|MadSci Network: Chemistry|
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.
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