You have to break the covalent bonding throughout the whole structure. In order to melt graphite, it isn't enough to loosen one sheet from another. Graphite has a high melting point, similar to that of diamond. As the delocalized electrons move around in the sheet, very large temporary dipoles can be set up which will induce opposite dipoles in the sheets above and below - and so on throughout the whole graphite crystal. So what holds the sheets together? In graphite you have the ultimate example of van der Waals dispersion forces. The atoms within a sheet are held together by strong covalent bonds - stronger, in fact, than in diamond because of the additional bonding caused by the delocalized electrons. There is, however, no direct contact between the delocalized electrons in one sheet and those in the neighboring sheets. The important thing is that the delocalized electrons are free to move anywhere within the sheet - each electron is no longer fixed to a particular carbon atom. They are no longer associated directly with any particular atom or pair of atoms, but are free to wander throughout the whole sheet. These "spare" electrons in each carbon atom become delocalized over the whole of the sheet of atoms in one layer. That leaves a fourth electron in the bonding level. The Chemistry and Physical Properties of Diamond Graphite and the FullerenesĮach carbon atom uses three of its electrons to form simple bonds to its three close neighbors.
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