Introduction
Pi bonds are responsible for many of the most important structures and reactions in organic chemistry, including alkenes, alkynes, conjugation, and resonance. But what exactly is a pi bond, and how does it form?
In this lesson, we'll build directly on the concepts used to form sigma bonds. Using the linear combination of atomic orbitals (LCAO), we'll see how p orbitals combine constructively and destructively to produce bonding and antibonding molecular orbitals. Understanding this process reveals what a pi bond really is at the molecular orbital level and why every orbital overlap produces both a bonding and an antibonding orbital.
From Sigma Bonds to Pi Bonds
If you've already seen how two s orbitals overlap to form a sigma bond, the process for creating a pi bond will feel very familiar.
The core idea remains the same:
- Orbitals are solutions to mathematical equations.
- Orbitals can be combined using linear combination of atomic orbitals.
- The overlap must occur both constructively and destructively.
- The number of orbitals going in must equal the number of orbitals coming out.
When two p orbitals overlap, we use exactly the same rules.
Constructive Overlap of p Orbitals
To form a pi bonding orbital, two p orbitals overlap so that matching phases align with one another.
Each p orbital contains:
- One lobe of positive phase
- One lobe of negative phase
When positive overlaps positive and negative overlaps negative, the overlap is maximally constructive.
Just like constructive interference between waves, this increases the amplitude of the resulting orbital.
Because orbital amplitude is related to the probability of finding electrons, constructive overlap increases electron density in the region where the orbitals overlap.
The Pi Bonding Orbital
The result of constructive overlap is the pi (π) bonding orbital.
In this molecular orbital:
- Electron density is increased between the nuclei.
- The probability of finding electrons in the overlap region is higher.
- The molecular orbital is lower in energy than the original atomic orbitals.
This lower-energy orbital is the bonding orbital formed from the overlap of the two p orbitals.
Why There Must Be a Second Orbital
The pi bonding orbital cannot be the whole story.
According to linear combination of atomic orbitals, if we start with two atomic orbitals, we must finish with two molecular orbitals.
So far we've only produced one molecular orbital through constructive overlap.
To satisfy LCAO, we must also perform the destructive combination.
Destructive Overlap of p Orbitals
Destructive overlap occurs when opposite phases are aligned.
One p orbital is combined with the negative phase of the other so that:
- Positive overlaps negative
- Negative overlaps positive
This creates destructive interference.
As the orbitals cancel each other, the probability of finding an electron between the nuclei drops to zero in a specific region.
The Nodal Plane and Pi-Star Orbital
The destructive overlap creates a nodal plane.
A nodal plane is a region where:
- The mathematical solution equals zero.
- The probability of finding an electron is zero.
This molecular orbital is called the pi-star (π) antibonding orbital*.
Unlike the bonding orbital, the π* orbital is higher in energy and does not stabilize the interaction between the two atoms.
Bonding and Antibonding Orbitals Always Come Together
One of the most important lessons from molecular orbital theory is that orbitals never overlap to create just one new orbital.
Whenever two atomic orbitals interact, they always produce:
A Bonding Orbital
- Lower energy
- Constructive overlap
- Increased electron density
- Stabilizing
An Antibonding Orbital
- Higher energy
- Destructive overlap
- Contains a node or nodal plane
- Destabilizing
This is true for sigma bonds, pi bonds, and any other molecular orbital interaction.
What a Pi Bond Really Is
At its most fundamental level, a pi bond is the bonding molecular orbital produced when two p orbitals overlap constructively.
Using linear combination of atomic orbitals:
- Two atomic p orbitals combine constructively to form a lower-energy π bonding orbital.
- The same two p orbitals combine destructively to form a higher-energy π* antibonding orbital.
- Both molecular orbitals are required because the number of orbitals is conserved.
This bonding-antibonding pair is the molecular orbital explanation for how p orbitals create pi bonds in organic molecules.
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