Covalent Bonding
7 marks · HardThis question is about how the physical properties of substances are linked to their structures and electrical conductivity.
The table shows data for some simple molecular and giant covalent substances:
| Substance | Relative molecular mass / Mr | Melting point / °C | Boiling point / °C |
|---|---|---|---|
| Chlorine (Cl₂) | 71 | -101 | -34 |
| Bromine (Br₂) | 160 | -7 | 59 |
| Iodine (I₂) | 254 | 114 | 184 |
| Diamond | – | 3550 | 4830 |
Using data from the table, describe and explain how the melting and boiling points of the halogens change as the relative molecular mass increases.
Explain why the melting point of diamond is much higher than those of the halogens listed, using ideas about structure and bonding.
State whether graphite conducts electricity. Suggest, based on your knowledge of its structure, why this is different from diamond and the halogens.
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.
Show answer
(b) For the halogens shown (Cl₂, Br₂, I₂), both melting point and boiling point increase as the relative molecular mass increases. This is because there are stronger intermolecular forces (van der Waals forces) between molecules with higher mass, so more energy is needed to separate them.
(c) Diamond has a much higher melting point because it has a giant covalent structure, where very strong covalent bonds extend throughout the crystal. Breaking these bonds requires much more energy than overcoming the weak intermolecular forces between halogen molecules.
(d) Graphite does conduct electricity. This is due to its structure: each carbon atom is bonded to only three others, leaving one electron per atom delocalised, which is free to move and carry charge through the layers. In diamond, all electrons are locked in covalent bonds, so it cannot conduct. Halogens do not conduct electricity because their molecules have no free electrons or ions.
Common mistakes: Only mentioning melting point, not both melting and boiling points for halogens; Confusing 'giant covalent' and 'simple molecular' bonding in explanations of diamond; Vague statements on 'free electrons' without relating to bonding context
Show mark scheme and worked solution
Marking points
M1 (b) · 1 mark · Describes that the melting and boiling points both increase as Mr increases for halogens.
A1 (b) · 1 mark · Explains increase in terms of stronger intermolecular forces needing more energy.
M2 (c) · 1 mark · States that diamond has a giant covalent structure with many strong covalent bonds extending throughout the lattice.
A2 (c) · 1 mark · Explains that breaking covalent bonds in diamond requires much more energy than breaking intermolecular forces in halogens.
M3 (d) · 1 mark · States graphite conducts electricity, unlike the others.
A3 (d) · 1 mark · Explains that graphite conducts due to delocalised electrons.
A4 (d) · 1 mark · Explains why diamond and halogens do not conduct (no free electrons or mobile charge carriers).
Worked solution
a) Table provided. No answer needed.
b) As we go from Cl₂ to Br₂ to I₂, both melting and boiling points increase. This is because the relative molecular mass increases, making the intermolecular forces (van der Waals forces) stronger, so more energy is needed to overcome them.
c) Diamond's melting point is much higher than those of the halogens because diamond is a giant covalent structure. Its atoms are joined by strong covalent bonds in a 3D network that requires much more energy to break than the weak forces between simple molecules in halogens.
d) Graphite conducts electricity because each C atom has one delocalised electron per atom due to bonding to three other C atoms, and these electrons are free to move within the layers. In contrast, diamond has all electrons tightly held in covalent bonds, and halogen molecules have no free electrons or ions, so neither diamond nor the halogens can conduct electricity.