At $298 \, K$,the bond energies of $C-H$,$C-C$,$C=C$,and $H-H$ bonds are $414$,$347$,$615$,and $435 \, kJ \, mol^{-1}$ respectively. What is the enthalpy change for the reaction $H_2C=CH_{2(g)} + H_{2(g)} \rightarrow H_3C-CH_{3(g)}$ at $298 \, K$?

Which of the following fuels will have the highest calorific value $(kJ/kg)$?

Enthalpy of formation of methane is $-75 \ kJ / mol$. What is the enthalpy change for formation of $24 \ g$ of methane (in $kJ$)?

The enthalpy of reaction for the reaction: $2 H_{2(g)} + O_{2(g)} \to 2 H_{2}O_{(l)}$ is $\Delta_{r}H^{\theta} = -572 \ kJ \ mol^{-1}$. What will be the standard enthalpy of formation of $H_{2}O_{(l)}$?

Calculate the standard enthalpy change for the reaction,$C_2H_5OH_{(\ell)} + 3O_{2_{(g)}} \rightarrow 2CO_{2_{(g)}} + 3H_2O_{(\ell)}$. Given: $\Delta_{f}H^{\circ}(C_2H_5OH) = -280 \ kJ \ mol^{-1}$,$\Delta_{f}H^{\circ}(CO_2) = -390 \ kJ \ mol^{-1}$,and $\Delta_{f}H^{\circ}(H_2O) = -285 \ kJ \ mol^{-1}$.

Given that,$C_{(s)} + O_{2(g)} \longrightarrow CO_{2(g)} ; \Delta H^{\circ} = -x \ kJ \ mol^{-1}$ and $2CO_{(g)} + O_{2(g)} \longrightarrow 2CO_{2(g)} ; \Delta H^{\circ} = -y \ kJ \ mol^{-1}$. The enthalpy of formation of $CO$ will be: