The energy absorbed by each molecule $(A_2)$ of a substance is $4.4 \times 10^{-19} \ J$ and bond energy per molecule is $4.0 \times 10^{-19} \ J.$ The kinetic energy of the molecule per atom will be $...... \times 10^{-20} \ J$

When $10 \ g$ of methane is completely burnt in oxygen,the heat evolved is $560 \ kJ$. What is the heat of combustion (in $kJ \ mol^{-1}$) of methane?

Which of the following equations correctly represents the standard heat of formation $(\Delta H_f^o)$ of methane?

Calculate the standard enthalpy change of the reaction: $C_2H_{2(g)} + \frac{5}{2}O_{2(g)} \rightarrow 2CO_{2(g)} + H_2O_{(\ell)}$ given the following standard enthalpies of formation:
$\Delta_fH^{\circ}(CO_2) = -393 \ kJ \ mol^{-1}$
$\Delta_fH^{\circ}(H_2O) = -286 \ kJ \ mol^{-1}$
$\Delta_fH^{\circ}(C_2H_2) = 227 \ kJ \ mol^{-1}$

If standard enthalpy of formation $(\Delta_{f} H^{\circ})$ of $CO_2, H_2 O$ and $CH_4$ are $-393, -286$ and $-74.0 \ kJ \ mol^{-1}$ respectively,the standard enthalpy of combustion of methane in $kJ \ mol^{-1}$ is

The enthalpy of the reaction,$H_{2(g)} + \frac{1}{2}O_{2(g)} \to H_2O_{(g)}$ is $\Delta H_1$ and that of $H_{2(g)} + \frac{1}{2}O_{2(g)} \to H_2O_{(l)}$ is $\Delta H_2$. Then: