Electrolyte solutions deviate from ideality quite a bit, as evidenced from colligative property experiments (such as freezing point), molar conductivity, and other sources. Strong electrolytes are almost completely dissociated in aqueous solution. Deviations between observed and ideally expected solution properties are NOT due to ion pairing for these strong electrolytes. Instead, the electrical interactions which extend among the ions over very long ranges give rise to large deviations from ideal solution behavior.

Lewis and Randall found an empirical relationship between the non-ideality of electrolyte solutions and the concentration of the ions. They defined the ionic strength (I) as:

where m represents molality and Z ionic charge. For an electrolyte like NaCl, I = m. However, for Na₂SO₄,

There are two take away messages here.

Next, when studying chemical systems, it is important to control this effect. Studies are undertaken at constant ionic strength. Indeed, in many studies salts are added simply to control ionic strength. But, adding a salt adds the complication of potential reacting species. So, the choice of salts is important. Sometimes salts are added to maintain constant ionic strength and are not otherwise involved in the chemical reactions. Potassium nitrate and sodium perchlorate are frequently used for this purpose. To be certain that neither the added cation nor anion are involved in chemical reactions, special salts such as tetramethyl ammonium tetraphenyl borate are used. (See earlier, net ionic reactions.)