At $60\,^{\circ}C$,the concentration of $[H_3O^{+}]$ in pure water is $10^{-6.7} \, \text{mol/L}$. What is the value of $K_W$ at $60\,^{\circ}C$?

Which one is a mixed salt?

What is the $[OH^-]$ in the final solution prepared by mixing $20.0 \ mL$ of $0.050 \ M \ HCl$ with $30.0 \ mL$ of $0.10 \ M \ Ba(OH)_2$?

At $25^{\circ}C$,the dissociation constants of acetic acid and $HCN$ are $1.5 \times 10^{-5}$ and $4.5 \times 10^{-10}$ respectively. What is the equilibrium constant for the following reaction? $(CN^- + CH_3COOH \rightleftharpoons HCN + CH_3COO^-)$

When $100 \ mL$ of $1.0 \ M \ HCl$ was mixed with $100 \ mL$ of $1.0 \ M \ NaOH$ in an insulated beaker at constant pressure,a temperature increase of $5.7^{\circ} C$ was measured for the beaker and its contents (Expt. $1$). Because the enthalpy of neutralization of a strong acid with a strong base is a constant $\left(-57.0 \ kJ \ mol ^{-1}\right)$,this experiment could be used to measure the calorimeter constant. In a second experiment (Expt. $2$),$100 \ mL$ of $2.0 \ M$ acetic acid $\left(K_a=2.0 \times 10^{-5}\right)$ was mixed with $100 \ mL$ of $1.0 \ M \ NaOH$ (under identical conditions to Expt. $1$) where a temperature rise of $5.6^{\circ} C$ was measured.
(Consider heat capacity of all solutions as $4.2 \ J \ g ^{-1} K ^{-1}$ and density of all solutions as $1.0 \ g \ mL ^{-1}$)
$1.$ Enthalpy of dissociation (in $kJ \ mol ^{-1}$) of acetic acid obtained from the Expt. $2$ is
$(A) \ 1.0 \ (B) \ 10.0 \ (C) \ 24.5 \ (D) \ 51.4$
$2.$ The $pH$ of the solution after Expt. $2$ is
$(A) \ 2.8 \ (B) \ 4.7 \ (C) \ 5.0 \ (D) \ 7.0$
Give the answer for question $1$ and $2.$

In which case is the change in $pH$ maximum?