$A$ weak monoprotic acid of $0.1 \ M$ ionizes to $1\%$ in solution. What will be the $pH$ of the solution?

$A$ weak acid,$HA,$ has a $K_a$ of $1.00 \times 10^{-5}.$ If $0.100 \ mol$ of this acid is dissolved in one litre of water,the percentage of acid dissociated at equilibrium is closest to $...\%$

For two acids $A$ and $B$,the values of $pK_{a1} = 1.2$ and $pK_{a2} = 2.8$ are given. Which of the following is correct?

The $pH$ of $0.1 \ M$ acetic acid is $3$. The dissociation constant of the acid will be:

The degree of dissociation of acetic acid $(0.1 \ mol \ L^{-1})$ in water ($K_a$ of acetic acid is $10^{-5}$) is

$A$ weak base is $1.3\%$ dissociated in its aqueous solution. If $K_b$ for the weak base is $1.69 \times 10^{-5}$ at $298 \text{ K}$, find the concentration of the aqueous solution of the weak base. (in $\text{ M}$)