Edexcel IGCSE Chemistry · Spec 1.49-1.51
Giant Covalent Structures
Giant covalent structures such as diamond, graphite and silicon dioxide.
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Giant Covalent Structures
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Explained
Diamond and graphite, same element, opposite properties
A giant covalent structure is a continuous network of atoms joined by strong covalent bonds, with no separate molecules anywhere in it. Diamond, graphite and silicon dioxide are the three you need.
Diamond and graphite are both pure carbon, yet one is the hardest natural substance and the other is soft enough to write with. Every difference between them comes from how many neighbours each carbon atom is bonded to.
Diamond
Each carbon atom forms four covalent bonds, giving a rigid three dimensional network.
It is extremely hard, because there is no weak direction in the structure and any force has to break covalent bonds. It has a very high melting point, because a great deal of energy is needed to break the many strong covalent bonds. And it does not conduct electricity, because all four outer electrons are used in bonding, so none are free to move.
Its hardness is why it is used in cutting tools and drill tips.
Graphite
Each carbon atom forms only three covalent bonds, so the atoms form flat layers of hexagons.
The layers are held to each other by weak forces, so they can slide over one another. That is why graphite is soft and slippery, why it is used as a lubricant, and why a pencil leaves a mark.
The fourth outer electron on each atom is not used in bonding. These electrons are delocalised and free to move between the layers, so graphite conducts electricity. It is the only non metal that does, which is why it is used for electrodes in electrolysis.
Its melting point is still very high, because melting it means breaking the covalent bonds within the layers, not just separating the layers.
What the mark scheme accepts and rejects
An Edexcel International GCSE Chemistry mark scheme for a six mark graphite question lists seven possible points but states that two of them must be included for full marks: that the layers can slide over each other, and that the electrons are free to move. Structure alone is not enough; the consequence has to be stated.
Two notes show how the wording is policed. If a response refers to weak intermolecular forces, or to layers of molecules, the mark for weak forces between the layers is not awarded. And the terms free electron, spare electron and unbonded are ignored, so delocalised is the word that earns credit.
The neighbouring question is stricter still. Its note says that any mention of intermolecular forces, forces between molecules or ions, ionic bonding or metallic bonding scores zero out of three.
The reason is that graphite contains no molecules at all. Calling the forces between its layers intermolecular says there are molecules there, which contradicts the whole idea of a giant structure. The safe phrase is weak forces between the layers.
Silicon dioxide and the fullerenes
Silicon dioxide, found in sand and quartz, has a structure like diamond, with each silicon atom bonded to four oxygen atoms and each oxygen to two silicons. It is therefore hard, has a very high melting point, and does not conduct.
C sixty is different. It is a fullerene, a hollow ball of sixty carbon atoms, and it is a molecule rather than a giant structure. Because separate molecules are held to each other only by weak intermolecular forces, C sixty is soft and melts at a low temperature, even though the bonds inside each ball are strong covalent ones.
This is the one carbon structure where the word intermolecular is correct, and knowing why it is correct here and wrong for graphite is the point of the comparison.
Spec 1.49-1.51
What you need to know
- Describe the structures of diamond and graphite
- Explain their properties from those structures
- Recognise C sixty, a fullerene
Active recall
Quick check
Answer each question before opening the answer.
Compare the structure and properties of diamond and graphite.
Diamond — each carbon bonded to 4 others, very hard, does not conduct. Graphite — each carbon bonded to 3 in layers that slide (soft, lubricant), and has delocalised electrons so it conducts.
Why do giant covalent structures have very high melting points?
They have many strong covalent bonds that need a lot of energy to break.
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