Explain why graphite can conduct electricity but most other giant covalent structures, such as diamond, cannot.

OCR GCSE Chemistry A: Gateway Science (J248) — C2.2 Bonding · Explain · 4 marks · View as Markdown

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

Graphite is used as an electrode material in industrial electrolysis cells. Diamond, despite also being a giant covalent structure made entirely of carbon atoms, is used as an electrical insulator in cutting tools.

Model answer (4 marks)

Graphite has delocalised electrons – one free electron per carbon atom – that can move throughout the layers.
These delocalised electrons act as charge carriers, allowing electrical conduction.
In diamond each carbon atom forms four covalent bonds, so all electrons are involved in bonding.
Without delocalised electrons there are no charge carriers, so diamond cannot conduct electricity.

Examiner tips

Use the word ‘delocalised’ and explain that they are free to move; link this to charge carriers. Show the contrast: diamond’s four‑bonded carbons leave no free electrons. Keep the answer concise – 4 points, 1 sentence each.

Common mistakes

Confusing graphite with diamond as having the same structure. Failing to mention that graphite’s electrons are delocalised. Using vague terms like ‘free electrons’ without linking to conduction.

Mark scheme (4 marks)

Graphite has delocalised electrons (one per carbon atom) that are free to move
These delocalised electrons can carry charge through the structure, allowing electrical conduction
In diamond, each carbon atom forms four covalent bonds, so all electrons are involved in bonding / there are no delocalised electrons
Without delocalised electrons there are no charge carriers, so diamond cannot conduct electricity

Key terms in this question

giant covalent structure