Explain why graphite can conduct electricity but diamond cannot, despite both being giant covalent structures made entirely of carbon atoms.

OCR GCSE Chemistry A: Gateway Science (J248) — C2.3 Properties of materials · 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 allotropes of carbon used in very different industrial applications. Graphite is used in electrodes, whilst diamond is used as an electrical insulator in cutting tools.

Model answer (4 marks)

In graphite each carbon atom is bonded to three others, leaving one valence electron that is not involved in a covalent bond. This electron is delocalised over the layers and can move freely, carrying charge and allowing graphite to conduct electricity.

In diamond each carbon atom is bonded to four others, so all valence electrons are used in covalent bonds. There are no free or delocalised electrons, so diamond cannot conduct electricity.

Examiner tips

Use the terms ‘delocalised electron’ and ‘free electron’ to show understanding of conductivity.
Explain the difference in bonding (three vs. four bonds) to justify the presence or absence of free electrons.
Mention that graphite’s layered structure allows electron mobility, while diamond’s rigid tetrahedral network does not.

Common mistakes

Confusing graphite with diamond by stating both have the same bonding; they differ in the number of bonds per carbon.
Failing to mention that the free electron in graphite is delocalised over the layers, not just a single electron per atom.
Using vague language such as ‘more electrons’ instead of specifying delocalised vs. bonded electrons.

Mark scheme (4 marks)

In graphite, each carbon atom forms three covalent bonds (with three other carbon atoms), leaving one electron per carbon atom not used in bonding.
This leaves a delocalised electron per carbon atom that is free to move and carry charge, allowing graphite to conduct electricity.
In diamond, each carbon atom forms four covalent bonds with four other carbon atoms.
All electrons in diamond are involved in covalent bonds, so there are no free/delocalised electrons to carry charge, meaning diamond cannot conduct electricity.

Key terms in this question

giant covalent structure