Consider the following thermodynamic data (all values at 298 K):
$$
\begin{aligned}
\Delta H_f^\circ(\text{NO}) &= +90.3\ \text{kJ mol}^{-1}, & \Delta S_f^\circ(\text{NO}) &= +210\ \text{J mol}^{-1}\text{K}^{-1},\
\Delta H_f^\circ(\text{NO}_2) &= +33.2\ \text{kJ mol}^{-1}, & \Delta S_f^\circ(\text{NO}_2) &= +150\ \text{J mol}^{-1}\text{K}^{-1},\
\Delta H_f^\circ(\text{N}_2\text{O}) &= +120\ \text{kJ mol}^{-1}, & \Delta S_f^\circ(\text{N}_2\text{O}) &= +250\ \text{J mol}^{-1}\text{K}^{-1},\
\Delta S_f^\circ(\text{O}_2) &= +205\ \text{J mol}^{-1}\text{K}^{-1}.
\end{aligned}
$$
For the reaction
$$
\text{N}_2\text{O}(g) + \tfrac{1}{2}\,\text{O}_2(g) \rightarrow 2\,\text{NO}_2(g)
$$
calculate:
(i) the standard enthalpy change $\Delta H^\circ$,
(ii) the standard entropy change $\Delta S^\circ$,
(iii) the standard Gibbs energy change $\Delta G^\circ$ at 298 K,
(iv) the equilibrium constant $K_p$ at 298 K, and
(v) the temperature at which $\Delta G^\circ$ would become zero (assume $\Delta H^\circ$ and $\Delta S^\circ$ are temperature‑independent).

Question

Accepted Answer

(i) $\Delta H^\circ = -53.6\ \text{kJ mol}^{-1}$\n(ii) $\Delta S^\circ = -52.5\ \text{J mol}^{-1}\text{K}^{-1}$\n(iii) $\Delta G^\circ_{298\text{K}} \approx -38.0\ \text{kJ mol}^{-1}$\n(iv) $K_p \approx 4.6\times10^{6}$\n(v) $T \approx 1.0\times10^{3}\ \text{K}$

Gce Model Questions Vol1 2021 Chemistry Question 31

Answer Options

About This Question

Master the Exam!