2.5:
Chemical Equilibria: Redefining Equilibrium Constant
Recall that the solubility of a sparingly soluble salt increases with the addition of an inert salt in a phenomenon known as the salt effect.
The extent of the salt effect depends on the ionic strength of the solution.
To account for the deviation from ideality, the equilibrium constant must relate the activity—which incorporates the partial molar Gibbs energy, or chemical potential—of each species.
This equilibrium constant, called the thermodynamic or standard equilibrium constant, expresses the Gibbs energy change of the process and incorporates the effects of ionic strength.
At negligible ionic strength, the activity coefficient is approximately unity, and the concentration equilibrium constant is approximately equal to the thermodynamic equilibrium constant.
This approximation is valid for solutions of singly charged ions or non-dissociating species with ionic strengths lower than 0.01 moles per liter.
However, for solutions with ionic strengths greater than 0.01 moles per liter or multiply charged ions, activity coefficient corrections are needed to avoid significant errors.
2.5:
Chemical Equilibria: Redefining Equilibrium Constant
The effect of an inert salt on the solubility of a sparingly soluble salt is known as the salt effect. The degree of the salt effect varies with the ionic strength of the solution, which in turn depends on the activity of the species in the solution. The activity is expressed as the product of concentration and the activity coefficient of the species.
To calculate the equilibrium constants of solutions of moderately high ionic strength, one must account for the salt effect. This redefined equilibrium constant is also called the thermodynamic equilibrium constant or standard equilibrium constant, as it expresses the Gibbs energy change of the process. The thermodynamic equilibrium constant incorporates the ionic strength of the solution.
In solutions of low ionic strength (nearly an ideal solution), the activity coefficient is close to 1. Thus, the thermodynamic equilibrium constant is approximately equal to the concentration equilibrium constant.
The activity coefficient corrections are often ignored to simplify the experimental calculations of equilibrium constants. This approximation is valid for dilute solutions containing singly charged ions or non-dissociating species with ionic strengths lower than 0.01 mol/L. Activity coefficient corrections become more critical for solutions with ionic strengths greater than 0.01 mol/L or of multiply charged ions. Ignoring the activity coefficient in such cases results in significant errors in calculations.