Dipole - Dipole Forces

Attraction between permanent dipoles

When molecules already carry a permanent dipole — like HCl, where chlorine pulls electron density towards itself — their oppositely-charged ends attract neighbouring molecules. The ends of such a dipole carry only partial charges, written with the Greek letter delta ($\delta$); these are always smaller in magnitude than a full electronic charge ($1.6\times10^{-19}\,\text{C}$).

Figure 5.2(a) shows the lopsided electron-cloud distribution inside a single HCl molecule — more density piled up on the chlorine end ($\delta^{-}$), leaving hydrogen electron-poor ($\delta^{+}$). Figure 5.2(b) then shows two such HCl molecules lining up head-to-tail so that the $\delta^{+}$ end of one sits near the $\delta^{-}$ end of the other — this alignment is the dipole-dipole interaction.

How strong, and how does it fall off?

Because only partial charges are involved, dipole-dipole attraction is:

• stronger than dispersion (London) forces, but
• weaker than the full ion-ion interaction between actual charged ions.

The attraction weakens as the molecules move apart, and — interestingly — the exact distance-dependence depends on whether the molecules can rotate freely:

• for stationary polar molecules (as in a solid), interaction energy $\propto \dfrac{1}{r^{3}}$
• for freely rotating polar molecules, interaction energy $\propto \dfrac{1}{r^{6}}$ …