Graphite is a giant covalent structure. Explain why graphite can conduct electricity but diamond cannot.

AQA GCSE Chemistry (8462) — 4.2.3 Structure and bonding of carbon · Explain · 4 marks · View as Markdown

Written & reviewed by James Millett — Biology (Imperial College London), PGCE Science (University of Cambridge).

Model answer (4 marks)

In graphite each carbon atom forms three covalent bonds with other carbon atoms, leaving one valence electron that is delocalised over the layers. These free electrons can move under an electric field, carrying current. In diamond each carbon atom forms four covalent bonds, so all valence electrons are involved in bonding and none are free to move, preventing electrical conduction.

Examiner tips

Use the term "delocalised electrons" to show understanding of conductivity.
Mention the difference in bonding (3 vs 4) to justify the presence/absence of free electrons.
Keep the answer concise and directly linked to the marks scheme points.
Use UK spelling (e.g., "delocalised").

Common mistakes

Confusing graphite with diamond as having the same bonding; not recognising the free electron in graphite.
Using vague phrases like "more electrons" without specifying delocalisation.
Failing to mention that diamond has no free electrons due to four‑coordinate bonding.

Mark scheme (4 marks)

In graphite, each carbon atom forms three covalent bonds (with three other carbon atoms)
This leaves one electron per carbon atom that is delocalised / free to move through the structure
These delocalised electrons can carry electrical charge / current through graphite
In diamond, each carbon atom forms four covalent bonds so there are no delocalised / free electrons to carry charge

Key terms in this question

giant covalent structure