Effect of dilution on the pH of a buffer

The effects of dilution on the pH of a strong acid (unbuffered solution) and a weak acidic buffer are shown in the diagram below.

It is evident from the diagram that the pH of the buffer solution is resistant to dilution. The pH of the systems are expressed by the following formulae:


Approximate pH formula

More accurate pH formula


—————       pH=-log[H^+]     (link)        ————–
CH3COOH/CH3COO buffer pH=pK_a+log\frac{[A^-]_s}{[HA]_r}     (link) 10^{-pH}+[A^-]=\frac{K_w}{10^{-pH}}+\frac{K_a[HA]}{10^{-pH}+K_a}     (link)

With reference to the above formulae, [H+] of a strong acid like HCl, decreases upon dilution and hence pH increases. According to the Henderson-Hasselbalch equation, the ratio of the concentration of the salt to the concentration of the acid is unchanged upon dilution and we would expect the buffer pH curve in the above diagram to be a horizontal line. However, the curve shows that the pH of a buffer solution appears relatively constant up to a thousand-fold dilution and subsequently increases at lower buffer concentrations. This is because the data points of the buffer curve in the diagram is calculated using the more accurate pH formula (i.e. the complete pH titration curve formula) instead of the Henderson-Hasselbalch equation, which only provides an approximation of the buffer’s pH. The Henderson-Hasselbalch equation is relatively accurate at high buffer concentrations or when Ka < 10-3, but unreliable at low buffer concentrations or when Ka > 10-3, where the assumptions used to derive it are no longer valid.



Explain the resistance in the change in pH of the CH3COOH/CH3COO buffer upon dilution using Le Chatelier’s principle.


CH_3COOH(aq)+H_2O(l)\rightleftharpoons CH_3COO^-(aq)+H_3O^+(aq)

The above equilibrium shifts to the right to counteract the increase in [H2O], thereby increasing [H3O+]. Since the volume of the solution also increases, the net effect is that the pH of the buffer remains relatively unchanged.



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