EXERCISES

Enthalpies of formation of $\mathrm{CO}\left(\mathrm{g}\right), {\mathrm{CO}}_2\left(\mathrm{g}\right), {\mathrm{N}}_2\mathrm{O}\left(\mathrm{g}\right)$ and ${\mathrm{N}}_2{\mathrm{O}}_4\left(\mathrm{g}\right)$ are $–110, – 393, 81$ and $9.7 \mathrm{kJ} {\mathrm{mol}}^{–1}$ respectively. Find the value of ${∆}_{\mathrm{r}}\mathrm{H}$ for the reaction: 

${\mathrm{N}}_2{\mathrm{O}}_{4\left(\mathrm{g}\right)} + 3{\mathrm{CO}}_{\left(\mathrm{g}\right)} \to {\mathrm{N}}_2{\mathrm{O}}_{\left(\mathrm{g}\right)} + 3{\mathrm{CO}}_{2\left(\mathrm{g}\right)}$

Given

${\mathrm{N}}_2\left(\mathrm{g}\right)+3{\mathrm{H}}_2\left(\mathrm{g}\right)\to 2{\mathrm{NH}}_3\left(\mathrm{g}\right) ; {∆}_{\mathrm{r}}\mathrm{H}°=-92.4 \mathrm{kJ} {\mathrm{mol}}^{-1}$ 

What is the standard enthalpy of formation of ${\mathrm{NH}}_3$ gas?

Calculate the standard enthalpy of formation of ${\mathrm{CH}}_3\mathrm{OH}\left(\mathrm{l}\right)$ from the following data:

${\mathrm{CH}}_3\mathrm{OH} \left(\mathrm{l}\right) +\frac{3}{2}{\mathrm{O}}_2\left(\mathrm{g}\right)\to {\mathrm{CO}}_2\left(\mathrm{g}\right) + 2{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right) ; {∆}_{\mathrm{r}}\mathrm{H}° = –726 \mathrm{kJ} {\mathrm{mol}}^{–1}$

$\mathrm{C}\left(\mathrm{graphite}\right)+ {\mathrm{O}}_2\left(\mathrm{g}\right) \to {\mathrm{CO}}_2\left(\mathrm{g}\right) ; {∆}_{\mathrm{c}}\mathrm{H}° = –393 \mathrm{kJ} {\mathrm{mol}}^{–1}$

${\mathrm{H}}_2\left(\mathrm{g}\right)+\frac{1}{2}{\mathrm{O}}_2\left(\mathrm{g}\right)\to {\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right) ; {∆}_{\mathrm{f}}\mathrm{H}°= –286 \mathrm{kJ} {\mathrm{mol}}^{–1}.$

Calculate the enthalpy change for the process
${\mathrm{CCl}}_4\left(\mathrm{g}\right)\to \mathrm{C}\left(\mathrm{g}\right) + 4 \mathrm{Cl}\left(\mathrm{g}\right)$
and calculate bond enthalpy of $\mathrm{C} – \mathrm{Cl}$ in ${\mathrm{CCl}}_4\left(\mathrm{g}\right)$.
${∆}_{\mathrm{vap}}\mathrm{H}°\left({\mathrm{CCl}}_4\right) = 30.5 \mathrm{kJ} {\mathrm{mol}}^{–1}.$
${∆}_{\mathrm{f}}\mathrm{H}°\left({\mathrm{CCl}}_4\right) = –135.5 \mathrm{kJ} {\mathrm{mol}}^{–1}.$
${∆}_{\mathrm{a}}\mathrm{H}° \left(\mathrm{C}\right) = 715.0 \mathrm{kJ} {\mathrm{mol}}^{–1}$ , where ${∆}_{\mathrm{a}}\mathrm{H}°$is enthalpy of atomisation
${∆}_{\mathrm{a}}\mathrm{H}°\left({\mathrm{Cl}}_2\right) = 242 \mathrm{kJ} {\mathrm{mol}}^{–1}$

For the reaction at $298 \mathrm{K}$,

$2\mathrm{A} + \mathrm{B} \to \mathrm{C}$ 

$∆\mathrm{H}=400 \mathrm{kJ} {\mathrm{mol}}^{–1 }$and $∆\mathrm{S}=0.2 \mathrm{kJ} {\mathrm{K}}^{–1} {\mathrm{mol}}^{–1}$ 

At what temperature will the reaction become spontaneous considering $∆\mathrm{H}$ and $∆\mathrm{S}$ to be constant over the temperature range.

For the reaction

$2 \mathrm{A}\left(\mathrm{g}\right) + \mathrm{B}\left(\mathrm{g}\right) \to 2\mathrm{D}\left(\mathrm{g}\right)$ 

Calculate $∆\mathrm{G}°$ for the reaction, and predict whether the reaction may occur spontaneously.

The equilibrium constant for a reaction is $10$. What will be the value of $∆{\mathrm{G}}^{⊝}$? $\mathrm{R}=8.134{\mathrm{JK}}^{-1} {\mathrm{mol}}^{-1}, \mathrm{T}=300\mathrm{K}$.

Calculate the entropy change in surroundings when $1.00 \mathrm{mol}$ of ${\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)$ is formed under standard conditions. ${∆}_{\mathrm{f}}\mathrm{H}°=-286 \mathrm{kJ} {\mathrm{mol}}^{-1}$.