Graphite and diamond are two allotropes of carbon that exhibit vastly different physical properties. Which statement correctly explains these differences in terms of their structure and bonding?
In graphite, each carbon atom is covalently bonded to 333 others in planar layers, creating delocalised electrons that permit electrical conductivity, while weak intermolecular forces between layers allow them to slide. In diamond, each carbon atom is covalently bonded to 444 others in a rigid, three-dimensional giant covalent lattice, rendering it extremely hard and non-conductive.
Graphite has a much lower melting point than diamond because only the weak intermolecular forces holding the layers together must be overcome to melt it, whereas strong covalent bonds must be broken to melt diamond.
Graphite conducts electricity because carbon atoms in the hexagonal sheets form ionic bonds that release mobile ions, whereas diamond is an electrical insulator because all of its carbon–carbon bonds are highly polar covalent bonds.
Diamond is harder than graphite because the individual carbon–carbon covalent bonds in its tetrahedral lattice are significantly stronger than the covalent bonds within graphite's hexagonal sheets.