Molecules are 3D, and their shape is not random. Electron pairs are all negative, so they push each other as far apart as they can. Work out how many groups of electrons surround the central atom and the shape falls out automatically — that idea is called VSEPR.
📚 What you need to know
VSEPR = Valence Shell Electron Pair Repulsion. Electron pairs repel and spread out as far as possible in 3D.
An electron domain is one region of electron density: a single bond, a double bond, a triple bond, or a lone pair.
A multiple bond counts as one domain, no matter how many pairs are in it.
Lone pairs repel more strongly than bonding pairs, so they squeeze the bond angles down.
Electron domain geometry counts everything; molecular shape only describes where the atoms are.
The idea in one sentence
All electron pairs are negative, and negative things repel. So the pairs around a central atom arrange themselves to get as far away from each other as possible. Once you know how many groups there are, geometry does the rest — there is only one way to space out two groups, one way to space out three, and one way to space out four.
VSEPR runs on three rules:
🧩 The three VSEPR rules
All electron domains spread out as far apart as possible — bonding pairs and lone pairs alike.
Lone pairs repel more strongly than bonding pairs, so they push the bonds closer together.
A multiple bond counts as a single domain, because all its pairs sit in the same region of space.
Rule 3 is the one people forget. In CO2 there are four shared pairs around the carbon, but they sit in only two regions — one on each side. Two domains means linear, and that is why CO2 is a straight molecule.
Domain geometry is not the same as molecular shape
This distinction wins and loses a lot of marks, so get it straight now.
Electron domain geometry — the arrangement of everything around the central atom, lone pairs included.
Molecular shape — what the molecule actually looks like, which only depends on where the atoms sit. Lone pairs are invisible.
Ammonia is the perfect example. It has four domains, so its domain geometry is tetrahedral. But one of those domains is a lone pair, and you cannot see a lone pair — so the shape you describe is trigonal pyramidal.
Two domains — linear
Two groups get as far apart as possible by sitting on opposite sides of the central atom. That gives a straight line and a bond angle of 180°. Examples: CO2, BeCl2, ethyne (HC≡CH).
Three domains — trigonal planar
Three groups spread out in a flat triangle, all in the same plane, at 120° to each other. If all three are bonding pairs, the shape is trigonal planar. Examples: BF3, ethene, methanal.
Two domains push to opposite sides; three domains spread into a flat triangle.
Three domains with one lone pair — bent
Swap one of those three domains for a lone pair and things change. The lone pair pushes harder, so the two remaining bonds get squeezed to about 118°. And because you cannot see the lone pair, the molecule looks bent, not triangular.
SO2 is the standard example. It has two S=O double bonds and one lone pair on the sulfur — three domains in total. (Sulfur ends up with 10 electrons here, which is fine: it is in period 3 and can expand its octet.)
Four domains — tetrahedral
Four groups cannot all fit in a flat plane and stay far apart, so they push out into 3D and form a tetrahedron with angles of 109.5°. When all four are bonding pairs the shape is tetrahedral: CH4, CCl4, NH4+.
The orange lobe is a lone pair. In SO₂ it squeezes the two bonds down to about 118°. In methane there are no lone pairs, so the full 109.5° is kept. Bold wedges point towards you, hashed bonds point away.
Four domains with lone pairs
Now start replacing bonds with lone pairs and watch the angle shrink each time:
One lone pair → trigonal pyramidal, ~107°.NH3 is the example. Three H atoms form a tripod with the lone pair sitting on top like an invisible fourth leg.
Two lone pairs → bent, ~104.5°.H2O is the example. Two lone pairs pushing means the angle is squeezed even further.
Same tetrahedral arrangement of domains in both, but the lone pairs are invisible — so we describe the shapes as trigonal pyramidal and bent.
Why lone pairs push harder
A bonding pair is held between two nuclei, so it is pulled from both sides and stays fairly tucked in. A lone pair is attached to one nucleus only, so nothing pulls it away — it sits closer to the central atom and takes up more room. More room taken means more pushing, which is why the order of repulsion goes:
Two lone pairs repel each other most strongly of all, and two bonding pairs least.
Every time you swap a bonding pair for a lone pair, the bond angle drops by roughly 2.5°. That is the pattern behind the three most-tested molecules in the whole topic:
Methane, ammonia and water all have four electron domains. The only thing that changes is how many of them are lone pairs.
Summary table
These six rows cover everything you need at Standard Level. If you can reproduce this table from memory, this topic is done.
Bonding pairs
Lone pairs
Total domains
Domain geometry
Molecular shape
Bond angle
2
0
2
linear
linear
180°
3
0
3
trigonal planar
trigonal planar
120°
2
1
3
trigonal planar
bent
~118°
4
0
4
tetrahedral
tetrahedral
109.5°
3
1
4
tetrahedral
trigonal pyramidal
~107°
2
2
4
tetrahedral
bent
~104.5°
🧩 How to answer any shape question
Draw the Lewis formula first. Always. This is where you spot the lone pairs.
Count the domains on the central atom — each bond (single, double or triple) is one, each lone pair is one.
Read off the domain geometry: 2 → linear, 3 → trigonal planar, 4 → tetrahedral.
Ignore the lone pairs and describe where the atoms are — that is your molecular shape.
Adjust the angle down by about 2.5° for each lone pair present.
WORKED EXAMPLE
Deduce the shape and bond angle of the ammonium ion, NH4+.
Step 1 — Lewis formulaNitrogen makes four N–H bonds; the lone pair has been used to form the coordinate bond.Step 2 — count domains: 4 bonding, 0 loneStep 3 — 4 domains → tetrahedral geometryStep 4 — no lone pairs, so the shape is the sametetrahedral, 109.5°The angle is the full 109.5° because there is no lone pair left to squeeze it.
WORKED EXAMPLE
CO2 and H2O both have a central atom bonded to two other atoms. Explain why one is linear and the other is bent.
CO₂ — carbon has 2 double bonds, no lone pairsEach double bond counts as ONE domain, so there are only 2 domains.linear, 180°H₂O — oxygen has 2 bonds plus 2 lone pairsThat is 4 domains → tetrahedral geometry, but 2 of them are invisible.bent, ~104.5°The difference is entirely down to the lone pairs on the oxygen.
💡 Exam tip
For the bent shape the IB accepts several names: bent, V-shaped, angular or non-linear. Any of them will do.
Never guess a shape from the formula alone. Draw the Lewis structure — H2O and BeCl2 are both AB2 but have completely different shapes.
If a question asks for the shape and the angle, give both. Two separate marks.
Write ~107° or about 107° for shapes with lone pairs — those angles are approximate, unlike the exact 180°, 120° and 109.5°.
⚠️ Common mix-up
Domain geometry vs molecular shape. NH3 is tetrahedral in its domains but trigonal pyramidal in shape. If the question says “shape of the molecule”, lone pairs don’t count.
A double bond is one domain, not two. Count regions, not pairs.
Don’t call NH3 “tetrahedral” and don’t call H2O “tetrahedral” — you can’t see the lone pairs, so you can’t name the shape after them.
Trigonal pyramidal is not the same as trigonal planar. Planar is flat with 120°; pyramidal is 3D with ~107°.
That completes the drawing-and-shapes half of covalent bonding. Next comes Bond Polarity — what happens when the two atoms in a bond don’t pull on the shared pair equally.
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