Shapes of molecules
Shapes of molecules
- To predict a shape, use VSEPR theory.
- Electron pairs around the central atom push apart as far as possible.
- Lone pairs push hardest, squeezing the angles.
Practice
VSEPR theory says the shape of a molecule is set by:
Like charges repel, so the electron pairs push apart to minimise repulsion, fixing the shape.
Lone pairs push harder
- A lone pair (not in a bond) repels more strongly than a bonding pair.
- Each lone pair squeezes the bond angle by about $2.5°$.
- That's why $\text{NH}_3$ ($107°$) and $\text{H}_2\text{O}$ ($104.5°$) have smaller angles than $\text{CH}_4$ ($109.5°$).
Practice
Why is the bond angle in water (104.5°) smaller than in methane (109.5°)?
Lone pairs repel more than bonding pairs; water's two lone pairs squeeze the H–O–H angle below 109.5°.
The shapes
| Molecule | Shape | Angle |
|---|---|---|
| $\text{CO}_2$ | linear | $180°$ |
| $\text{BF}_3$ | trigonal planar | $120°$ |
| $\text{CH}_4$ | tetrahedral | $109.5°$ |
| $\text{NH}_3$ | pyramidal | $107°$ |
| $\text{H}_2\text{O}$ | bent | $104.5°$ |
| $\text{SF}_6$ | octahedral | $90°$ |
Practice
Methane (CH₄) is:
Four bonding pairs and no lone pairs give a tetrahedral shape at 109.5°.
Practice
Match each molecule to its shape.
CO₂ = linear (180°); BF₃ = trigonal planar (120°); NH₃ = pyramidal (107°).
You've got it
Key idea
- VSEPR: electron pairs repel and spread out as far as possible
- lone pairs push harder than bonding pairs (≈ $2.5°$ each)
- linear $180°$, trigonal planar $120°$, tetrahedral $109.5°$, pyramidal $107°$, bent $104.5°$, octahedral $90°$