GraphiteOxford AQA IGCSE Chemistry: Revision notes
Section 1
What is graphite?
Graphite is another allotrope of carbon. Each carbon atom bonds covalently to three others (not four, as in diamond), forming flat layers of hexagonal rings. There are no covalent bonds between the layers — only weak forces hold the layers together.
Section 2
Why is graphite soft and slippery?
Because there are no covalent bonds between the layers, only weak forces hold them together. This means the layers can easily slide over each other, which is why graphite feels soft and slippery and is used as a lubricant and in pencil 'lead'.
Students often say graphite has 'weak covalent bonds' between layers — this is wrong. There are NO covalent bonds between layers at all, only weak intermolecular forces.
Section 3
Why does graphite conduct electricity?
Each carbon atom in graphite only forms three covalent bonds, so one electron per atom is not used in bonding. This spare electron becomes delocalised and is free to move along the layers, allowing graphite to conduct electricity and heat.
This is unusual for a non-metal and makes graphite useful as an electrode.
When explaining conductivity, always link it to 'one delocalised electron per carbon atom' — a vague answer like 'it has free electrons' will lose marks.
Must Know
- Graphite is a giant covalent structure of carbon in which each atom bonds to three others
- Atoms are arranged in layers of hexagonal rings
- There are no covalent bonds between layers, only weak forces, so layers slide over each other
- This makes graphite soft and slippery
- One delocalised electron per carbon atom allows graphite to conduct electricity and heat
- Graphite has a high melting point because the covalent bonds within each layer are strong
That's the notes covered.
Carry on to the next subtopic.