Why is water polar

Polarity describes how electric charge is distributed within a molecule. If the centers of positive and negative charge do not coincide, the molecule possesses a dipole moment and is called polar. Polar molecules interact strongly with electric fields and with other dipoles, which explains many of water’s unique properties such as high surface tension and solvent power.

Molecular shape and electronegativity

Oxygen is far more electronegative (3.44 on the Pauling scale) than hydrogen (2.20), so each O–H bond is strongly polar, with oxygen acquiring a partial negative charge (δ‑) and hydrogen a partial positive charge (δ+). However, polarity of individual bonds does not guarantee a polar molecule; the three‑dimensional arrangement of those bonds determines whether the bond dipoles cancel. In water the two O–H bonds are separated by a 104.5° angle, a geometry that prevents cancellation and leaves a net dipole pointing from the hydrogens toward the oxygen.

Key factors that make water polar:

• Large electronegativity difference between O and H
• Bent (V‑shaped) molecular geometry
• Bond angle of 104.5° preventing dipole cancellation

How to determine polarity of any molecule:

1. 1Identify all bonds and assign partial charges based on electronegativity differences.
2. 2Draw the molecular geometry using VSEPR or experimental data.
3. 3Represent each bond dipole as a vector pointing from positive to negative charge.
4. 4Add the vectors vectorially; a non‑zero resultant indicates a polar molecule.

A concrete dipole‑moment calculation illustrates the concept. Each O–H bond carries about 0.33 e of charge separation and has a bond length of 0.96 Å. The dipole moment of one bond is $\mu = q\times d = 0.33\,e \times 0.96\,\text{Å} \approx 1.58\,\text{D}$ (1 D = 3.336×10⁻³⁰ C·m). Because the two bond vectors are separated by 104.5°, the resultant dipole is $\mu_{total}=2\mu\cos(52.25^\circ ) \approx 1.85\,\text{D}$, matching the experimental value for water.

Partial charges and bond lengths in water: