Ostwald’s Dilution Law

Based on Arrhenius' theory of electrolytic dissociation, Ostwald’s dilution law explains the equilibrium in aqueous electrolyte solutions.

Consider an electrolyte AB at a concentration ‘c’ mol per liter, which dissociates reversibly into ions.

If α is the degree of dissociation, each ion has an equilibrium concentration of cα, while the undissociated fraction (1 − α) has a concentration of c(1 − α). These terms define the equilibrium constant K, called the dissociation constant.

Ostwald’s dilution law applies only to weak electrolytes like acetic acid and NH₄OH.

Strong electrolytes like HCl and NaF are almost completely dissociated, with α close to one. Under these conditions, (1 − α) approaches zero, making the law mathematically invalid.

The failure of the law for strong electrolytes arises because, on dissolution, water’s high dielectric constant weakens electrostatic forces, leading to near-complete dissociation. As a result, no ion–molecule equilibrium exists, and Ostwald’s dilution law becomes inapplicable.