Ions of opposite charge are naturally attracted to each other by the electrostatic force. This is described by Coulomb's law:
where  is the force of attraction,  and  are the magnitudes of the electrical charges,  is the dielectric constant of the medium and  is the distance between the ions. For ions in solution this is an approximation because the ions exert a polarizing effect on the solvent molecules that surround them, which attenuates the electric field somewhat. Nevertheless, some general conclusions can be inferred.
:Ion association will increase as:
:*the magnitude(s) of the electrical charge(s)  and  increase,
:*the magnitude of the dielectric constant  decreases,
:*the size of the ions decreases so that the distance  between cation and anion decreases.
The equilibrium constant  for ion-pair formation, like all equilibrium constants, is related to the standard free-energy change:
where  is the gas constant and  is the temperature in kelvins. Free energy is made up of an enthalpy term and an entropy term:
The coulombic energy released when ions associate contributes to the enthalpy term, . In the case of contact ion pairs, the covalent interaction energy also contributes to the enthalpy, as does the energy of displacing a solvent molecule from the solvation shell of the cation or anion. The tendency to associate is opposed by the entropy term, which results from the fact that the solution containing unassociated ions is more disordered than a solution containing associates. The entropy term is similar for electrolytes of the same type, with minor differences due to solvation effects. Therefore, it is the magnitude of the enthalpy term that mostly determines the extent of ion association for a given electrolyte type. This explains the general rules given above.