The enthalpy changes for the following processes are listed below :

Cl₂(g) = 2Cl(g), 242.3 kJ mol^–1
I₂(g) = 2I(g), 151.0 kJ mol^–1
ICl(g) = I(g) + Cl(g), 211.3 kJ mol^–1
I₂(s) = I₂(g), 62.76 kJ mol^–1

Given that the standard states for iodine and chlorine are I₂(s) and Cl₂(g), the standard enthalpy of formation for ICl(g) is :

Consider an endothermic reaction, X $$\to$$ Y with the activation energies E_b and E_f for the backward and forward reactions, respectively. In general :

Consider the reaction: N₂ + 3H₂ $$\to$$ 2NH₃ carried out at constant temperature and pressure. If $$\Delta H$$ and $$\Delta U$$ are the enthalpy and internal energy changes for the reaction, which of the following expressions is true?

If the bond dissociation energies of XY, X₂ and Y₂ (all diatomic molecules) are in the ratio of 1:1:0.5 and $$\Delta H_f$$ for the formation of XY is -200 kJ mole^-1. The bond dissociation energy of X₂ will be :