Compare the electrical conductivity of graphite and diamond, and explain the difference in terms of their structures and bonding.

Cambridge International IGCSE Chemistry (0620) — 2.6 Giant covalent structures · Compare · 4 marks · View as Markdown

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

Model answer (4 marks)

Graphite conducts electricity, whereas diamond does not.

In graphite each carbon atom is bonded to three others by covalent bonds, leaving one valence electron that is not involved in bonding. These electrons are delocalised over the layers and can move freely, allowing charge to be carried.

In diamond every carbon atom forms four covalent bonds, using all four valence electrons. There are no delocalised electrons, so no charge carriers are available and diamond is an electrical insulator.

Examiner tips

  • State the conductivity difference first, then explain using bonding and delocalisation.
  • Use the terms "delocalised electrons" and "covalent bonds" as per the mark scheme.
  • Show the contrast between 3‑fold and 4‑fold coordination clearly.

Common mistakes

  • Confusing graphite with diamond as both being forms of carbon; not mentioning delocalised electrons.
  • Using vague terms like "free electrons" without linking to bonding structure.
  • Failing to specify that diamond has no free electrons because all are used in bonds.

Mark scheme (4 marks)

  1. Graphite conducts electricity but diamond does not
  2. In graphite, each carbon atom forms only three covalent bonds, leaving one outer (delocalised) electron per atom
  3. These delocalised electrons in graphite are free to move through the structure and carry charge
  4. In diamond, every outer electron is used in four covalent bonds so there are no free/delocalised electrons to carry charge

Key terms in this question

electrical conductivity

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