How Lone Pairs Change Molecular Shape

Methane, ammonia and water all have four pairs of electrons around their central atom. Yet methane is a tetrahedron, ammonia is a pyramid and water is bent. The difference is lone pairs: electron pairs that belong to the central atom but aren’t used in bonding. Lone pairs are invisible in the final shape — you only “see” atoms — but they shape the molecule as surely as bonds do. This article explains why lone pairs matter, how much space they take, and where they go in larger molecules.

What is a lone pair?

A lone pair (or non-bonding pair) is a pair of valence electrons on an atom that isn’t shared with another atom. You can find them by drawing a Lewis structure (see Lewis dot structures):

Atom (typical molecule)	Valence electrons	Bonds	Lone pairs
C (CH₄)	4	4	0
N (NH₃)	5	3	1
O (H₂O)	6	2	2
F (HF)	7	1	3

A useful rule for main-group atoms with a full octet: lone pairs = 4 − number of bonds.

Why lone pairs affect shape

According to VSEPR theory, all regions of electron density around a central atom — bonding pairs and lone pairs — repel each other and spread out as far as possible (see VSEPR and molecular geometry). So a lone pair occupies a position in the electron arrangement just like a bond does.

But when we name a molecule’s shape, we describe only where the atoms are. The lone pairs are part of the arrangement but aren’t counted in the shape’s name. That’s why the same electron arrangement can produce several different shapes (see electron geometry vs molecular geometry).

Why lone pairs take up more room

A bonding pair is held between two nuclei, stretched out along the bond. A lone pair is held by only one nucleus. It sits closer to the central atom and spreads out sideways, occupying a fatter, wider region of space.

As a result:

lone pair–lone pair repulsion > lone pair–bonding pair repulsion > bonding pair–bonding pair repulsion

Lone pairs push bonding pairs together, making bond angles smaller than the ideal values (see bond angles explained).

Four regions: tetrahedral family

With four regions of electron density, the electron arrangement is tetrahedral. Replacing bonds with lone pairs gives:

Bonding pairs	Lone pairs	Shape	Angle	Example
4	0	Tetrahedral	109.5°	CH₄, CCl₄, NH₄⁺
3	1	Trigonal pyramidal	about 107°	NH₃, H₃O⁺, PCl₃
2	2	Bent (V-shaped)	about 104.5°	H₂O, H₂S, OF₂
1	3	Linear (only two atoms)	—	HF, HCl

Picture a tetrahedron. Hide one corner (a lone pair) and the three remaining atoms plus the central atom form a pyramid. Hide two corners and you’re left with a bent, V-shaped molecule.

Three regions: trigonal planar family

Bonding pairs	Lone pairs	Shape	Angle	Example
3	0	Trigonal planar	120°	BF₃, CO₃²⁻, NO₃⁻
2	1	Bent	about 117–119°	SO₂, O₃, NO₂⁻

In SO₂ and ozone, the single lone pair on the central atom pushes the two bonds slightly closer than 120°. Ozone’s angle is about 117°.

Five regions: where lone pairs choose to go

With five regions, the electron arrangement is a trigonal bipyramid. It has two different kinds of position:

• Axial (top and bottom): each has three neighbours at 90°.
• Equatorial (around the middle): each has only two neighbours at 90° (the other two are at 120°).

Repulsions at 90° are much stronger than at 120°. So lone pairs always go to the equatorial positions, where they have fewer close neighbours.

Bonding	Lone	Shape	Example
5	0	Trigonal bipyramidal	PF₅, PCl₅
4	1	Seesaw	SF₄
3	2	T-shaped	ClF₃, BrF₃
2	3	Linear	XeF₂, I₃⁻

The results can be surprising: XeF₂ has three lone pairs and five regions of electron density, yet it’s a perfectly linear molecule. The three lone pairs sit around the “equator”, and the two fluorines point straight up and down.

Six regions: octahedral family

With six regions, the electron arrangement is octahedral, and all six positions are equivalent.

Bonding	Lone	Shape	Example
6	0	Octahedral	SF₆
5	1	Square pyramidal	BrF₅, IF₅
4	2	Square planar	XeF₄, ICl₄⁻

With two lone pairs, they go opposite each other (at 180°), minimising their mutual repulsion. That leaves four atoms in a flat square around the central atom.

Seeing lone pairs in real data

Lone pairs can’t be photographed directly, but their effects show up clearly in measurements. The series CH₄ → NH₃ → H₂O shows bond angles closing from 109.5° to about 107° and then 104.5° as lone pairs replace bonds. The ammonium ion, NH₄⁺, is a revealing check: when ammonia’s lone pair is used to bond an extra hydrogen ion, the lone pair disappears and the angle opens back up to exactly 109.5°. The same happens when water becomes the oxonium ion, H₃O⁺: with one lone pair left instead of two, its angle widens towards the 107° of ammonia.

Lone pairs also explain why some molecules turn inside out. The nitrogen in ammonia flips rapidly through a flat arrangement, like an umbrella blowing inside out, billions of times a second. The lone pair passes from one side of the molecule to the other. This “inversion” is so regular that the first atomic clock, built in 1949, used the ammonia inversion frequency as its reference.

Lone pairs and polarity

Lone pairs also affect whether a molecule is polar. In water and ammonia, lone pairs make the charge distribution uneven and help produce a large dipole moment. In symmetrical cases like XeF₂ and XeF₄, lone pairs sit symmetrically, and the molecule is non-polar (see how to tell if a molecule is polar).

Lone pairs in chemistry beyond shape

Lone pairs don’t just set shapes. They’re also responsible for:

• Basicity: ammonia accepts H⁺ using its lone pair (see dative covalent bonds).
• Hydrogen bonding: lone pairs on O, N and F accept hydrogen bonds from δ+ hydrogens (see hydrogen bonding).
• Ligand behaviour: water, ammonia and chloride ions bind to metal ions through lone pairs.
• Nucleophilic attack: lone pairs on O and N attack δ+ carbon atoms in many organic reactions.
• Weak bonds: repulsion between lone pairs on neighbouring atoms weakens F–F, O–O and N–N single bonds (see bond length and bond strength).

Common misconceptions

• “Lone pairs aren’t part of the shape, so ignore them.” They’re essential for working out the shape, even though they aren’t named in it.
• “Ammonia is trigonal planar.” Its lone pair makes it a pyramid.
• “Lone pairs take up the same space as bonds.” They take up more.
• “In five-region molecules lone pairs go anywhere.” They go equatorial.

Key takeaways

• A lone pair is a non-bonding pair of valence electrons on an atom.
• Lone pairs count as regions of electron density in VSEPR, but the shape names only atom positions.
• Lone pairs repel more strongly than bonding pairs and reduce bond angles.
• Tetrahedral family: tetrahedral → trigonal pyramidal → bent.
• In trigonal bipyramids lone pairs go equatorial (seesaw, T-shaped, linear); in octahedra a second lone pair goes opposite the first (square planar).

Practise with VSEPR practice questions.