VSEPR Theory: Why Molecules Have the Shapes They Do
Imagine you're in a crowded room. Everyone wants their personal space. If you're standing with a few friends, you'll naturally spread out so no one is too close to anyone else. That's exactly what happens inside a molecule.
The Core Intuition
Electron pairs are negatively charged. They repel each other. In a molecule, the electron pairs around a central atom will arrange themselves as far apart as possible — just like those people in the room. This simple idea is the entire foundation of VSEPR (pronounced "ves-per") Theory.
VSEPR stands for Valence Shell Electron Pair Repulsion. The name tells you exactly what it's about: the repulsion between electron pairs in the valence shell.
The Precise Statement
VSEPR Theory states that the geometry around a central atom is determined by minimizing the repulsion between all electron pairs (both bonding and lone pairs) in its valence shell.
Two key points to hold onto:
- All electron pairs repel — whether they are shared (bonding pairs) or unshared (lone pairs).
- Lone pairs repel more strongly than bonding pairs. A lone pair is "fatter" — it's only attracted to one nucleus, so it spreads out more and pushes harder on its neighbours.
How to Predict Shape in 3 Steps
Step 1: Count the total electron pairs around the central atom.
Add the number of atoms bonded to the central atom plus the number of lone pairs on it. This gives you the steric number.
Step 2: Arrange those pairs as far apart as possible.
This gives you the electron-pair geometry — the shape if you pretend all pairs are identical.
Step 3: Replace lone pairs with "invisible" space.
The actual molecular geometry is the shape formed by the atoms alone, ignoring lone pairs.
The Common Geometries at a Glance
| Steric Number | Electron-Pair Geometry | Lone Pairs | Molecular Geometry | Example | Bond Angle |
|---|
| 2 | Linear | 0 | Linear | CO2 | 180° |
| 3 | Trigonal planar | 0 | Trigonal planar | BF3 | 120° |
| 3 | Trigonal planar | 1 | Bent | SO2 | ~119° |
| 4 | Tetrahedral | 0 | Tetrahedral | CH4 | 109.5° |
| 4 | Tetrahedral | 1 | Trigonal pyramidal | NH3 | ~107° |
| 4 | Tetrahedral | 2 | Bent | H2O | ~104.5° |
| 5 | Trigonal bipyramidal | 0 | Trigonal bipyramidal | PCl5 | 90°, 120° |
| 6 | Octahedral | 0 | Octahedral | SF6 | 90° |
A common mistake: thinking that NH3 is tetrahedral. It has tetrahedral electron-pair geometry, but because one position is a lone pair, the molecular shape is trigonal pyramidal. The bond angle is 107°, not 109.5°.
Why Lone Pairs Squeeze Bond Angles
Take water (H2O). The central oxygen has 4 electron pairs: 2 bonding (to H atoms) and 2 lone pairs. The ideal tetrahedral angle is 109.5°. But the two lone pairs push harder on the bonding pairs, compressing the H–O–H angle to about 104.5°.
In ammonia (NH3), there's only one lone pair, so the compression is less — the H–N–H angle is about 107°. …