is typical of carbon compounds that have double bonds between atoms.
The carbon-carbon bond length in ethylene is 1.34Å, compared
to 1.54Å in ethane, and the bond energy, or energy required
to pull the carbon atoms apart, is 147 kcal mole,
rather than 83 kcal mole
for ethane. Furthermore, the carbon-carbon double bond is rigid.
No rotation is possible around the bond, and the two carbon atoms
and four hydrogen atoms all are constrained to lie in one plane.
The H-C-H bond angle at either end of the molecule is 117.
How can MO theory account for these characteristics of ethylene?
Ethylene has 12 outer-shell atomic orbitals involved in bonding:
one s and three p orbitals from each carbon, and a Is from each
of the four hydrogens. It also has 12 outer-shell electrons to place
in MO's: four each from the carbons and one each from the hydrogens.
The Is carbon orbitals are filled with electron pairs, do not overlap
appreciably, and play no part in bonding. One solution to the bonding
problem would be to begin with sp3 hybrid orbitals around the carbons,
and to assume that each carbon atom shares two such tetrahedral
orbitals with the other, as shown at the lower right. This is unlikely,
because of the severely bent bonds that would result between carbons.
It is also wrong, because it predicts a H-C-H bond angle of 109.5
instead of the observed 117.