For the formation of carbon monoxide $(CO)$ from its elements at $298 \, K$,the value of $(\Delta H - \Delta U)$ in $J \, mol^{-1}$ will be: $(R = 8.314 \, J K^{-1} mol^{-1})$

An ideal gas,$\overline{C}_{V} = \frac{5}{2} R$,is expanded adiabatically against a constant pressure of $1 \ atm$ until it doubles in volume. If the initial temperature and pressure are $298 \ K$ and $5 \ atm$,respectively,then the final temperature is . . . . . . $K$ (nearest integer). [$\overline{C}_{V}$ is the molar heat capacity at constant volume]

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)

Which of the following expressions represents the first law of thermodynamics?

What is the value of $\Delta H - \Delta U$ for the formation of $2$ moles of ammonia from $H_{2(g)}$ and $N_{2(g)}$?

When a system absorbs $8 \ kJ$ of heat and does $2.2 \ kJ$ of work on the surroundings,calculate the internal energy change. (in $kJ$)