Answer:
In graphite, carbon is \[s{{p}^{2}}\]-hybridized
and each carbon is linked to three other carbon atoms by forming hexagonal
rings. Each carbon is now left with one unhybridized p-orbital which undergoes
sideways overlap to form three \[p\pi -p\pi \] double bonds. Thus, graphite has
two-dimensional sheet like (layered) structure consisting of a number of
benzene rings fused together. Silicon, on the other hand, does not form an
analogue of carbon because of the following reason:
Due to bigger size and smaller electro negativity
of Si than C, it does not undergo \[s{{p}^{2}}\]-hybridization and hence it
does not form \[p\pi -p\pi \] double bonds needed for graphite like structure.
Instead, it prefers to undergo only \[s{{p}^{3}}\]-hybridization and hence
silicon has diamond like three-dimensional network structure.
You need to login to perform this action.
You will be redirected in
3 sec
