The $pH$ of a monoacidic weak base is $10.9$. Calculate the percent dissociation in a $0.02 \ M$ solution. (in $\%$)

Calculate the $pH$ of a solution containing $6.0 \ g$ of acetic acid $(CH_3COOH)$ in $250 \ mL$ of water. Given: $K_a = 1.8 \times 10^{-5}$ at $298 \ K$,and atomic masses are $C = 12, H = 1, O = 16$. (in $.57$)

$A$ weak acid of dissociation constant $10^{-5}$ is being titrated with aqueous $NaOH$ solution. The $pH$ at the point of one-third neutralization of the acid will be

Degree of dissociation is $10\%$ for $10^{-3} \ M$ solution of $H_2CO_3$. Then,the $pH$ of the solution is:

Ionisation constant of $CH_3COOH$ is $1.7 \times 10^{-5}$ and concentration of $H^{+}$ ions is $3.4 \times 10^{-4} \ M$. Find the initial concentration of $CH_3COOH$ molecules.

What is the $pH$ of a $1 \ M$ $CH_3COOH$ solution? Given that the $K_a$ of acetic acid $= 1.8 \times 10^{-5}$.