How do pi bonds work?

ORBITAL PERSPECTIVE

Pi (#pi#) bonds are made by overlapping two atomic orbitals sidelong, as shown below. In contrast, sigma (#sigma#) bonds are made by overlapping two atomic orbitals head-on.

Either way, this overlap can either be in-phase or out-of-phase.

• The in-phase one (same colors overlapping) is lower in energy and is called the bonding #pi# overlap. It generates a #pi# molecular orbital.
• The out-of-phase one (opposite colors overlapping) is higher in energy and is called the antibonding #pi# overlap. It generates a #pi^"*"# molecular orbital.

A double bond has one #sigma# and one #pi# bond, while a triple bond has one #sigma# and two #pi# bonds.

MOLECULAR ORBITAL DIAGRAM PERSPECTIVE

The MO diagram depiction is:

where:

• #pi_(npx)# is the bonding molecular orbital formed by the in-phase overlap of an #np_x# with an #np_x# atomic orbital.
• #pi_(npy)# is the bonding molecular orbital formed by the in-phase overlap of an #np_y# with an #np_y# atomic orbital.
• #pi_(npx)^"*"# is the antibonding molecular orbital formed by the out-of-phase overlap of an #np_x# with an #np_x# atomic orbital.
• #pi_(npy)^"*"# is the antibonding molecular orbital formed by the out-of-phase overlap of an #np_y# with an #np_y# atomic orbital.

We have three common ways that we can occupy the #pi# and #pi^"*"# molecular orbitals.

• When the #pi# molecular orbitals are filled but the #pi^"*"# ones are not, we have a #pi# bond.
• When both kinds of molecular orbitals are filled, those electrons are nonbonding and are lone pairs.
• When neither kind of molecular orbital is filled, there is no lone pair or bond.

The #sigma# overlaps are the ones that are head-on, and are not our focus (though you should know those as well).