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: Bonding revision notes

Tier note: questions and parts marked (0620 Extended only, 5070 required) — ionic bonding beyond Group I with Group VII (e.g. magnesium oxide), explaining the properties of ionic compounds and of simple molecular compounds in terms of intermolecular forces, metallic bonding and silicon(IV) oxide — are not required for a 0620 Core candidate. All other questions are answerable at Core.

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Questions

1. Define ionic and covalent bonding, and (0620 Extended only, 5070 required) metallic bonding. [6]

2. Describe, using dot-and-cross reasoning, what happens to the electrons when sodium reacts with chlorine. [3]

3. (0620 Extended only, 5070 required) Explain why magnesium oxide has a much higher melting point than sodium chloride. [3]

4. (0620 Extended only, 5070 required) Explain why ionic compounds conduct electricity when molten or dissolved but not when solid. [2]

5. Compare diamond and graphite.

(a) State how many covalent bonds each carbon forms in each. [2] (b) Explain why graphite conducts electricity. [3] (c) Explain why graphite is soft and used as a lubricant. [2] (d) Explain why both have very high melting points. [2]

6. (0620 Extended only, 5070 required) A student states: “Carbon dioxide has a low boiling point because its covalent bonds are weak.” Explain what is wrong with this. [3]

7. (0620 Extended only, 5070 required) Explain why metals conduct electricity and are malleable. [4]

8. (0620 Extended only, 5070 required) Silicon(IV) oxide, SiO₂, is a giant covalent structure, unlike the simple molecular carbon dioxide compared in question 6.

(a) State one structural similarity between SiO₂ and diamond that explains why the two substances have similar physical properties. [2]

(b) Predict, giving a reason, whether SiO₂ conducts electricity. [2]

9. (0620 Extended only, 5070 required) Explain why metals are ductile (can be drawn into wires) as well as malleable (can be hammered into sheets), in terms of the same structural feature. [2]

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Answers

1. Ionic — the electrostatic attraction [1] between oppositely charged ions [1]. Covalent — a shared pair of electrons [1] attracted to the nuclei of both atoms [1]. (0620 Extended only, 5070 required) Metallic — the electrostatic attraction [1] between positive metal ions and a sea of delocalised electrons [1].

2. Sodium loses one electron from its outer shell, forming Na⁺ [1]. Chlorine gains that electron, forming Cl⁻ [1]. Both ions now have a full outer shell, and they are held together by electrostatic attraction [1].

Dot-and-cross diagram (outer shells only): Na’s one outer electron is shown as a dot (•) and Cl’s own seven outer electrons as crosses (×). The dot transfers to chlorine and stays a dot, so [Cl]⁻ has eight outer electrons — seven crosses and one dot — and [Na]⁺ has no outer electrons left:

             ××                             ××
 Na•   +   ××Cl×      →      [Na]⁺      [ ××Clו ]⁻
             ××                             ××

Na: 1 outer electron (•), lost      Cl⁻: 7 crosses (its own) + 1 dot (from Na)
[Na]⁺: no outer electrons left      = 8 outer electrons, a full outer shell

3. Mg²⁺ and O²⁻ carry double the charge of Na⁺ and Cl⁻ [1], so there is a greater electrostatic force of attraction between the oppositely charged ions [1]; this stronger attraction requires much more energy to overcome, giving MgO a much higher melting point [1].

4. In the solid the ions are held in fixed positions in the lattice and cannot move [1]; when molten or dissolved the ions are free to move and carry charge [1].

5. (a) Diamond four [1]; graphite three [1]. (b) 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]. (c) The layers are held together only by weak forces between them [1], so they can slide over one another [1]. (d) Both are giant covalent structures [1] in which very many strong covalent bonds must be broken [1].

6. (0620 Extended only, 5070 required) Boiling overcomes the weak intermolecular forces between CO₂ molecules [1], not the covalent bonds within them [1]. The C=O covalent bonds are in fact very strong and remain intact — the gas still consists of CO₂ molecules [1].

7. (0620 Extended only, 5070 required) They contain delocalised electrons that are free to move through the structure and carry charge [1] [1]. The layers of positive ions can slide over one another without breaking the metallic bonding [1], because the delocalised electrons continue to hold the structure together [1].

8. (0620 Extended only, 5070 required) (a) Like diamond, every atom is joined to its neighbours by strong covalent bonds throughout a rigid three-dimensional lattice, with no weak points to overcome separately [1]; this is why both have very high melting points and great hardness [1]. (b) SiO₂ does not conduct electricity [1], because, unlike graphite, all of the outer electrons are localised, either in a bonding pair or a lone pair (each silicon forms four bonds, each oxygen forms two and keeps two lone pairs), leaving none delocalised and free to carry charge [1].

9. (0620 Extended only, 5070 required) Ductility, like malleability, comes from the fact that the layers of positive ions can slide past one another into a new position [1], while the delocalised “sea” of electrons continues to hold the structure together regardless of shape — so the metal can be drawn into a wire (ductile) or hammered into a sheet (malleable) without shattering [1].

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Where marks are usually lost

• Omitting “electrostatic attraction” from a bonding definition.
• Saying covalent bonds break when a simple molecular substance boils.
• Saying metals are malleable “because they are soft”.
• Assuming every giant covalent structure conducts electricity like graphite — diamond and SiO₂ do not, because all their outer electrons are localised in bonding pairs (and, in SiO₂’s case, lone pairs too), leaving none delocalised.
• Confusing SiO₂’s structural similarity to diamond (a fully-bonded 3D covalent lattice) with a compositional one — they share a structure type, not a formula or the same elements.
• Forgetting that “high melting point” alone doesn’t distinguish giant covalent from ionic — the reasoning must specifically reference covalent bonds being broken throughout the structure, not electrostatic attraction between ions.

Questions 8 and 9 draw on the giant covalent structures and metallic bonding sections of the Ionic, Covalent and Metallic Bonding study guide — question 8 extends the diamond/graphite comparison in question 5 to a third giant covalent substance, and question 9 completes question 7’s account of malleability with the closely related property of ductility, neither previously tested on this page.