In an adiabatic expansion of an ideal gas,which of the following relations holds true?

Which one of the following equations does not correctly represent the first law of thermodynamics for the given processes involving an ideal gas? (Assume non-expansion work is zero)

Assertion : For a reaction $2NH_{3(g)} \to N_{2(g)} + 3H_{2(g)}$ ; $\Delta H > \Delta E$.
Reason : Enthalpy change is always greater than internal energy change.

Calculate the change in internal energy of the system if $20 \ kJ$ work is done on the system and it releases $10 \ kJ$ heat in a particular reaction. (in $kJ$)

Explain the relationship between the change in heat at constant pressure and constant volume.

For the gas phase reaction $A_{2(g)} + B_{2(g)} \rightarrow 2AB_{(g)}$ taking place at constant pressure and temperature where all gases are behaving ideally,which of the following must be correct?