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Edited on Wed Nov-02-11 02:10 AM by eppur_se_muova
Silicon forms long chain molecules which are more reactive under Earthly conditions (air and water) because of the very strong bonds which form between Si and O, and the relative weakness of the Si-H (and even Si-Si) bonds. The ability of Si and heavier elements to form "hypervalent" intermediates -- which have no second-row analogues -- also leads to *faster* bond-cleavage reactions relative to carbon. It's not that hard to make polysilanes in the lab IF you work under an argon line. If they have too many Si-H bonds they tend to be pyrophoric (ignite spontaneously on exposure to air), but in the absence of air and water (and at lower temperatures) they're fine. That's why most of the Si on Earth is in the form of SiO2 or silicates. Even elemental Si is not found naturally on Earth. {ETA: SiC is found as a very rare mineral, Moissanite, in meteorites and some minerals of volcanic origin. That's the closest thing to a reduced form of Si found naturallyon Earth.} (There are, by the way, polygermanes, -stannanes, and even -plumbanes, although the stability of the M-M bond decreases drastically in that order.)
Silicon does not form strong double (or triple bonds), so it cannot form anywhere near the variety of compounds that carbon does. Many molecular architectures adopted by carbon compounds (often with the help of other second-row elements, N & O) have no third-row, Si, analogs, except for a few, highly reactive, lab-created examples. (Most notably lacking are the many "aromatic" molecules formed by C,N, and O rings. All of the nucleic acid (DNA & RNA) bases fall in this category.) Ditto P versus N. So even in a strongly reducing, oxygen-free envirionment, silicon would be a poorer basis for complex structures than carbon.
If you had to design the chemistry of life from scratch, you would start with a second-row element (relatively strong single bonds, also double and triple bonds, no d orbitals so no hypervalency), and one that forms more than one bond at a time (eliminating Li and F). You would want to avoid empty p orbitals (this eliminates Be and B) because these provide a mechanism for attack by other compounds and cleavage of bonds, just as empty d orbitals do; you want your molecules to last. You would want to avoid too many lone pairs because of the 'beta lone pair effect' which leads to weaker bonds, and the 'acetal effect' which also leads to facile fragmentation -- so no long chains or networks containing too many N or O atoms. C only makes up a few hundred ppm of the Earth's crust, but it's the only element that fills the bill. No surprise that we haven't seen any signs of non-carbon based life yet, even under "unearthly" conditions.
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