Which statement correctly contrasts diamond and graphite in terms of conductivity?

• A. Graphite conducts electricity, diamond does not
• B. Both conduct equally
• C. Diamond conducts more than graphite
• D. Neither conducts electricity

Answer: (A)

The key idea here is how the structure and bonding of carbon affect the presence of mobile charge carriers. Graphite conducts electricity because its carbon atoms form layers of hexagonal rings with sp2 bonds. In each carbon, three bonds use up hybrid orbitals, leaving one electron in a p orbital that becomes delocalized over the layer. These delocalized electrons can move and carry charge, so electrical conductivity occurs within the layers (and to some extent along them). Diamond, however, forms a giant three-dimensional network of sp3 bonds where each carbon is tetrahedrally bonded to four others, so all electrons are localized in strong covalent bonds and there are no free charge carriers to move. Hence diamond does not conduct electricity.

So the correct contrast is that graphite conducts electricity while diamond does not. The other statements don’t fit: if both conducted equally or if neither conducted, that would ignore the delocalized electrons in graphite; diamond conducting more contradicts its lack of free electrons.