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 whilst diamond cannot, and explain why both substances have high melting points.
Written & reviewed by James Millett — Biology (Imperial College London), PGCE Science (University of Cambridge).
Graphite is used as an electrode material in electrolysis because it conducts electricity. Diamond, despite also being made of carbon, is used as a cutting tool rather than an electrical conductor.
Model answer (5 marks)
In graphite each carbon atom is bonded to three others by strong covalent bonds, leaving one valence electron that is not involved in a bond. These electrons are delocalised over the layers and can move freely, allowing graphite to conduct electricity. In diamond every carbon atom forms four covalent bonds, so all valence electrons are tied up in bonds and there are no free electrons; therefore diamond does not conduct.
Both graphite and diamond are giant covalent networks. The bonds throughout the structure are very strong covalent bonds, so a large amount of energy is required to break them. This gives both materials very high melting points.
Both graphite and diamond are giant covalent networks. The bonds throughout the structure are very strong covalent bonds, so a large amount of energy is required to break them. This gives both materials very high melting points.
Examiner tips
- Use the word ‘delocalised’ for graphite’s free electrons
- Mention the 3‑bonding in graphite vs 4‑bonding in diamond
- State that strong covalent bonds give high melting points
- Keep the answer concise and use correct terminology
Common mistakes
- Confusing graphite’s layers with diamond’s structure
- Saying graphite conducts because it is metallic instead of due to delocalised electrons
- Forgetting to mention the 3‑ vs 4‑bonding difference
Mark scheme (5 marks)
- In graphite, each carbon atom forms three covalent bonds (with three other carbon atoms)
- Graphite has one delocalised electron per carbon atom which is free to move and carry charge
- In diamond, each carbon atom forms four covalent bonds so there are no delocalised electrons
- Both structures contain many strong covalent bonds throughout the structure
- A large amount of energy is needed to break these covalent bonds, resulting in high melting points
Key terms in this question
giant covalent structure · graphite · diamond
Related
- All OCR A-Level Chemistry B: Salters (H433) revision notes →
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- Decode the mark scheme abbreviations →
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