Bond length

Open Question

Chemistry:
Why are Triple Bond lengths shorter than Single and Double Bonds?
In Chemistry class we were told to hypothesize why we think bond lengths get shorter as multiple bonds increase, my question is, can someone explain it to me?I know that bond lengths in pecometers for C - C Single bonds is 154, in C - C Double bonds is 133, and 120 in C - C Triple bonds.

My assumption is that as more electrons become "shared" between the Carbon atoms, the orbitals become closer together. But why?
I know in the single bonds they are sp3, the double bond is sp2 and the triple bond is sp, does orbital shape have anything to do with why they are closer?

And I aslo know there are two pi bonds in the sp hybridized, C - C triple bonds, only one in the sp2 double bond and no pi bonds in the single bonded Carbon atoms. Does pi bonds havre anything to d with bond length?
I guess all this information just has me condused, maybe someone can set me straight?
Thnaks alot!

Answers (1)
by Timothy


There is an important electrostatic potential energy component to chemical bonding. When a bond forms, more electron density accumulates in the region between the positively charged nuclei than would be there if the electron "clouds" of the atoms just stayed the same as they are in isolated atoms. Since the electrons are negatively charged, when the probability of finding the electrons is greater in the region between the nuclei, the positively-charged nuclei are drawn together as they are attracted to the negatively charged electron density.

Thus, the electrons are the negatively charged "glue" that brings the positively charged nuclei together. In double and triple bonds, more electron density piles up in the region between the nuclei, so the nuclei are drawn together even more closely by their attraction to the electron density.

However, π bonds don't build up electron density directly between the nuclei, so π bonds don't add as much to the bond strength as σ bonds which do. π bonds shorten the bonds too, but their shortening effect is not as great as the effect of the σ bonds. Of course, π bonds always add to the effect of the σ bonds and therefore double bonds (σ + one π) and triple bonds (σ + two π) are stronger and shorter than single σ bonds alone.

Q: "I know in the single bonds they are sp³, the double bond is sp² and the triple bond is sp, does orbital shape have anything to do with why they are closer?"

A: Actually, yes. s orbitals are a little more contracted than p orbitals and the optimal distance between atoms for accumulating bonding electron density is consequently a little shorter when sp² hybrids are involved that when sp³ hybrids are involved; when sp hybrids are involved, the bonds a shortened a little bit more than when sp² hybrids are involved.

This small shortening can be seen by comparing the C–H bond distances in acetylene (H–C≡C–H, d(sp-CH) = 106 pm), ethylene (H₂C=CH₂, d(sp²-CH) = 108.7 pm) and ethane (H₃C–CH₃, d(sp³-CH) = 109.4 pm).

The differences in hybridization are less important than the difference due to bond order, however.

Full disclosure: This argument is fine as far as it goes, but there is also an important electronic kinetic energy contribution to bonding that requires a more sophisticated knowledge of the quantum mechanics of chemical bonds. I haven't attempted to explain that...
Source(s):
Chemistry professor