Explain why graphite can conduct electricity whereas diamond cannot, despite both being allotropes of carbon.

IB DP Chemistry Standard Level (2023 syllabus) — S2.4 From models to materials · Explain · 4 marks · View as Markdown

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

Model answer (4 marks)

In graphite each C atom forms three σ bonds with neighbours, leaving one valence electron unpaired. These electrons are delocalised over the planar sheets, creating a mobile electron system that can carry charge.
In diamond each C atom forms four σ bonds in a tetrahedral network, using all four valence electrons. All electrons are localised in bonds, so there are no mobile charge carriers and diamond does not conduct.

Examiner tips

Mention the delocalised π‑electron system in graphite
Explain that diamond’s sp³ bonds use all valence electrons
Show the link between mobile electrons and conductivity
Use the exact terms "delocalised" and "localised"

Common mistakes

Confusing graphite with graphene or mentioning only σ bonds
Saying diamond conducts at high temperatures
Using vague phrases like "more electrons" instead of explaining localisation

Mark scheme (4 marks)

In graphite, each carbon atom forms three covalent bonds, leaving one electron per carbon atom not involved in sigma bonding / forming a delocalised electron system.
These delocalised electrons are free to move through the layers and carry charge, so graphite conducts electricity.
In diamond, each carbon forms four covalent bonds in a tetrahedral arrangement, using all four valence electrons.
Because all electrons in diamond are localised in bonds, there are no mobile charge carriers, so diamond cannot conduct electricity.

Key terms in this question

allotrope