Oops. You are right, ZenBeam. I was just trying to make the point that the orientation of the forth orbital is determined by its p character, which by necessity has it pointing away from the other carbon atom. But it is incorrect to call this a p orbital. It is sp.
Well, I dug into some of the journals we have handy here, and it would seem that C[sub]2[/sub] is a molecule that occurs in astrophysical environments, among others. It’s bound by about 6.3 eV, with a bond length of about 1.2 Angstroms.
Numbers from Bartlett and Watts, JCP 96 (1992), pg 6073. (They did some theoretical work, and the experimental numbers are given here.)
Thinking about this some more, without any hybridization, the three P orbitals of each atom could overlap at their ends, and the S orbital electrons could be shared as well. While this is almost certainly unstable in practice, I would think this could be modeled, and a bond energy calculated.
I think 1.2 Angstroms is consistent with an sp-hybridized, C-C triple bond, but the fact that it’s a radical might change that.
Yeah, I don’t understand that either. That’s a VERY short double bond, but that’s what everyone tells me it is. I haven’t been able to figure out how they know.
I am doing some research work right now with diatomic carbon in carbon stars which are low temperature stars that exhibit some peculiar phenomenon. All these chemists seem to think the chemical world revolves around stability at 1 atm, room temperature systems, failing to take into consideration that a lot of the universe does not fit that mold.
Perhaps the OP should have been more clear about what was meant by “possible”. Indeed many, many things we think are “impossible” are “possible” … it just requires the right conditions!