Question

Buffers based on di- or triprotic acids may have multiple pH regions over which they are stable. That is, they exhibit some stability as the pH of solution is equivalent to each of their pKas (pKa₁, pKa₂, pKa₃). If you have a di- or triprotic buffer, state below whether you see evidence of this in the form of your pH curve. If you do not see evidence of this, explain why this should be the case.

Answer 1

Ans. Using Henderson-Hasselbalch equation, pH = pKa + log [A^-] / [AH]

When, [A^-] = [AH], the equation becomes –

            pH = pKa + log 1 = pKa + 0

Hence, pH = pKa

That is, when [A^-] = [AH], the pH = pKa.

# The buffering capacity is maximum when pH = pKa.

At pH = pKa, [A^-] = [AH].

So, at pH = pKa, both the conjugate acid and its protonated form is in equi-molar concentrations. In this region, the buffer requires relatively greater amount of base to neutralize all AH. Similarly, in this region, the buffer would also require relatively greater amount of acid to convert all A^- into AH.

So, addition of large amount of acid or base at pH = pKa results little change in the slope of titration curve. These horizontal regions represent the region of pKa of the weak acid.

# Therefore, number of horizontal regions with little slope in the titration curve would be equal to the number of pKa values associated with the acid in the buffer.

An example of the diprotic weak acid is shown in the figure with explanation.

#2. If very strong titrant (acid or base) is added to the buffer in flask, the horizontal regions may not be clearly visible. It may give a false perception of the acid being monoprotic.

So, it is advised to use relatively dilute titrant. The more diluted is the titrant (solution taken in burette), the more prominent is the buffering region and so are the pKa regions.

Therefore, if you don’t see the horizontal regions depicting the pKa values, use more dilute titrant for titration.