Calculate the value of the dissociation constant $(K_a)$ of a weak acid,which dissociates to $0.01 \%$ in its $0.1 \ M$ solution.

$pH$ of $0.05 \ M$ aqueous solution of diethyl amine is $10$. The value of equilibrium constant $(K_b)$ is:

At $T$ $(K)$,if the ionisation constant of ammonia in solution is $2.5 \times 10^{-5}$,the $pH$ of $0.01 \ M$ ammonia solution and the ionisation constant of its conjugate acid respectively at that temperature are $(\log 2 = 0.30)$

$A$ weak base $MOH$ of $0.1 \ N$ concentration shows a $pH$ value of $9$. What is the percentage degree of ionization of the base ? $.......\%$

At what concentration of $0.2 \ M$ $CH_3COOH$ will its degree of dissociation become double (in $M$)? (Given: $K_a = 1.8 \times 10^{-5}$ for $CH_3COOH$)

Nicotinic acid $(K_a = 10^{-5})$ is represented by the formula $HNiC$. Calculate the percentage of dissociation in a solution containing $0.1 \ mol$ of nicotinic acid in $2 \ L$ of solution.