More good news! The properties of diamond and graphite and silicon dioxide have’t changed!
On a more serious note though, you need to be able to explain the properties of giant covalent networks / lattices in terms of their structure and their bonding – they are not the same things. Structure is about the arrangement of atoms, bond angles, how many bonds are being made. Bonding is covalent vs intermolecular (as in graphite – we have strong covalent bonds within layers, weak London forces / instantaneous dipole – induce dipole bonds between layers).
Long answer exam questions commonly ask you to contrast and compare the properties of simple molecules with giant network structures. And with all long answers, make sure you aren’t contradicting yourself or rambling – the quickest ways to lose marks. A good technique is to use bullet points to structure your answer (bullet points written out in full).
Giant covalent networks
Carbon, boron and phosphorus are elements that form macromolecular networks – different structural forms of the same element, such as graphite, diamond and buckminsterfullerene in carbon, are known as allotropes.
Diamond:
Diamond has an exceptionally high melting point and is one of the hardest substances due to the very high amounts of energy needed to break all the strong covalent C-C bonds in the network.
Graphite:
It is relatively easy to slide the layers from each other which makes graphite a useful lubricant that won’t combust under high temperatures (unlike oil and grease). It also conducts electricity due to the delocalised ∏-system of electrons over each layer.
You can read more about hybrid molecular orbitals and bonding theory here 😊.
Buckminsterfullerene:
Buckminsterfullerene, C60, is one member of a large family of fullerenes. It is a soft solid like graphite with weak intermolecular bonds between molecules, and behaves as a semi-conductor.
The story of the discovery of buckminsterfullerene is brilliantly told in the BBC Horizon documentary ‘Molecules with Sunglasses’. Go watch it!
Silicon dioxide:
There are a huge number of different structures of silica, SiO2, the most common of which is quartz. Different structural forms of the same compound are known as polymorphs. As might be expected from its diamond-like structure, it is a very hard mineral.
Polymers also have covalent macromolecular structures.
Simple molecules
Simple molecules include water, carbon dioxide and nitrogen. They have strong covalent bonds between atoms within the molecules (intramolecular bonds) but weak intermolecular bonds between molecules. The chemistry of these substances, whether elements or compounds, is governed by the nature of the intermolecular bonds, but they typically have very low melting and boiling points, are soft, do not conduct electricity and dissolve in non-polar solvents.
Practice questions
- Carbon and silicon both form oxides with the formula XO2. Explain why silicon dioxide is a solid at room temperature but carbon dioxide is a gas.
- Describe 2 features of the structure and bonding in graphite that are similar to that in diamond and 2 features that are different.
- (a) Explain why the boiling point of bromine is different to that of magnesium in terms of the structure and bonding in each element.
(b) Suggest why magnesium is liquid over a much greater range of temperature compared to bromine.
| Element | Melting point / K | Boiling point / K |
| Magnesium | 923 | 1383 |
| Bromine | 266 | 332 |
Answers
- SiO2 is a giant covalent network / lattice / held together by strong covalent bonds. CO2 has a simple molecular structure with weak instantaneous dipole – induced dipoles / London forces between molecules. Less energy is needed to separate out CO2 molecules than to break the covalent bonds in the SiO2 network.
- Similarities: consist only of carbon atoms, giant network / lattice structure, covalent bonding within layers of graphite same as covalent bonding between atoms in diamonds structure. Differences: in diamond every carbon atom is bonded to 4 others but in graphite each carbon atom is bonded to 3 others, carbon atoms in graphite are arranged in layers but in diamond the atoms are arranged tetrahedrally, in diamond the bond angle is 109.5° but 120° in graphite, no delocalised electrons in diamond but there are between the layers in graphite*.
* always make a full comparison in longer answers – don’t leave it hanging or give the examiner an opportunity not to give you the full marks!
- (a) Bromine has a simple molecular structure, magnesium is a giant metallic lattice. There are weak London forces / instantaneous dipole – induced dipole forces between bromine molecules that take little energy to break, hence low boiling point. The giant metallic lattice in magnesium is the result of strong electrostatic attraction between Mg2+ cations and a sea of delocalised electrons which takes a large amount of energy to overcome, hence the high boiling point.
(b) Magnesium is liquid over a greater range because the forces of attraction in the giant metallic lattice are so much stronger.