pH of 0.01 M-(NH4)2 SO4 and 0.02 M -NH4 OH buffer (p Ka. of NH4^+=9.26 ) is (a) 9.26+log2 (b) 9.

step 1 So in this question, M. O .H. X gives M plus X plus OH negative. So this is 10 .st to power minu

PH of 0.01 M-(NH4)2 SO4 and 0.02 M -NH4 OH buffer (p Ka. of...

$\mathrm{pH}$ of $0.01 \mathrm{M}-\left(\mathrm{NH}_{4}\right)_{2} \mathrm{SO}_{4}$ and $0.02 \mathrm{M}$
$-\mathrm{NH}_{4} \mathrm{OH}$ buffer $\left(\mathrm{p} K_{\mathrm{a}}\right.$ of $\mathrm{NH}_{4}^{+}=9.26$ ) is
(a) $9.26+\log 2$
(b) $9.26-\log 2$
(c) $4.74+\log 2$
(d) $9.26$

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Key Concepts

Acid-Base Buffers

Buffers are solutions that consist of a weak acid and its conjugate base (or a weak base and its conjugate acid) which help maintain a relatively constant pH when small amounts of an acid or base are added. The buffer mechanism stabilizes pH by shifting the equilibrium between the acid and its conjugate base, absorbing excess hydrogen or hydroxide ions.

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation, given by pH = pKa + log([base]/[acid]), is a key tool in calculating the pH of buffer solutions. It relates the pH of a buffer directly to the ratio of the concentrations of the conjugate base to the acid, and is derived from the equilibrium expression of the acid dissociation reaction in a solution.

Acid Dissociation Constant (pKa)

The pKa value is the negative logarithm of the acid dissociation constant (Ka) and provides a measure of the strength of an acid in solution. It signifies the pH at which half of the acid molecules are dissociated into their conjugate base form, playing a crucial role in the equilibrium calculations of buffer systems.

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