Covalent Bonding
8 marks · HardThis question is about the differences in properties between substances with simple molecular structures and those with giant covalent structures.
Explain why substances with simple molecular structures, such as chlorine (Cl₂), have low melting and boiling points.
Diamond and graphite are both forms of carbon, but diamond is much harder than graphite. Explain, in terms of structure and bonding, why diamond is harder than graphite.
State and explain whether Cl₂ can conduct electricity. Then compare this with the ability of diamond to conduct electricity, explaining your reasoning.
Work this out on paper, then compare with the answer and worked solution below. With a free account you can submit written answers for marking.
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(a) Simple molecular substances like Cl₂ have low melting and boiling points because weak intermolecular forces (van der Waals forces) hold the molecules together. Little energy is required to overcome these weak forces when the substance is melted or boiled. The strong covalent bonds within molecules are not broken; only the weak forces between molecules are overcome.
(b) Diamond is harder than graphite because, in diamond, each carbon atom forms four strong covalent bonds to four other carbon atoms, creating a three-dimensional giant lattice. The strong covalent bonds extend throughout the structure, making it very rigid and hard to break. In graphite, each carbon atom is bonded to only three others in flat layers, with weak forces between the layers, so the layers can slide over each other, making graphite much softer.
(c) Cl₂ cannot conduct electricity because it consists of neutral molecules with no charged particles or free electrons to carry charge. Diamond also does not conduct electricity because all its electrons are held tightly in covalent bonds and none are free to move. In both structures, there are no mobile charge carriers.
Common mistakes: Confusing breaking covalent bonds with breaking intermolecular forces in simple molecules; Stating 'Cl₂ conducts electricity as a gas or liquid' (it does not); Confusing diamond's structure with graphite's
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Marking points
M1 (a) · 1 mark · Identifies weak intermolecular forces between molecules as the reason for low melting/boiling points in simple molecular structures.
A1 (a) · 1 mark · Explains that only the weak intermolecular forces are broken, not covalent bonds.
M2 (b) · 1 mark · Describes the extended network of strong covalent bonds in diamond.
A2 (b) · 1 mark · Explains that this makes diamond very rigid/hard.
M3 (b) · 1 mark · Compares with graphite, noting only three bonds per C atom and weak forces between layers allow movement/layer sliding (so softer).
M4 (c) · 1 mark · States that Cl₂ does not conduct electricity because there are no charged particles or mobile electrons.
M5 (c) · 1 mark · States that diamond also does not conduct electricity and explains that all electrons are held tightly in covalent bonds.
A3 (c) · 1 mark · Compares both, stating neither has mobile charge carriers.
Worked solution
a) Simple molecular substances, like Cl₂, have low melting and boiling points because the only forces holding the molecules together are weak intermolecular (van der Waals) forces. These are easily overcome by a small amount of energy, so the substance changes state at a low temperature. The strong covalent bonds within the Cl₂ molecules are not broken, only the weak forces between them.
b) In diamond, each carbon atom forms four strong covalent bonds to four other carbons, producing a 3D tetrahedral network that extends throughout the crystal. This structure is very stable and rigid, making diamond extremely hard. In contrast, graphite's carbon atoms are joined in layers by three strong covalent bonds, but the forces between these layers are weak so the layers slide, making graphite soft.
c) Cl₂ cannot conduct electricity because its molecules do not have an overall charge and all electrons are in fixed positions in the molecule. There are no ions or delocalised electrons. Similarly, diamond does not conduct electricity because all electrons are tightly bound in the covalent-bonded network, so none are free to move and carry charge. Therefore, neither substance conducts electricity.