Graphite and diamond are both giant covalent structures made entirely of carbon atoms, yet they have very different properties. Describe the structure of graphite and explain why graphite can conduct electricity whilst diamond cannot.

Pearson Edexcel International GCSE Chemistry (4CH1) — 1.7 Covalent substances · Describe and Explain · 4 marks · View as Markdown

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

Model answer (4 marks)

Graphite consists of layers of hexagonal rings; each carbon atom forms three covalent bonds with three other carbons, leaving one electron free.
There are no covalent bonds between the layers, so the free electrons can move.
Thus graphite conducts electricity.
In diamond every carbon atom uses all four outer electrons in covalent bonds, so there are no free electrons and it does not conduct.

Examiner tips

  • Use the word ‘layers’ and ‘hexagonal rings’ to show structure. Mention the free electron per carbon. Explain the lack of inter‑layer bonds. State that diamond’s electrons are all bonded.
  • common_mistakes
  • :
  • Confusing graphite with diamond structure. Forgetting that diamond has no free electrons. Using vague terms like ‘strong bonds’ instead of specifying covalent bonds.

Mark scheme (4 marks)

  1. Each carbon atom in graphite forms three covalent bonds with three other carbon atoms, forming layers of hexagonal rings
  2. There are no covalent bonds between the layers
  3. Each carbon atom in graphite has one delocalised electron per carbon atom that is free to move
  4. In diamond, all four outer electrons are used in 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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