σ and π bonds: head-on versus sideways overlap

5 minAS Level3.4.2

AI-drafted, independently checked · not yet teacher-reviewed

By the end of this lesson

By the end you can describe how σ and π bonds form by orbital overlap, and count the σ and π bonds in $\text{H}_2$ , $\text{C}_2\text{H}_6$ , $\text{C}_2\text{H}_4$ , $\text{HCN}$ and $\text{N}_2$ .

A double bond is not two copies of a single bond

Dot-and-cross diagrams tell you how many pairs are shared. They don't tell you where those electrons are.

To see that, zoom in to orbitals. A covalent bond forms when an orbital from each atom overlaps, and the shared pair sits in the overlap.

Two ways to overlap

σ (sigma) bond: direct (head-on, end-on) overlap of orbitals between the bonding atoms. The electron density is concentrated on the line joining the two nuclei.
π (pi) bond: sideways overlap of two adjacent p orbitals. The electron density is above and below the σ bond, not on the line joining the nuclei.

Quick check

Multiple choiceQuestion 1 · Recall

Which statement describes how a σ bond forms?

ASideways overlap of two adjacent p orbitalsBDirect (head-on) overlap of orbitals between the two bonding atomsCComplete transfer of an electron from one atom to the otherDOverlap of orbitals above and below the line joining the nuclei

Fill in the blanksQuestion 2 · Recall

A π bond forms by the overlap of two adjacent orbitals, with the electron density above and below the σ bond.

Answer the 2 questions above to continue.

Previous lessonCoordinate (dative) bonds: one atom brings both electrons Next lessonHybridisation: sp³, sp² and sp orbitals

Head-on gives σ, between the nuclei. Sideways gives π, above and below.