Question

The rapid change of $pH$ near the stoichiometric point of an acid base titration is the basis of indicator detection. $pH$ of the solution is related to ratio of the concentrations of the conjugate acid ($HIn$) and base $(In^-)$ forms of the indicator given by the expression

• A. $\log \frac{[In^-]}{[HIn]} = pK_{In} - pH$
• B. $\log \frac{[HIn]}{[In^-]} = pK_{In} - pH$
• C. $\log \frac{[HIn]}{[In^-]} = pH - pK_{In}$
• D. $\log \frac{[In^-]}{[HIn]} = pH - pK_{In}$

Answer 1

Answer: (D)

$\log \frac{[In^-]}{[HIn]} = pH - pK_{In}$

Answer 2

Acid indicators are generally weak acid. The dissociation of indicator Hln takes place as follows $Hln \rightleftharpoons H^+ + ln^-$ $\therefore K_{ln} = \frac{[H^+][ln^-]}{[Hln]}$ or $[H^+] = K_{ln} . \frac{[Hln]}{[ln^-]}$ $\because pH = - \log [H^+]$ $\therefore pH = - \log \bigg (K_{ln} . \frac{[Hln]}{[ln^-]} \bigg )$ $= - \log K_{ln} + \log \frac{[ln^-]}{[Hln]}$ $= pK_{ln} + \log \frac{[ln^-]}{[Hln]}$ or $\log \frac{[ln^-]}{[Hln]} = pH - pK_{ln}$