The conversion of oxygen to ozone represented by the equation $3O_2 \to 2O_3$ is an endothermic reaction. The enthalpy change $\Delta H$ accompanying the reaction:

The $\Delta H_f$ for $CO_{2(g)}$,$CO_{(g)}$ and $H_2O_{(g)}$ are $-393.5$,$-110.5$ and $-241.8 \ kJ \ mol^{-1}$ respectively. The standard enthalpy change for the reaction is:
$CO_{2(g)} + H_{2(g)} \to CO_{(g)} + H_2O_{(g)}$

$AB$,$A_2$ and $B_2$ are diatomic molecules. If the bond enthalpies of $A_2$,$AB$ and $B_2$ are in the ratio $1:1:0.5$ and enthalpy of formation of $AB$ from $A_2$ and $B_2$ is $-100 \, kJ \, mol^{-1}$,what is the bond energy of $A_2$ in $kJ \, mol^{-1}$?

The standard enthalpy of formation $(\Delta_f H^o_{298})$ for methane,$CH_4$ is $-74.9 \ kJ \ mol^{-1}$. In order to calculate the average energy given out in the formation of a $C-H$ bond from this,it is necessary to know which one of the following?

Define thermochemical equations and explain the conventions used in them.

$2C + O_2 \to 2CO$; $\Delta H = -220 \ kJ$. Which of the following statements is correct for this reaction?