Why does graphite conduct electricity while diamond does not, even though both are allotropes of carbon?

Question

Accepted Answer

## Model Answer In **graphite**, each carbon atom is bonded to only **three** other carbon atoms, leaving one electron free. These free electrons can move through the layers, allowing graphite to conduct electricity. In **diamond**, each carbon atom is bonded to **four** other carbon atoms in a rigid 3D structure, leaving **no free electrons**. Hence, diamond cannot conduct electricity. *Source: Chapter 4, Allotropes of Carbon (More to Know)* --- ## Explanation - The key concept is **free/delocalized electrons**: graphite has them (one bond is a double bond, freeing an electron per atom); diamond has none (all four valence electrons are used in bonding). - Examiners expect you to clearly state the bonding difference (3 vs 4 bonds) AND directly link it to presence/absence of free electrons for conduction. - Don't just say "different structures" — explain *why* that structure affects conductivity.

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