The $\Delta H_f^o$ values for $ICl_{(g)}$,$Cl_{(g)}$,and $I_{(g)}$ are $17.57$,$121.34$,and $106.96 \, J \, mol^{-1}$ respectively. What is the bond dissociation energy of the $I-Cl$ bond in $J \, mol^{-1}$?

Given that:
$C_{(s)} + O_{2(g)} \to CO_{2(g)}, \Delta H = -394 \ kJ$
$2H_{2(g)} + O_{2(g)} \to 2H_2O_{(l)}, \Delta H = -568 \ kJ$
$CH_{4(g)} + 2O_{2(g)} \to CO_{2(g)} + 2H_2O_{(l)}, \Delta H = -892 \ kJ$
Calculate the heat of formation of $CH_{4(g)}$ in $kJ$.

If the enthalpies of atomization of methane and ethane are respectively $320 \ cal$ and $360 \ cal$,then the bond energy of the $C-C$ bond is......$cal$.

For the reaction $2H_{2(g)} + O_{2(g)} \to 2H_2O_{(l)}$,the enthalpy change is $\Delta H = -571 \ kJ$. If the $H-H$ bond energy is $435 \ kJ \ mol^{-1}$ and the $O=O$ bond energy is $498 \ kJ \ mol^{-1}$,calculate the average bond energy of the $O-H$ bond in $kJ \ mol^{-1}$.

The heat of combustion of ethanol into carbon dioxide and water is $-327 \ kcal$ at constant pressure. The heat evolved (in $cal$) at constant volume and $27^{\circ} C$ (assuming all gases behave ideally) is $\left( R = 2 \ cal \ mol^{-1} \ K^{-1} \right)$

The total enthalpies of reactants and products are $H_R$ and $H_P$ respectively. For an exothermic reaction,which of the following is true?