Graphite and diamond are both giant covalent structures made entirely of carbon atoms, yet they have very different properties. 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 neighbouring carbons, leaving one valence electron that is not involved in bonding. This electron is delocalised over the layers and can move freely, allowing graphite to conduct electricity.

In diamond each carbon atom forms four covalent bonds, using all four valence electrons. There are no delocalised electrons, so diamond cannot conduct electricity.

Examiner tips

  • Use the word ‘delocalised’ or ‘free’ electrons when describing graphite’s conductivity.
  • Mention the difference in bonding (3 vs 4 bonds) to explain the absence of free electrons in diamond.
  • Keep the answer concise – one sentence per point is enough for full marks.

Common mistakes

  • Failing to state that graphite has delocalised electrons; writing only that it has ‘free electrons’ without linking to the 3‑bond structure.
  • Confusing the number of bonds in diamond with graphite; e.g., saying diamond has 3 bonds.

Mark scheme (4 marks)

  1. In graphite, each carbon atom forms three covalent bonds (with three other carbon atoms)
  2. This leaves one electron per carbon atom that is delocalised / not used in bonding
  3. These delocalised electrons are free to move through the structure and carry electrical charge
  4. In diamond, each carbon forms four covalent bonds so there are no delocalised / free electrons to carry charge

Key terms in this question

giant covalent structure · graphite · diamond

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