The diamond and graphite structures (above),
which we first saw in chapter 4, now can be understood in terms of
orbitals and bonding. Each carbon atom in diamond uses sp3 hybridization
to form four bonds to neighbors, and all bonding electrons are localized.
In graphite, three electrons per carbon atom are localized in sp2-derived
bonds within one sheet, and the fourth electron is delocalized over
the entire sheet as though in a "two dimensional metal."
Diamond and graphite illustrate this behavior for carbon; diamond
is the nonmetallic allotrope, and graphite has some of the properties
of a metal (see above). The lack of color in diamond tells us that
it has no closely spaced electronic energy levels to absorb visible
light; all of its electrons are tied down in covalent bonds.
In graphite the electrons are delocalized and free to wander within
the layers of carbon atoms. The black color and metallic sheen of
graphite result from the absorption and reemission of many wavelengths
of light by these mobile electrons.