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Question

Why is graphite a good conductor of electricity but not a diamond?


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Solution

Graphite

  • Graphite and diamond are allotropes of carbon.
  • Each carbon in graphite is sp2 hybridized.
  • A hexagonal array in graphite is formed when each carbon atom is connected to three other carbon atoms in the same plane.
  • Structure of graphite:
Allotropes of Carbon - Properties, Structure of Carbon Allotropes

Electricity conductivity of graphite

  • In graphite, only three carbon atoms are directly bonded to one another through covalent bonds.
  • As a result, bonding requires just three of a carbon atom's four valence electrons, while the fourth is relatively free and can move from one carbon atom to another.
  • With these free electrons, graphite is a good conductor of electricity.

Diamond

  • Each carbon atom in a diamond is connected to four other carbon atoms by strong covalent bonds, forming a huge covalent structure.
  • In diamond a regular tetrahedral network structure is formed by the carbon atoms.
  • Structure of diamond:
Diamond | Properties of Diamond - Applications and Origins | Chemistry

Electricity conductivity of Diamond

  • Diamonds have no free electrons that can move and conduct electricity because of their tetrahedral arrangement of covalently bound carbon atoms.
  • Since the existence of free electrons is required for electrical conduction.
  • As a result, diamond is a poor electrical conductor.

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