Calculate the value of the dissociation constant $(K_b)$ of a weak monoacidic base if it dissociates to $2 \%$ in a $0.1 \ M$ solution.

The $pH$ of $0.05 \ M$ acetic acid is $(K_a = 2 \times 10^{-5})$.

$K_b$ for $NH_4OH$ is $1.8 \times 10^{-5}$. The $[OH^{-}]$ of $0.1 \ M \ NH_4OH$ is

The first and second ionization constants of $H_2X$ are $2.5 \times 10^{-8}$ and $1.0 \times 10^{-13}$ respectively. The concentration of $X^{2-}$ in $0.1 \ M$ $H_2X$ solution is . . . . . . $\times 10^{-15} \ M$. The value of $Y$ is:

The dissociation constant of a monoacidic weak base $MOH$ is $1.8 \times 10^{-5}$. What is the concentration of $OH^{-}$ ions in its $0.1 \, M$ solution?

The $pH$ of a $1.0 \ M$ monobasic acid $HX$ is $2$. The van't Hoff factor for the aqueous solution of the acid will be ........