Graphite is a much better conductor of heat and electricity than diamo

Graphite is a much better conductor of heat and electricity than diamond. This is due to the fact that each carbon atom in graphite:

[amp_mcq option1=”undergoes sp² hybridization and forms three sigma bonds with three neighbouring carbon atoms” option2=”undergoes sp³ hybridization” option3=”is tetrahedrally bonded” option4=”is free from van der Waals force” correct=”option1″]

This question was previously asked in

UPSC NDA-2 – 2015

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The correct answer is A) undergoes sp² hybridization and forms three sigma bonds with three neighbouring carbon atoms.

Graphite’s structure consists of layers of carbon atoms arranged in hexagonal lattices. Each carbon atom in graphite is sp² hybridized and forms three strong sigma bonds with three adjacent carbon atoms within the same layer. The remaining unhybridized p-orbital on each carbon atom overlaps sideways with p-orbitals of neighbouring atoms, forming a delocalized pi electron system across the layer. These delocalized electrons are free to move within the layers, making graphite a good conductor of heat and electricity.

In contrast, diamond has a tetrahedral structure where each carbon atom is sp³ hybridized and forms four strong sigma bonds with four neighbouring carbon atoms. All valence electrons are localized in these covalent bonds, so there are no free electrons to conduct electricity, making diamond an excellent electrical insulator. The strong covalent network also makes diamond an excellent thermal conductor, but the question specifically asks about electrical conductivity compared to diamond, and graphite’s delocalized electrons are key. Graphite’s thermal conductivity is also generally higher than diamond in some directions due to strong in-plane bonds and vibration modes. However, the primary reason for its *better* electrical conductivity than diamond is the presence of delocalized electrons, which stems from its sp² hybridization and bonding structure described in option A.

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