The standard enthalpies of formation of $\mathrm{CH}_4(\mathrm{~g}), \mathrm{CO}_2(\mathrm{~g})$ and $\mathrm{H}_2 \mathrm{O}(\mathrm{l})$ are $-74.8 \mathrm{~kJ} \mathrm{~mol}^{-1},-393.5 \mathrm{~kJ} \mathrm{~mol}^{-1}$ and $-285.8 \mathrm{~kJ} \mathrm{~mol}^{-1}$ respectively. Then the enthalpy change for the given reaction in $\mathrm{kJ} \mathrm{mol}^{-1}$ will be:

$$ 2 \mathrm{CH}_4(\mathrm{~g})+4 \mathrm{O}_2(\mathrm{~g}) \rightarrow 2 \mathrm{CO}_2(\mathrm{~g})+4 \mathrm{H}_2 \mathrm{O}(\mathrm{l}) $$

Identify the INCORRECT statement

Ozone is formed by the reaction

$$ \mathrm{O}_{2(g)}+\mathrm{O}_{(g)} \rightarrow \mathrm{O}_{3(g)}, \Delta \mathrm{H}=-107.2 \mathrm{~kJ} . $$

Given $\mathrm{O}=0$ bond energy is $498.0 \mathrm{~kJ} \mathrm{~mol}^{-1}$, the average bond energy of ozone is:

$\Delta \mathrm{H}$ and $\Delta \mathrm{S}$ for a reaction are $35.5 \mathrm{~kJ} \mathrm{~mol}^{-1}$ and $83.6 \mathrm{~J} \mathrm{~K}^{-1}$ respectively. Assuming that $\Delta \mathrm{H}$ and $\Delta \mathrm{S}$ do not vary with temperature, the reaction is spontaneous when: