almost entirely, so the bonding is ionic [1]. The smaller difference for AlCl₃ means the electrons are shared, but unequally, so the bonding is polar covalent [1]. (Why a metal chloride can be covalent in terms of polarisation is not assessed — syllabus 3.1.4.)

6. (a) Instantaneous dipole–induced dipole (id-id, London dispersion) forces, a type of van der Waals’ force [1].

(b) Sublimation overcomes the intermolecular forces between I₂ molecules [1], not the covalent bond within each molecule — the iodine vapour still consists of I₂ molecules [1].

7. Calcium (2,8,8,2) transfers one electron to each of two fluorine atoms [1] (2,7), forming Ca²⁺ (2,8,8) and two F⁻ (2,8) [1]. The resulting ions pack into a giant ionic lattice held together by electrostatic attraction extending in all directions [1].

8. A large electronegativity difference means one atom can pull the shared pair almost entirely to itself, effectively forming ions rather than a genuinely shared pair [1]. A small or zero difference instead means neither atom dominates, so the pair remains covalent [1].

9. The layers of positive ions can slide past one another without breaking any specific bond [1], because the bonding is not between any two fixed atoms but a delocalised electron sea extending through the whole structure [1], which simply redistributes around the ions as they move [1].

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Where marks are usually lost

• Omitting “electrostatic attraction” from a bonding definition.
• Saying graphite’s fourth electron is “unbonded” rather than delocalised.
• Saying covalent bonds break when a simple molecular substance melts or sublimes.
• Calculating an electronegativity difference but not saying what its size predicts about the bond type.
• Explaining MgO’s melting point by charge alone, without mentioning ionic radius.
• Describing metallic bonding as a bond between two specific atoms, rather than a force extending through the whole structure.
• Confusing “no sharp electronegativity cutoff” with there being no way to compare bond types — always reason from the relative size of the difference.