Orbital Overlap & Covalent Bond Types
This is Higher Level (HL) content.
Definitions and Orbital Overlap
Covalent bonds result from the overlap of atomic orbitals. The way these orbitals overlap determines whether a sigma or pi bond is formed.
σ (Sigma) Bonds (Head-on Overlap)
Formed by the direct, head-on (axial) combination of atomic orbitals. The electron density is concentrated directly between the nuclei, along the internuclear bond axis.
Can form from the overlap of two s orbitals, two p orbitals, an s and a p orbital, or hybrid orbitals.
π (Pi) Bonds (Lateral Overlap)
Formed by the lateral (side-on) combination of adjacent, unhybridised p orbitals. The electron density is concentrated in two regions above and below the plane of the bond axis.
Only unhybridised p orbitals can form pi bonds.
Sigma vs Pi: Key Differences
| Sigma (σ) Bond | Pi (π) Bond | |
|---|---|---|
| Overlap type | Head-on (axial, end-to-end) | Lateral (side-to-side) |
| Orbitals involved | s-s, s-p, p-p, or hybrid-hybrid | Unhybridised p-p only |
| Electron density | Along the bond axis | Above & below the axis |
| Bond strength | Stronger (greater degree of overlap) | Weaker (less orbital overlap) |
| Rotation | Free rotation | Prevents rotation |
Key Properties
Relative Strength (σ vs π)
Sigma bonds are stronger than pi bonds because the direct, head-on overlap of the orbitals allows for a greater degree of orbital overlap. Pi bonds are generally weaker due to a significantly lower degree of lateral overlap.
Bond Rotation & Rigidity
Molecules with only sigma bonds allow for free rotation around the bond axis. Pi bonds restrict rotation because rotating the atoms would break the parallel overlap of the p orbitals, disrupting the electron cloud. This restricted rotation gives rise to geometric (cis-trans) isomers.
Bond Composition
Rules for Counting σ and π Bonds
The first bond that forms between two atoms is always a sigma bond. Any additional bonds within a double or triple bond are pi bonds.
- Single bond = 1σ
- Double bond = 1σ + 1π
- Triple bond = 1σ + 2π
Worked Examples: Organic Molecules
Rules for Counting σ and π Bonds
Ethane (C₂H₆)
1 C-C single bond + 6 C-H single bonds
Total: 7σ bonds, 0π bonds
Rules for Counting σ and π Bonds
Ethene (C₂H₄)
1 C=C double bond + 4 C-H single bonds
Total: 5σ bonds, 1π bond
Rules for Counting σ and π Bonds
Ethyne (C₂H₂)
1 C≡C triple bond + 2 C-H single bonds
Total: 3σ bonds (1 from C-C + 2 from C-H) and 2π bonds (both in the C≡C)
Worked Examples: Inorganic Molecules
| Molecule | Structure | σ bonds | π bonds |
|---|---|---|---|
| H₂ | H-H single bond (s-s overlap) | 1 | 0 |
| O₂ | O=O double bond | 1 | 1 |
| N₂ | N≡N triple bond | 1 | 2 |
| CO₂ | O=C=O (two double bonds) | 2 | 2 |
Do Not Forget C-H Sigma Bonds When Counting
Every bond has exactly one σ component: the first bond formed is always σ. Pi bonds are additional bonds only. A "triple bond" is NOT "3 bonds of equal strength."
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