Polar bonds and dipole moment

When two identical atoms share a pair (as in H₂ or Cl₂), they share it equally. When two different atoms share, the more electronegative one pulls the pair closer, getting a small negative charge (δ–), and leaves the other slightly positive (δ+): H^δ+–F^δ–. Such a bond is polar. The bigger the difference in electronegativity, the more ionic the bond.

The polarity is measured by the dipole moment: µ = charge × distance between the charges. Its unit is the debye, D (1 D = 3.33564 × 10⁻³⁰ C m). Dipole moment has a direction as well as a size (it is a vector). Chemists draw it as an arrow (with a small cross at the tail) pointing from the positive to the negative end.

A molecule's total dipole moment is the vector sum of its bond dipoles, added like arrows, so it depends on the molecule's shape:

• In a symmetrical molecule the bond dipoles cancel: CO₂ (linear), BF₃ (trigonal planar) and CCl₄ (tetrahedral) all have zero dipole moment, although each has polar bonds.
• In a bent molecule such as water, they do not cancel: H₂O has a dipole moment of 1.85 D.

Lone pairs matter too. NH₃ (1.47 D) is much more polar than NF₃ (0.23 D), although both are pyramidal. In NH₃ the lone pair's dipole points the same way as the N–H bond dipoles and adds to them; in NF₃ it points against the N–F bond dipoles and partly cancels them.

Ionic bonds also have some covalent character. Fajans' rules: the smaller the cation and the higher its charge, the more it distorts (polarises) the electron cloud of the anion; and the larger the anion, the more easily it is distorted. More distortion means more covalent character.