Chemical Equilibria: Redefining Equilibrium Constant

The salt effect describes how solubility of a sparingly soluble salt increases when an inert salt is added to the solution. This phenomenon occurs because the added salt increases the ionic strength of the solution, which affects the activity coefficients of the dissolved species and alters their effective concentrations in the equilibrium system.

The thermodynamic equilibrium constant relates the activity of each species, which incorporates both concentration and activity coefficient, accounting for deviations from ideal behavior. The concentration equilibrium constant uses only molar concentrations. At low ionic strengths, activity coefficients approach unity, making the two constants approximately equal, but at higher ionic strengths, activity coefficient corrections become necessary.

Activity coefficient corrections can be ignored for dilute solutions containing singly charged ions or non-dissociating species with ionic strengths lower than 0.01 mol/L. In these cases, the activity coefficient is approximately unity, and the concentration equilibrium constant is approximately equal to the thermodynamic equilibrium constant, simplifying experimental calculations.

Ionic strength determines the magnitude of the salt effect and the deviation from ideal solution behavior. Solutions with ionic strengths greater than 0.01 mol/L or containing multiply charged ions require activity coefficient corrections to avoid significant errors. The higher the ionic strength, the greater the deviation from ideality and the more critical these corrections become.

Activity is the product of concentration and the activity coefficient of a species. It represents the effective concentration of a species in solution, accounting for interactions between ions and deviations from ideal behavior. The activity coefficient incorporates the partial molar Gibbs energy, or chemical potential, of each species in the system.

Multiply charged ions experience stronger electrostatic interactions with other ions in solution, making their activity coefficients more sensitive to changes in ionic strength. These ions require activity coefficient corrections even at lower ionic strengths to prevent significant errors in equilibrium constant calculations and to accurately predict solubility behavior.

The thermodynamic equilibrium constant incorporates the ionic strength of the solution and relates the activities of species, which are based on their chemical potential. By accounting for activity rather than concentration alone, this constant accurately expresses the Gibbs energy change of the process and reflects the true thermodynamic driving force of the reaction.