The standard enthalpies of formation of $CO$ and $CO_2$ are $-110 \, kJ \, mol^{-1}$ and $-394 \, kJ \, mol^{-1}$ respectively. What will be the heat of combustion of $1 \, mol$ of graphite in $kJ$?

The values of bond enthalpy of $H_2$,$Cl_2$ and $HCl$ are $104$,$58$ and $103 \ kcal/mol$ respectively. Find the enthalpy of formation of $HCl_{(g)}$. Reaction: $\frac{1}{2}H_{2(g)} + \frac{1}{2}Cl_{2(g)} \to HCl_{(g)}$

Calculate the standard enthalpy of combustion of carbon monoxide $(CO)$ given that $\Delta_f H^{\circ}(CO) = -110 \text{ kJ mol}^{-1}$ and $\Delta_f H^{\circ}(CO_2) = -393 \text{ kJ mol}^{-1}$.

Use the data from the table to estimate the enthalpy of formation of $CH_3CHO$.

Bond	Bond Enthalpy $(kJ \ mol^{-1})$	Enthalpy of formation $(kJ \ mol^{-1})$
$C-H$	$400$	$C(g): 700$
$C-C$	$350$	$H(g): 200$
$C=O$	$700$	$O(g): 250$

The bond energy of $Cl-Cl$ in $Cl_2$ is $242 \, kJ \, mol^{-1}$. What is the wavelength (in $nm$) of light required to break a single $Cl-Cl$ bond?

The bond dissociation energies of $XY$,$X_2$,and $Y_2$ (all diatomic molecules) are in the ratio $1 : 1 : 0.5$. If the enthalpy of formation of $XY$ is $\Delta_fH = -200 \ kJ \ mol^{-1}$,find the bond dissociation energy of $X_2$ in $kJ \ mol^{-1}$.