Bond Angles: Why Water Is Bent at 104.5°

Why is a water molecule bent at 104.5°, and not straight (180°), or at the tetrahedral angle of 109.5°? The answer comes from a simple idea — electron pairs repel each other — refined by a few extra factors: lone pairs, electronegativity, the size of the central atom and multiple bonds. This article explains each factor with real measured angles, so you can predict and justify bond angles rather than just memorise them.

What a bond angle is

A bond angle is the angle between two bonds that share an atom — for example, the H–O–H angle in water. Bond angles are measured using techniques such as microwave spectroscopy (for gases), electron diffraction and X-ray crystallography. They’re averages, because molecules are always vibrating.

Factor 1: electron pairs spread out

The foundation is VSEPR theory: regions of electron density around a central atom (bonding pairs and lone pairs) repel each other and arrange themselves as far apart as possible (see VSEPR and molecular geometry).

Regions	Arrangement	Ideal angle
2	Linear	180°
3	Trigonal planar	120°
4	Tetrahedral	109.5°
5	Trigonal bipyramidal	90°, 120°
6	Octahedral	90°

The tetrahedral angle is exactly 109.47° — the angle you get by pointing four lines from the centre of a tetrahedron to its corners. Methane, CH₄, with four identical bonding pairs and no lone pairs, has exactly this angle.

Factor 2: lone pairs repel more than bonding pairs

Water’s oxygen is surrounded by four regions of electron density: two bonding pairs (to H) and two lone pairs. So the electron pairs are arranged roughly tetrahedrally. But the H–O–H angle is 104.5°, not 109.5°. Why smaller?

A lone pair is held by only one nucleus, so it sits closer to the central atom and spreads out more widely than a bonding pair, which is stretched out between two nuclei. A lone pair therefore takes up more space and repels more strongly. The order of repulsion is:

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

The lone pairs push the bonding pairs closer together. Each lone pair reduces the bond angle by roughly 2–2.5°:

Molecule	Bonding pairs	Lone pairs	Shape	Bond angle
CH₄	4	0	Tetrahedral	109.5°
NH₃	3	1	Trigonal pyramidal	107°
H₂O	2	2	Bent	104.5°

This pattern — 109.5°, 107°, 104.5° — is one of the most commonly examined in chemistry (see methane, ammonia and water compared and the shape of a water molecule).

Factor 3: electronegativity of the atoms

Bond angles also depend on where the bonding electrons sit.

Electronegative outer atoms → smaller angles

If the atoms bonded to the centre are very electronegative, they pull the bonding pairs away from the central atom. The bonding pairs are then further from each other and repel less, so the angle can close up.

Molecule	Angle
NH₃	107.8°
NF₃	102.2°
H₂O	104.5°
OF₂	103.1°

Fluorine pulls the bonding electrons towards itself, so the F–N–F angle is smaller than H–N–H.

Electronegative central atom → larger angles

If the central atom is more electronegative, it pulls bonding pairs closer to itself, where they repel each other more strongly, opening the angle. Compare hydrides down group 15 and group 16:

Molecule	Angle	Molecule	Angle
NH₃	107.8°	H₂O	104.5°
PH₃	93.5°	H₂S	92.1°
AsH₃	92.0°	H₂Se	91.0°
SbH₃	91.7°	H₂Te	90.3°

Going down the group, the central atom becomes larger and less electronegative. Its bonding pairs are further apart and further from the nucleus, so they repel less, and the angle closes towards 90°.

Factor 4: the size of the central atom and hybridisation

The dramatic drop from NH₃ (107.8°) to PH₃ (93.5°), and from H₂O to H₂S, is too large to explain by electronegativity alone. For larger central atoms, there’s another effect: they use less hybridisation.

Hybridising s and p orbitals costs energy. For small period-2 atoms like N and O, the gain in bond strength makes it worthwhile, so they use roughly sp³ hybrids with tetrahedral-like angles. For larger atoms like P and S, the s and p orbitals are more different in size and energy, so mixing them gives less benefit. The bonds are formed mainly from p orbitals, which are at 90° to each other. That’s why PH₃ and H₂S have angles close to 90° (see hybridisation explained).

Factor 5: multiple bonds

A double or triple bond contains more electron density than a single bond, so it repels neighbouring bonds more strongly. This opens the angles next to it and closes the others.

Methanal, H₂C=O: the ideal trigonal planar angle is 120°. Measured values:

• H–C=O: about 121.8° (larger)
• H–C–H: about 116.5° (smaller)

The C=O double bond pushes the C–H bonds closer together. The same effect appears in ethene (H–C–H about 117°) and in carboxylic acids and esters (see single, double and triple bonds compared).

Putting it all together: water

So why exactly 104.5° for water?

1. Four electron regions around oxygen → roughly tetrahedral arrangement (starting point 109.5°).
2. Two lone pairs, which repel more strongly than bonding pairs → angle squeezed below 109.5°.
3. Oxygen is small and very electronegative → it hybridises effectively and holds the bonding pairs close, so the angle stays well above the 90° seen in H₂S.

The balance of these effects gives 104.5°. It’s worth remembering that 104.5° isn’t a number that VSEPR theory calculates exactly — VSEPR predicts “a bit less than 109.5°”, and the precise value comes from measurement (or detailed quantum calculations).

Why bond angles matter

• Polarity: water is polar because it’s bent; a linear H–O–H would be non-polar. That polarity makes water a great solvent and gives it hydrogen bonding (see dipole moments).
• Ice: water’s near-tetrahedral angle lets molecules hydrogen-bond into an open hexagonal lattice, which is why ice is less dense than liquid water and floats.
• Biology: the exact geometry of carbon atoms (109.5°) and flat peptide bonds (about 120°) controls how proteins and DNA fold.
• Ring strain: in cyclopropane, C–C–C angles are forced to 60°, far from 109.5°, making the ring strained and reactive.

Common misconceptions

• “Water is bent because hydrogen atoms repel.” It’s bent because of the lone pairs on oxygen.
• “All four-region molecules have 109.5° angles.” Only when all four regions are identical bonding pairs.
• “VSEPR predicts exact angles.” It predicts shapes and trends; exact values come from measurement.
• “Bigger central atoms give bigger angles.” For hydrides, bigger central atoms give smaller angles.

Key takeaways

• Electron pairs arrange to minimise repulsion, giving ideal angles of 180°, 120°, 109.5°, 90°.
• Lone pairs repel more: CH₄ 109.5° → NH₃ 107° → H₂O 104.5°.
• Electronegative outer atoms close angles (NF₃ < NH₃); less electronegative, larger central atoms also close them (PH₃, H₂S near 90°).
• Multiple bonds push neighbouring bonds apart.
• Water’s 104.5° results from two lone pairs on a small, electronegative oxygen.