$A$ weak monoacidic base dissociates to $1.5 \%$ in $0.001 \ M$ solution at $298 \ K$. Calculate the dissociation constant of the weak base.

Dimethyl amine $(CH_3)_2NH$ is a weak base and its ionization constant is $5.4 \times 10^{-5}$. Calculate $[OH^{-}]$,$[H_3O^{+}]$,$pOH$ and $pH$ of its $0.2 \ M$ solution at equilibrium.

The $K_a$ of monobasic acid $A, B$ and $C$ are $10^{-6}, 10^{-8}$ and $10^{-10}$ respectively. The concentrations of $A, B$ and $C$ are respectively $0.1 \ M, 0.01 \ M$ and $0.001 \ M$. Which of the following is correct for $pOH$ of $A, B$ and $C$?

$A$ weak monobasic acid is $0.04 \%$ dissociated in $0.25 \text{ M}$ solution. What is the $pH$ of the solution?

$A$ solution of weak acid $HA$ containing $0.01 \ mol/L$ of acid has $pH = 4$. The percentage degree of ionisation of the acid and the ionisation constant of acid are respectively.

What is the $pH$ of a weak dibasic acid,that is $2 \%$ dissociated in its $M/100$ solution at $298 \ K$?