carbon-di-oxid and silicon di oxide!!?

Question

though C and Si belong to the same group, CO2 do not polymerise but Si O2 polymerises why?
please do help me to find the answer.

Answer 1

Here is ur ans in detail................

While silicon dioxide is chemically described as SiO2, you never actually encounter any in quite that form. You always encounter huge conglomerations of SiO2 molecules, linked together in extensive chemical 3-dimensional matrixes. This kind of linkage, for some simple single molecule like SiO2, is known as "polymerized". Each individual molecule is called a "monomer"; ordinary quartz (pure SiO2, remember) is a polymer. All substances known as "silicates" are also polymers, thanks to seemingly endless linkages of silicon-to-oxygen-to-silicon-to-oxygen-to-silicon-...

Okay, it is widely known that carbon and silicon are "related" atoms with similar chemical properties. Almost all the more widely-known polymers (polyethylene, polyester, polyvinyl chloride or PVC, polyurethane, etc.) are in fact based on chains of carbon atoms. Silicones, in fact, are silicon polymers that were designed BECAUSE of the fact that carbon molecules form polymeric chains so easily -- and silicon is chemically similar to carbon.

So...if polymerized silicon dioxide is so common, why don't we ever hear of polymerized carbon dioxide? When I first thought of this idea, a bunch of years ago, I could hardly believe that no one else had thought of it, but I couldn't find any references (which of course doesn't mean that there aren't any). I suppose, even today, that the stuff, if any was ever made, would prove to be unstable. Carbon dioxide may simply "prefer", energetically speaking, to exist in the form of individual molecules. (This sort of thing is well-known to be true for nitrogen gas molecules, N2. Most chemical explosives depend partly on the fact that if given half a chance, nitrogen will break loose from certain molecules, such as potassium nitrate, in order to recreate N2 molecules. That initial break-up leads to the release of other molecules, such as oxygen, which can fan the flames of the main explosive event.)

On the other hand, perhaps polymerized CO2 is stable, but requires such special conditions to create that you never find any in Nature. If this is true, what might its physical properties be? Well, since we are talking about a perfect analogue of quartz, it should be able to exist in a crystalline form. This is a remarkable form, because it is the same form that pure crystalline carbon takes -- otherwise known as diamond. And we all know that diamond is a remarkable material.

Pure crystalline silicon will also possess the same structure as diamond, but it is nowhere near as hard a substance as diamond. Why? Two reasons: One, the silicon-silicon bond is much weaker than the carbon-carbon bond; Two, carbon is a significantly smaller-than-average atom, and there are vastly more carbon-carbon bonds per unit of volume in a diamond, than than bonds-per-volume in any other substance.

Now consider quartz: It is harder than crystalline silicon, but not as hard as a diamond. This is because the silicon-oxygen bond is quite a strong chemical bond (stronger, if I recall right, than even the carbon-carbon bond). However, there are still a lot fewer chemical bonds per unit of volume in quartz than in diamond, so diamond remains the champion in hardness.

There is another substance, rather harder than quartz, known as "silicon carbide", which also possesses the diamond structure, and is among the hardest of all substances -- except for diamond. Due to the size of the silicon atom, there are again fewer carbon-silicon bonds per unit of volume than carbon-carbon bonds in diamond. (They are also not quite as strong as the carbon-carbon bond, if I recall right.)

With the background in place, we can now consider the suggested polymerized carbon dioxide. The carbon-oxygen bond is a VERY strong bond, significantly stronger than the carbon-carbon bond in diamond. However, oxygen atoms are larger than carbon atoms, and so there will again be fewer bonds per unit of volume than in diamond. Will the bond-strength make up the difference? Will poly-CO2 be as hard as, or even harder than diamond? THAT might explain why we never see any of these crystals in Nature! We all know what Nature has to do to make a diamond, and they are pretty rare....

Not to mention that ordinary CO2 is a gas, while ordinary pure carbon is a solid (graphite, for example). HOW in Nature will it collect, pressurize, and polymerize carbon dioxide???

However, we humans collect and pressurize pure CO2 everywhere we make soda-pop. One might think that if we added a bunch of heat, and significantly more pressure (and maybe some electrical discharges or blasts of laser light at the right frequency), we might find ourselves possessed of poly-CO2.

Will it be stable? I dunno. A friend of mine at work told me that he once saw something on the Web about some other dude who had made some, and it decomposed and destroyed his laboratory, but I know that this friend often exaggerates for entertainment value. He might have made the story up. Certainly I haven't been able to find any reference to the alleged incident.

IF it is stable, though, then simply because it will probably be relatively easier to make than diamond, we could find a lot of uses for it. For example, it is known that ordinary glass gets STRONGER under pressure (up to the limit where it collapses, of course); some submersibles are simply two glass hemispheres and a rubber seal. Poly-CO2 should be stronger yet. This is, in fact, the stuff of which vacuum balloons might be made....

Answer 2

The reason lies in the fact that Si uses its d orbital in the hybridised orbitals which prevents assuming an sp2 state necessary for multiple bonds.