When working geometry problems by the VSEPR method, a
double or triple bond acts just like a single bond for geometry purposes: the
added electrons don't have any effect on the molecular geometry. This is because
the extra electrons in multiple bonds do not reside in hybrid
For example, the ethylene molecule has the Lewis diagram below.
bond angles around carbon in ethylene are 120o, which implies an AX3
geometry and sp2 hybridization. One of the two shared electron pairs
on a carbon in the molecule (blue) is in an sp2 hybrid orbital as are
the two electron pairs that form bonds to the hydrogen, the other (red) is not.
The bonding pair that is in the hybrid orbital is said to reside in a s bond.
So what type of orbital are the other two electrons in? It turns out that
they are in the remaining unhybridized p orbitals on the carbon. (The
sp2 hybrid uses one s and two p orbitals, leaving a p orbital on each
carbon.) These two p orbitals overlap, forming a p
bond. This orbital sticks up out of the plane of the molecule, as in the picture
In the case of acetylene, HCCH, we have the Lewis structure below
we have sp hybridization and an AX2
geometry around the carbons. There are two p orbitals left over on each carbon:
both overlap and thus form two p bonds (red) along with
the (blue) s bond that results from the electrons in sp
hybrid orbitals. Some pictures of the two sets of p
bonds are below