Explain why graphite can conduct electricity but diamond cannot, even though both substances are forms of carbon.

Edexcel GCSE Chemistry (1CH0) — 1.7 Types of substance · Explain · 4 marks · View as Markdown

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

Graphite and diamond are both giant covalent structures made entirely of carbon atoms, yet they have very different electrical properties.

Model answer (4 marks)

In graphite each carbon atom is bonded to three others, leaving one valence electron unshared. These electrons are delocalised over the layers and can move freely, carrying charge and allowing electrical conduction.
In diamond every carbon atom forms four covalent bonds, using all valence electrons. There are no delocalised or free electrons, so diamond does not conduct electricity.

Examiner tips

  • Use the word ‘delocalised’ to show understanding of electron mobility.
  • Explain the difference in bonding (3 vs 4 bonds) to justify the presence/absence of free electrons.
  • Keep the answer concise – each point should be a short sentence.

Common mistakes

  • Confusing graphite with diamond as having the same bonding; forgetting to mention the free electron in graphite.
  • Using vague terms like ‘free electrons’ without linking to delocalisation.
  • Over‑expanding the answer with unnecessary detail about crystal structure.

Mark scheme (4 marks)

  1. In graphite, each carbon atom forms bonds with only three other carbon atoms (leaving one electron not used in bonding)
  2. Graphite has delocalised electrons that are free to move through the structure
  3. These delocalised electrons carry charge through graphite, allowing it to conduct electricity
  4. In diamond, each carbon atom forms four covalent bonds so all electrons are used in bonding / there are no free or delocalised electrons in diamond

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