Calculate the enthalpy change for the reaction $H_{2(g)} + Br_{2(g)} \rightarrow 2HBr_{(g)}$ in $kJ$. The bond energies of $H-H$,$Br-Br$,and $H-Br$ are $435$,$192$,and $364 \, kJ \, mol^{-1}$ respectively.

The heat change for the reaction $H_2 + \frac{1}{2} O_2 \to H_2 O$ is called

Given $C + O_2 \rightarrow CO_2 + 94.2 \ kcal$,$\Delta H = -94.2 \ kcal$; $H_2 + 1/2 O_2 \rightarrow H_2O + 68.3 \ kcal$,$\Delta H = -68.3 \ kcal$ and $CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O + 210.8 \ kcal$,$\Delta H = -210.8 \ kcal$. Calculate the heat of formation of methane $(CH_4)$ in $kcal$.

$A$,$B$,$C$ and $D$ are some compounds. The enthalpy of formation of $A_{(g)}$,$B_{(g)}$,$C_{(g)}$ and $D_{(g)}$ is $9.7, -110, 81$ and $-393 \ kJ \ mol^{-1}$ respectively. What is $\Delta_r H$ (in $kJ \ mol^{-1}$) for the given reaction?
$A_{(g)} + 3B_{(g)} \longrightarrow C_{(g)} + 3D_{(g)}$

Which of the following has standard enthalpy of formation equal to zero?

The heat of formation of methane $C_{(s)} + 2H_{2(g)} \to CH_{4(g)}$ at constant pressure is $-18500 \ cal$ at $25 \ ^oC$. The heat of reaction at constant volume would be (in $cal$)