Bonding
The type of bonding can be explained in terms of orbital hybridization. In ethylene each carbon atom has three sp2 orbitals and one p-orbital. The three sp2 orbitals lie in a plane with 120° angles. The p-orbital is perpendicular to this plane. When the carbon atoms approach each other, two of the sp2 orbitals overlap to form a sigma bond. At the same time, the two p-orbitals approach (again in the same plane) and together they form a pi-bond. For maximum overlap, the p-orbitals have to remain parallel, and, therefore, rotation around the central bond is not possible. This property gives rise to cis-trans isomerism. Double bonds are shorter than single bonds because p-orbital overlap is maximized.
| 2 sp2 orbitals (total of 3 such orbitals) approach to form a sp2-sp2 sigma bond | Two p-orbitals overlap to form a pi-bond in a plane parallel to the sigma plane |
With 133 pm, the C=C bond length is shorter than the C−C length in ethane with 154 pm. The double bond is also stronger, 636 (KJ/mol) versus 368 kJ/mole but not twice as much as the pi-bond is weaker than the sigma bond due to less effective pi-overlap.
In an alternative representation, the double bond results from two overlapping sp3 orbitals as in a bent bond.
Read more about this topic: Double Bond
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