When $CO_2$ dissolves in water,the following equilibrium is established:
$CO_2 + 2H_2O \rightleftharpoons H_3O^{+} + HCO_3^-$
For which the equilibrium constant is $3.8 \times 10^{-7}$ and $pH = 6.0$. The ratio of $[HCO_3^-]$ to $[CO_2]$ would be:

The first ionization constant of $H_2S$ is $9.1 \times 10^{-8}$. Calculate the concentration of $HS^{-}$ ion in its $0.1 \ M$ solution. How will this concentration be affected if the solution is $0.1 \ M$ in $HCl$ also? If the second dissociation constant of $H_2S$ is $1.2 \times 10^{-13}$,calculate the concentration of $S^{2-}$ under both conditions.

At $27^{\circ} C$,the degree of dissociation of $HA$ (weak acid) in $0.5 \ M$ of its solution is $1 \%$. The concentrations of $H_3O^{+}$,$A^{-}$,and $HA$ at equilibrium (in $mol \ L^{-1}$) are respectively

$20 \ mL$,$0.1 \ M \ HA$ $[(K_b)_{A^-} = 10^{-10}]$ is titrated against $0.2 \ M \ NaOH$. Select the correct statement [Given : $\log 2 = 0.3, \log 3 = 0.48$ ]

If $pK_{a} = -\log K_{a} = 4$ for a weak acid $HX$ and $K_{a} = C\alpha^{2}$,then the Van't Hoff factor when $C = 0.01 \ M$ is

What is the $pH$ of a solution containing $10 \, mL$ of $0.1 \, M \, NaOH$ and $10 \, mL$ of $0.05 \, M \, H_2SO_4$?