Dipole-dipole forces act between the molecules possessing permanent dipole. Ends of the dipoles possess “partial charges” and these charges are shown by Greek letter delta (δ).
Partial charges are always less than the unit electronic charge (1.6×10–19 C). The polar molecules interact with neighbouring molecules. Fig 5.2 (a) shows electron cloud distribution in the dipole of hydrogen chloride and Fig. 5.2 (b) shows dipole-dipole interaction between two HCl molecules.
This interaction is stronger than the London forces but is weaker than ion-ion interaction because only partial charges are involved. The attractive force decreases with the increase of distance between the dipoles.
As in the above case here also, the interaction energy is inversely proportional to distance between polar molecules.
Dipole-dipole interaction energy between stationary polar molecules (as in solids) is proportional to 1/r 3 and that between rotating polar molecules is proportional to 1/r 6, where r is the distance between polar molecules.
Besides dipole-dipole interaction, polar molecules can interact by London forces also. Thus cumulative effect is that the total of intermolecular forces in polar molecules increase.