I asked my chem teacher about this, and he didn’t know. Does a quarduple-bonded carbon pair exist, or would it be too unstable? Is there a name for it? I.e., alkyne=3
Thanks!
I believe a quadruple-bonded carbon-carbon pair is not possible not just because of the high level of strain imposed on the 3 bonding p-orbitals but also because as the number of bonds between a pair of atoms increases, the distance between the nuclei decreases. In other words, you’d be trying to bring the two positively charged nuclei much too close to each other and there would be too much strain due to like charge repulsion for the molecule to exist.
I’ll second that opinion. Too much steric(sp?) strain. I’m just a chem student, though.
I know from nothing about chemistry, but I know that carbon is present in only three forms and that isn’t one of them.
It’s been almost a quarter century since I looked at a chem text, but I do remember that carbon bonds tetrahedrally. So I can envision the structural hookup for three of the bonds, but the fourth would have to be a(n impossible) reach around.
What beatle said. This isn’t exactly accurate, but think of it like this: Take two tetrahedra and get three of their points to overlap. IOW, take two pyramids and put them base to base. The two fourth points are then pointing directly away from each other, and so cannot overlap.
As I said, that’s not entirely accurate, what with orbital hybridization and all, but it’s along those lines.
This thread addressed the same issue.
For what it’s worth, C[sub]2[/sub] can exist, but it’s only doubly bonded. I don’t remember the name for it though.
Really? C[sub]2[/sub] by itself? As a stable molecule? Not C[sub]2[/sub]H[sub]6[/sub]? If it’s only doubly bonded, is it an unstable transition molecule that soon picks up something to plug into loose ends?
Who said anything about stability? 
Actually, I’ll have to try to remember to dig up some references on it. I KNOW that theoretical chemists play around with this system periodically, and I could swear that it’s something along the lines of a metastable gas phase beast. I might, however, be wrong. In any event, regardless of whether it’s been seen experimentally, it does have a bound state. It might be so reactive that you’d be hard pressed to see it except as an intermediate (I suspect this is the case).
What beatle guessed is right. C[sub]2[/sub] is a species that exists transiently in, for example, flames that come from burning carbon-containing material. I’m pretty sure it’s a radical and not a stable molecule and not quadruply bonded. I think it’s responsible for the color of a candle flame. It gives off some particular wavelength of light that’s used in studying combustion, anyway. People use that wavelength to analyze the performance of rocket engines.
Actually, there are more than three stable forms of carbon. The common ones are diamond, graphite and fullerene, but there is at least one other. In acetylenic carbon the atoms are in a straight chain with alternating single and triple bonds.
For those of you who can’t get enough of this subject and are clamoring for more, here’s a good page about carbon allotropes. It talks about another allotrope I’d never heard of before, called lonsdaleite.
I’m not chemist or even student of chemistry, but I think the other common form of carbon is the benzene ring. Am I right?
quote:
"I asked my chem teacher about this, and he didn’t know. "
holy shit your chemistry teacher is dense. no it is not possible.
Actually, benzene (C[sub]6[/sub]H[sub]6[/sub]) contains hydrogen. But in fullerenes, all of the carbons belong to benzene rings that are saturated with other benzene rings instead of hydrogen.
::dusting off chemistry degree::
Carbon atoms usually undergo sp[sup]3[/sup] hybridization to bond. While they can bond without it, the bonds are weak and have resonance structures to keep them stable. In sp[sup]3[/sup] hybridization, the bond angles tend to be around 109 degrees. Looking at the “perfect example” of sp[sup]3[/sup] hybridization, methane, we have for equal bonds that repel each other equally until they have the largest 3d space between them. This gives the tetrahedral shape with bond angles of ~109.
Considering the Valence Orbital theory of bonding, these areas of high electron density must overlap to form a bond (region of high density in C and region of high density in H). The maximum number of orbitals that can overlap between two carbons (or any two sp[sup]3[/sup] hybridizing atoms) is three. The fourth bond can not form.
Now, what if we energize the electrons to the point of moving them so that the fourth regions of high density can overlap and form a bond. Here we get into the realm of possibilities. However, it takes gobs of energy and thus is incredibly unstable - the fourth bond breaking as a result of its forming.
No need to invoke sp[sup]3[/sup] hybridization, because that only happens when carbon has 4 single bonds. If you start with a triply bonded carbon pair, then the 4th bond must come from p orbitals that are pointing 180 degrees away from each other. There is no overlap and therefore no possiblity of a bond.
This isn’t to say that quadruple bonds are impossible, but they don’t happen with just s and p orbitals. D orbitals are needed.
bad niobium knight! bad! we dont use naughty words in GQ, now do we? go sit in the corner for five minutes and think about what you’ve done.
Sorry for any confusion, Dr. Lao; I mentioned the hybridization to demonstrate the bond angles and orientations in space.
If I’m reading my chemistry book correctly, for triple bonded carbon, hybridization between one P orbital and the S orbital takes place. The large lobes of one SP orbital from each atom overlap to form a sigma bond. The two unhybridized P orbitals from each atom are oriented perpendicular to the line between the carbon atoms, and form pi bonds. It’s the large lobes of the second SP orbital from each atom that are pointed away from each other, preventing a fourth bond.