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 orbitals

For example, the ethylene molecule has the Lewis diagram below.

The bond angles around carbon in ethylene are 120^o, 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 below

In the case of acetylene, HCCH, we have the Lewis structure below

Here, 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





 

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