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?
$(\log 4.242=0.6275) ;(\log 0.04242=-1.372) \quad (K_a=1.8 \times 10^{-5})$

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

The $pH$ of a $0.01 \ M$ weak acid $HX$ $(K_{a}=4 \times 10^{-10})$ is found to be $5$. Now the acid solution is diluted with excess of water so that the $pH$ of the solution changes to $6$. The new concentration of the diluted weak acid is given as $x \times 10^{-4} \ M$. The value of $x$ is $...........$ (nearest integer)

The $K_a$ of a weak acid $HA$ is $1.00 \times 10^{-5}$. If $0.100 \ mol$ of this acid is dissolved in $1 \ L$ of water,what is the percentage dissociation of the acid at equilibrium?

If the degree of ionisation is $0.01$ for a decimolar solution of a weak acid $HA$,then the $pKa$ of the acid is:

The ionization constant of $0.1$ $M$ weak acid is $1.74 \times 10^{-5}$ at $298$ $K$ temperature. Calculate the $pH$ of its $0.1$ $M$ solution. (in $.88$)