If the pH of a buffer solution containing 0.1 | vedclass

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(A) For a basic buffer,the Henderson-Hasselbalch equation is given by: $pOH = pK_b + \log \frac{[	ext{salt}]}{[	ext{base}]}$.Given $pH = 10.5$,we ca

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(A) For a basic buffer,the Henderson-Hasselbalch equation is given by: $pOH = pK_b + \log \frac{[	ext{salt}]}{[	ext{base}]}$.Given $pH = 10.5$,we calculate $pOH$ as: $pOH = 14 - 10.5 = 3.5$.Substituting the values into the equation: $3.5 = pK_b + \log \frac{0.01}{0.1}$.$3.5 = pK_b + \log(10^{-1})$.$3.5 = pK_b - 1$,which gives $pK_b = 4.5$.Using the relation $pK_a + pK_b = 14$,we find the $pK_a$ of the conjugate acid: $pK_a = 14 - 4.5 = 9.5$.

If the $pH$ of a buffer solution containing $0.1 \ M$ of monoacidic base and $0.01 \ M$ of its salt is $10.5$,the $pK_a$ of the conjugate acid is: (in $.5$)

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