VSEPR & Molecular Geometry

Electron pairs repel — so they spread out as far as possible. That single idea predicts the 3-D shape and bond angles of almost any molecule.

High schoolIntro Gen ChemUni Year 1
⏱️ About 22 min

Tie a few balloons together at their necks and they automatically fan out into a neat, symmetric shape — they can't help it, because they push each other apart. Electron pairs around an atom do exactly the same thing, and that stubborn mutual repulsion is all you need to predict a molecule's shape.

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The big idea: VSEPR (Valence-Shell Electron-Pair Repulsion) says the groups of electrons around a central atom arrange themselves as far apart as possible. Count the electron domains, and the geometry — and its bond angles — follows directly. Lone pairs push harder than bonds, bending the shape.
🎯 By the end, you'll be able to
  • Count the electron domains (bonded atoms + lone pairs) around a central atom
  • Map domain count to geometry: linear, trigonal planar, tetrahedral, and beyond
  • State the ideal bond angles: 180°, 120°, 109.5°
  • Explain how lone pairs distort a shape (bent water, pyramidal ammonia)
📎 Helpful to know first

Count the domains, get the shape

Start from the Lewis structure. Around the central atom, count the electron domains — each domain is either a bonded atom or a lone pair. A crucial shortcut: a double or triple bond still counts as one domain, because all that electron density points in a single direction.

Since domains repel, they space themselves as far apart as geometry allows. The number of domains alone fixes the arrangement:

🔑 Domains → geometry → angle
2 domains → linear, 180°. 3 domains → trigonal planar, 120°. 4 domains → tetrahedral, 109.5°. 5 domains → trigonal bipyramidal (90° & 120°). 6 domains → octahedral, 90°. This is the electron geometry — the arrangement of all domains, bonds and lone pairs alike.

Lone pairs bend the picture

The molecular shape is what we actually see — the arrangement of the atoms only. Lone pairs are invisible to that shape, yet they still occupy a domain and still push. And they push harder than bonding pairs, because a lone pair is held by just one nucleus and spreads out more.

So lone pairs squeeze the remaining bond angles below their ideal values. Take four domains (tetrahedral, 109.5° ideal):

  • 0 lone pairs — methane CH₄: tetrahedral, angles 109.5°.
  • 1 lone pair — ammonia NH₃: trigonal pyramidal, angles squeezed to ~107°.
  • 2 lone pairs — water H₂O: bent, angles squeezed to ~104.5°.
✨ Same domain count, different shape names
CH₄, NH₃ and H₂O all have four electron domains and the same underlying tetrahedral electron geometry. What differs is how many of those domains are lone pairs — which is why we give the visible atom arrangement its own name (tetrahedral → pyramidal → bent) and why the bond angle shrinks step by step (109.5° → 107° → 104.5°).
\[ \underset{109.5^\circ}{\ce{CH4}}\quad>\quad\underset{\sim107^\circ}{\ce{NH3}}\quad>\quad\underset{\sim104.5^\circ}{\ce{H2O}} \]
As lone pairs replace bonds on a 4-domain centre, the bond angle is compressed step by step.
📝 Worked example: Predict the molecular shape and approximate bond angle of ammonia, NH₃.
  1. From the Lewis structure, nitrogen has 3 bonds (to H) plus 1 lone pair = 4 electron domains.
  2. Four domains arrange tetrahedrally (electron geometry = tetrahedral, ideal 109.5°).
  3. One domain is a lone pair, so the visible atom shape is trigonal pyramidal — the three H atoms form a tripod under the nitrogen.
  4. The lone pair pushes harder than the bonds, compressing the H–N–H angle from 109.5° to about 107°.
✓ Trigonal pyramidal, with H–N–H bond angles of roughly 107°.

Why CO₂ is straight but H₂O is bent

It's tempting to think every triatomic molecule looks the same. It doesn't. CO₂ has a central carbon with two double bonds and no lone pairs — just 2 domains — so it is linear at 180°. Water has a central oxygen with 2 bonds and 2 lone pairs — 4 domains — so it is bent at ~104.5°. Same three atoms, totally different shape, entirely because of the lone pairs.

✏️ Practice: In carbon dioxide (O=C=O), how many electron domains surround the central carbon atom? (Remember: each double bond counts as one domain, and carbon has no lone pairs here.)
domains
Solution
  1. Carbon is bonded to two oxygens; each is a double bond.
  2. A double bond counts as a single domain, and carbon has no lone pairs.
  3. So there are 2 domains — which is why CO₂ is linear (180°).

Check your understanding

1. For counting electron domains in VSEPR, a double bond counts as…
All the electron density of a double (or triple) bond points in one direction, so it counts as a single domain when predicting shape.
2. Water (H₂O) is bent rather than linear because…
Oxygen has 4 domains (2 bonds + 2 lone pairs). The lone pairs occupy space and repel the bonds, bending the molecule to about 104.5° instead of 180°.
3. What is the ideal bond angle for a molecule with four bonding domains and no lone pairs (like CH₄)?
Four equivalent domains spread out to a perfect tetrahedron, giving bond angles of 109.5°.
🔎 Related lessons
Previous: Lewis Structures Next: Hybridization & Polarity Intermolecular Forces
✅ Key takeaways
  • VSEPR: electron domains around a central atom repel and spread as far apart as possible.
  • Domain count sets geometry — 2: linear 180°, 3: trigonal planar 120°, 4: tetrahedral 109.5°.
  • A double or triple bond counts as one domain.
  • Lone pairs push harder than bonds, compressing angles: CH₄ 109.5° → NH₃ ~107° → H₂O ~104.5°.
  • Molecular shape describes the atoms only, so lone pairs turn tetrahedral into pyramidal or bent.
➡️ VSEPR predicts the shape; the next lesson explains the orbitals behind it — hybridization — and then uses that shape to decide the property that governs how molecules interact: polarity.
Want to test yourself on this? Try the Chemistry practice test →