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

AQA A-Level Chemistry (7405) — 3.1.3 Bonding · Explain · 5 marks · View as Markdown

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

Model answer (5 marks)

Graphite has delocalised electrons that are free to move.
In graphite each carbon atom forms three covalent bonds, so only three of its four outer electrons are used in bonding.
These delocalised electrons can carry charge and move through the structure when a voltage is applied.
In diamond every carbon atom forms four covalent bonds, so all four outer electrons are used in bonding.
Thus diamond has no delocalised or free electrons, so it cannot conduct electricity.

Examiner tips

  • Use the term "delocalised electrons" for graphite and "no free electrons" for diamond. Show the bonding difference: 3 bonds in graphite, 4 in diamond. Explain that free electrons allow charge transport. Keep the answer concise and to the point.
  • common_mistakes
  • :
  • Confusing graphite with diamond in terms of bonding. Forgetting to mention that diamond uses all four electrons. Using vague phrases like "electrons are free" without linking to conductivity.

Mark scheme (5 marks)

  1. Graphite has delocalised electrons (that are free to move)
  2. In graphite each carbon atom forms three covalent bonds / only three of its four outer electrons are used in bonding
  3. These delocalised electrons can carry charge / move through the structure when a voltage is applied
  4. In diamond every carbon atom forms four covalent bonds / all four outer electrons are used in bonding
  5. So diamond has no delocalised / free electrons, therefore it cannot conduct electricity

Related

More Bonding questions

▶ Try answering this question with AI marking (free) →