The $pH$ of an aqueous solution of $H_2CO_3$ is $3.3$. If ${K_{a_1}} = {10^{-3}}$ and ${K_{a_2}} = {10^{-13}}$,then the concentration of $[HCO_3^-]$ is:

If $1 \ mM$ solution of ethylamine produces $pH = 9$,then the ionization constant $(K_b)$ of ethylamine is $10^{-x}$. The value of $x$ is . . . . . . (nearest integer). [The degree of ionization of ethylamine can be neglected with respect to unity.]

What is the value of $pH$ of a $NaOH$ solution that dissociates $2 \%$ in its $0.01 \ M$ solution?

What is the $pH$ of $0.001 \,M$ aniline solution? The ionization constant of aniline can be taken from the table. Calculate the degree of ionization of aniline in the solution. Also,calculate the ionization constant of the conjugate acid of aniline.

Base	$K_{b}$
Dimethylamine,$(CH_{3})_{2}NH$	$5.4 \times 10^{-4}$
Triethylamine,$(C_{2}H_{5})_{3}N$	$6.45 \times 10^{-5}$
Ammonia,$NH_{3}$	$1.77 \times 10^{-5}$
Quinine	$1.10 \times 10^{-6}$
Pyridine,$C_{5}H_{5}N$	$1.77 \times 10^{-9}$
Aniline,$C_{6}H_{5}NH_{2}$	$4.27 \times 10^{-10}$
Urea,$CO(NH_{2})_{2}$	$1.3 \times 10^{-14}$

At $25^{\circ} C$,the percentage of ionization of $x \ M$ acetic acid is $4.242$. What is the pH of the acetic acid solution?
$(\log 4.242=0.6275) ;(\log 0.04242=-1.372) \quad (K_a=1.8 \times 10^{-5})$

The $pH$ of $0.1 \ M$ $HCN$ solution is $5.2$. Calculate the dissociation constant $K_a$ of this solution.