Sigma and Pi Bonds

Going back to the example of BeH 2 , one envisions an sp hybridized orbital overlapping with the s
orbitals of the two hydrogen atoms to form bonds. These kinds of bonds, which result from end-to-
end overlap of orbitals from the two bonded atoms, are known as sigma bonds (σ bonds). When
multiple bonds are involved, however, we do not have several σ bonds between two atoms. What we
have instead is another kind of bond known as a pi bond (π bond), which has a very different spatial
arrangement. A useful example to consider is that of the ethene molecule, H 2 C=CH 2 , in which the
carbon atoms are double-bonded to each other and each of them is also bonded to two hydrogen
atoms. Each carbon atom, then, is bonded to three groups—a CH 2 group and two H atoms, and
VSEPR theory would lead us to predict that the three groups are arranged about 120° apart:

This is in fact the case. The carbon atoms are both sp^2 hybridized. For each of the two carbon atoms,
two of the hybrid orbitals are used to form the bonds with the hydrogen atoms, with the remaining
one forming a σ bond with the leftover hybrid orbital from the other carbon atom.

Where does one get a double bond? It is not the case that each hybrid orbital gives a bond, thus
giving two: Two hybrid orbitals are needed to interact to give one bond. The second bond that is
needed to give a double bond comes from the interaction of the p orbitals that are left unused in
hybridization (one on each carbon atom). Remember that three sp^2 hybrid orbitals are formed by
mixing an s and two p orbitals. Since there are actually three p orbitals, one is left over, sticking out
of the plane that contains the hybrid orbitals. (In this case, then, the unhybridized p orbitals are
coming out at you.) These two unhybridized orbitals from different carbon atoms can interact and