Ionic Equilibrium in Solution

A $0.02 \mathrm{M}$ solution of pyridinium hydrochloride has $\mathrm{pH}=3.44$ calculate ionization constant of pyridine. (antilog $0.18=1.5$)

It has been found that the $\mathrm{pH}$ of a $0.01 \mathrm{M}$ solution of an organic acid is $4.15$. Calculate the concentration of the anion, the ionization constant of the acid and its ${\mathrm{pK}}_{\mathrm{a}}$. (antilog $0.85=7.08$; antilog $0.7=5$)

If a weak base has the dissociation constant, ${\mathrm{K}}_{\mathrm{b}}$, then the value of the dissociation constant, ${\mathrm{K}}_{\mathrm{a}}$, of its conjugate acid is given by:

The molarity of ${\mathrm{NH}}_3$ solution of $\mathrm{pH} 12.0 \mathrm{at} 25° \mathrm{C}$ is: $\left({\mathrm{K}}_{\mathrm{b}}\right.$ of ${\mathrm{NH}}_3$ at $25° \mathrm{C}$ is $\left.1.8\times {10}^{-5}\right)$

The precipitate $\mathrm{AgCl}$ dissolves in:

Sodium hydroxide cannot be used as a primary standard for acid base titration because:

Which of the following is not a strong electrolyte?

$10 \mathrm{mL}$ of $\frac{\mathrm{M}}{5}{\mathrm{CH}}_3\mathrm{COONa}$ solution is titrated with $\frac{\mathrm{M}}{5}\mathrm{HCl}$ solution. The $\mathrm{pH}$ value at equivalence point is: $\left({\mathrm{pK}}_{\mathrm{a}}\left({\mathrm{CH}}_3\mathrm{COOH}\right)=4.76\right)$

Ratio of $\left[{\mathrm{HA}}^{2-}\right]$ in $1 \mathrm{L}$ of $0.1 \mathrm{M} {\mathrm{H}}_3\mathrm{A}$ solution $\left({\mathrm{K}}_{{\mathrm{a}}_1}={10}^{-5}; {\mathrm{K}}_{{\mathrm{a}}_2}={10}^{-8} \& {\mathrm{K}}_{{\mathrm{a}}_3}={10}^{-11}\right) \&$ upon addition of $0.1 \mathrm{mole}$ $\mathrm{HCl}$ to it will be :

In a solution obtained by mixing $100 \mathrm{mL}$ of  $0.25 \mathrm{M}$  Triethylamine $\left({\mathrm{K}}_{\mathrm{b}}=6.4\times {10}^{-5}\right) \& 400 \mathrm{mL}$ of  $\frac{\mathrm{M}}{18}{\mathrm{NH}}_4$ $\mathrm{OH} \left({\mathrm{K}}_{\mathrm{b}}=1.8\times {10}^{-5}\right)$ $:$