Re: Why is Si, unlike carbon, not able to exist in a graphite-like structure?

Allow me to quote from my favorite general chemistry text:

"Silicon crystallizes with a diamondlike structure in which each silicon
atom is covalently bonded to four neighboring silicon atoms at the corners
of a regular tetrahedron.  Thus, a single crystal of silicon is a three-
dimensional giant molecule."

(from Nebergall, Holtclaw, and Robinson, "College Chemistry", 6th Ed., D.C.
Heath and Co., Lexington, MA, 1980.)

So silicon forms a three-dimensional structure similar to diamond.  Why,
then, does it not form a structure like graphite, in which the carbon
atoms are arranged in sheets?

The reason is that in graphite, the carbon atoms sit in flat six-membered
rings with delocalized pi-electrons above and below.  A single "sheet" of
graphite looks like benzene rings all bonded together.  Each carbon atom
is SP-2 hybridized, as in benzene.  

So why not silicon?

Silicon has a very difficult time forming double bonds.  Its outer shell
of electrons are in the 3S and 3P orbitals, and it takes a good deal more
energy to hybridize these into SP hybrids.  Thus, trying to force silicon
atoms into a flat ring with delocalize pi-orbitals simply takes too much
energy.  Silicon is much happier forming four sigma bonds to its neighbors.