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A Lewis structure tells you which atoms are connected and where the lone pairs are. It doesn’t tell you what the molecule looks like in three dimensions. Yet shape matters enormously: it decides whether a molecule is polar, how it packs into a solid, and whether a drug fits the protein it’s meant to block.
VSEPR theory — valence shell electron pair repulsion, usually pronounced “vesper” — predicts shape with one simple idea.
The idea
Electron groups around a central atom repel each other and spread out as far apart as possible.
An “electron group” (also called an electron domain) is any region of electrons around the central atom:
- a single bond,
- a double or triple bond (counts as one group),
- a lone pair.
Count the groups and you know how they arrange themselves.
Electron geometry: the five basic arrangements
| Electron groups | Arrangement | Angle |
|---|---|---|
| 2 | Linear | 180° |
| 3 | Trigonal planar | 120° |
| 4 | Tetrahedral | 109.5° |
| 5 | Trigonal bipyramidal | 90°, 120° |
| 6 | Octahedral | 90° |
Molecular shape: what you actually “see”
The molecular shape describes only where the atoms are. Lone pairs take up space and push bonds around, but they’re invisible when you describe the shape. So molecules with the same number of electron groups can have different shapes:
| Groups | Bonds | Lone pairs | Shape | Example |
|---|---|---|---|---|
| 2 | 2 | 0 | Linear | CO₂ |
| 3 | 3 | 0 | Trigonal planar | BF₃ |
| 3 | 2 | 1 | Bent (~119°) | SO₂ |
| 4 | 4 | 0 | Tetrahedral | CH₄ |
| 4 | 3 | 1 | Trigonal pyramidal (~107°) | NH₃ |
| 4 | 2 | 2 | Bent (~104.5°) | H₂O |
| 5 | 5 | 0 | Trigonal bipyramidal | PCl₅ |
| 5 | 4 | 1 | Seesaw | SF₄ |
| 5 | 3 | 2 | T-shaped | ClF₃ |
| 5 | 2 | 3 | Linear | XeF₂ |
| 6 | 6 | 0 | Octahedral | SF₆ |
| 6 | 5 | 1 | Square pyramidal | BrF₅ |
| 6 | 4 | 2 | Square planar | XeF₄ |
Lone pairs squeeze bond angles
Lone pairs are held by only one nucleus, so they spread out more and repel more strongly than bonding pairs. The order of repulsion is:
lone pair–lone pair > lone pair–bond pair > bond pair–bond pair
That’s why methane, ammonia and water — all with four electron groups — have shrinking angles:
- CH₄: 0 lone pairs → 109.5°
- NH₃: 1 lone pair → about 107°
- H₂O: 2 lone pairs → about 104.5°
Water’s bent shape is one of the most consequential facts in chemistry. If water were straight like CO₂, it wouldn’t be polar, wouldn’t hydrogen-bond the way it does, and wouldn’t be the liquid that makes life possible.
The method
- Draw the Lewis structure.
- Count electron groups on the central atom (bonds of any order count as one; each lone pair counts as one).
- Find the electron geometry from the number of groups.
- Find the molecular shape from how many of those groups are lone pairs.
- Adjust angles down slightly for each lone pair.
Worked examples
Carbon dioxide, CO₂. Carbon has two double bonds and no lone pairs → 2 groups → linear, 180°. The two C=O bonds point in opposite directions, so their polarities cancel and CO₂ is nonpolar overall.
Ammonia, NH₃. Nitrogen has three bonds and one lone pair → 4 groups → tetrahedral electron geometry, but the shape is trigonal pyramidal, like a tripod with nitrogen at the top.
Sulfur dioxide, SO₂. Sulfur has two bonds to oxygen and one lone pair → 3 groups → bent, just under 120°.
Xenon tetrafluoride, XeF₄. Xenon has four bonds and two lone pairs → 6 groups → octahedral arrangement. The two lone pairs sit opposite each other to be as far apart as possible, leaving the four fluorines in a flat square planar shape. (Yes, xenon — a noble gas — does form compounds. See the noble gases.)
Where the lone pairs go in 5-group molecules
In the trigonal bipyramid, the three “equatorial” positions (around the middle) have more room than the two “axial” positions (top and bottom). Lone pairs always go equatorial first. That’s why SF₄ is a seesaw, ClF₃ is T-shaped, and XeF₂ ends up perfectly linear with its three lone pairs around the middle.
Limits of VSEPR
VSEPR is a model, not a full theory. It works very well for main-group molecules but poorly for many transition metal compounds, and it doesn’t explain why bonds form — only how the electron groups arrange. For that, chemists use hybridization and molecular orbital theory. As a quick way to predict shape, though, nothing beats it.
Quick answers
Does a double bond count as two groups? No, as one. All the electrons in a multiple bond sit between the same two atoms.
What’s the difference between electron geometry and molecular shape? Electron geometry includes lone pairs; molecular shape describes only the atoms.
Why is water bent but CO₂ linear? Water’s oxygen has two lone pairs (four groups total); CO₂’s carbon has none (two groups).
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