Polarity of bonds and dipole moment

A bond between two identical atoms (H₂, O₂) is non-polar; when the atoms differ in electronegativity, the shared pair shifts towards the more electronegative atom and the bond becomes polar (as in HF).

Dipole moment μ = charge × distance between the charge centres; it is a vector, drawn from the positive to the negative end in the convention chemists use (crossed arrow).

The unit is the Debye: 1 D = 3.33564 × 10⁻³⁰ C m.

A molecule's net dipole moment is the vector sum of its bond dipoles, so symmetry can cancel it: BeF₂, BF₃, CO₂, CH₄ and CCl₄ have zero dipole moment.

Water is bent, so its two O–H bond dipoles add to a net value (about 1.85 D).

NH₃ has a larger dipole moment than NF₃: in NH₃ the lone-pair moment adds to the N–H bond moments, while in NF₃ it opposes the N–F bond moments.

Ionic bonds carry some covalent character too, and Fajans' rules say when it is larger: a small cation paired with a large anion, and a cation with a higher charge.

Of two cations with the same size and charge, a transition-metal type cation with an (n−1)dⁿ ns⁰ configuration (for example Cu⁺, 3d¹⁰) polarises the anion more than one with a noble gas ns² np⁶ configuration.