Why are the melting and boiling points of graphite, and diamond so high?
2 Answers
Because both graphite and diamond are non-molecular species, in which each constituent C atom is bound to other carbon atoms by strong chemical bonds.
Explanation:
Both diamond and graphite are network covalent materials. There are no discrete molecules, and vaporization would mean disrupting strong interatomic (covalent) bonds. I am not sure of the physical properties of buckminsterfullerene, 60 carbon atoms arranged in a football shape, but because this species is molecular, its melting/boiling points would be substantially lower than its non-molecular analogues. So, because we are physical scientists, this is your homework: find the melting points of the three carbon allotropes and rationalize them on the basis of their molecularity.
Both have a similar reason as their structure is very similar; the difference is graphite has layers "connected' by weak inter-molecular forces.
Explanation:
Diamond:
A lot of energy is needed to overcome the strong covalent bonds between the carbon atoms. Thus, it has high melting and boiling points.
Graphite:
Though little energy is needed to overcome the weak inter-molecular forces between the layers, a lot of energy is still needed to overcome the strong covalent bonds between the carbon atoms.
