2.8:
Acid–Base Equilibria: Activity-Based Definition of pH
The definition of the pH of a solution as the negative logarithm of the hydrogen ion concentration is valid only for an ideal solution.
In practice, the pH measurement considers the negative logarithm of the hydrogen ion activity rather than concentration.
So, the pH can be more accurately redefined as the negative logarithm of the product of the hydrogen ion concentration and its activity coefficient.
In the case of pure water at 25 degrees C, the extremely low ionic strength suggests that the activity coefficients of the ions are close to one.
However, adding a salt like potassium chloride to the water increases the ionic strength of the solution, thus decreasing the activity coefficients.
It follows here that the hydrogen ion activity is slightly increased, corresponding to a decrease in the pH of the solution.
Notably, the addition of potassium chloride causes a negligible change in pH but a significant increase in the hydrogen ion concentration.
2.8:
Acid–Base Equilibria: Activity-Based Definition of pH
For an ideal solution, the pH is defined as the negative logarithm of the hydrogen ion concentration. For a non-ideal solution, an accurate measurement of the pH must consider the negative logarithm of the hydrogen ion activity rather than concentration. In such a solution, the pH can be more accurately defined as the negative logarithm of a product of the hydrogen ion concentration and its activity coefficient.
In solutions of very low ionic strength—for example, pure water—the activity coefficient of the hydrogen ion is close to one when the ionic strength of the solution increases due to the addition of an electrolyte that does not donate or accept a proton. This results in a slight decrease in the pH of the solution. In other words, the addition of an electrolyte increases the hydrogen ion activity, or the effective hydrogen ion concentration in the solution, which decreases the pH of the solution.