The enthalpy of atomization of $PH_3(g)$ is $228 \, kcal \, mol^{-1}$ and that of $P_2H_4(g)$ is $355 \, kcal \, mol^{-1}$. The $P-P$ bond energy (in $kcal \, mol^{-1}$) is:

In the reaction for the transition of carbon in the diamond form to carbon in the graphite form,$\Delta H = -453.5 \ cal$. This points out that

$A_2 + B_2 \rightarrow 2AB$; $\Delta H_{r}^0 = -400\,kJ\,mol^{-1}$. $AB$,$A_2$ and $B_2$ are diatomic molecules. If the bond enthalpies of $A_2$,$B_2$ and $AB$ are in the ratio $1:0.5:1$,then the bond enthalpy of $A_2$ is $......\,kJ\,mol^{-1}$ (Nearest integer).

The average $C-H$ bond energy is $416 \ kJ \ mol^{-1}$. Which of the following equations correctly represents the bond dissociation of $CH_4$?

Which of these species has a standard enthalpy of formation equal to zero?

Given that $\Delta H_f(H) = 218 \ kJ/mol$, express the $H-H$ bond energy in $kcal/mol$.