How do double bonds consist of sigma and pi bonds?

Mar 1, 2017

Recall that a sigma bond is a single bond where two orbitals have significant overlap. For alkanes (hydrocarbons with single bonds only) all carbon atoms have an $s {p}^{3}$ orbital hybridization. The $s {p}^{3}$ arrangement creates 4 hybridized orbitals that are ${109.5}^{\circ}$ apart from one another with a tetrahedral geometry.

Consider Ethane ($C H 3 C H 3$)

Ethane has 4 sigma bonds for each carbon (one to the other carbon, three to each hydrogen). With the $s {p}^{3}$ hybridization, there is no way to a double bond between the carbons without causing significant strain and interference.

Consider the orbitals for ethene ($C {H}_{2} C {H}_{2}$), an alkene with a double bond between the two carbons.

The carbon atoms in ethene have an $s {p}^{2}$ hybridization meaning that there are 3 hybridized orbitals in a trigonal planar geometric arrangement with ${120}^{\circ}$ angles between them and one unybridized p orbital (the ${p}_{y}$ orbital). The $s {p}^{2}$ hybridization facilitates a double bond between the carbons because it allows an $s {p}^{2}$ orbital to overlap through a sigma bond and the unhybridized ${p}_{y}$ orbitals for each carbon to also overlap. This p orbital overlap is a pi bond.

Although a pi bond is higher energy than a sigma bond, the combination of a single sigma bond and a single pi bond in a double bond is much lower energy overall than trying to force two sigma bonds between two atoms.