For a hypothetical reaction $A_{2(g)} + B_{2(g)} \rightleftharpoons 2AB_{(g)}$,at $200 \ K$,$\Delta_r G$ and $\Delta_r S$ are $20 \ kJ \ mol^{-1}$ and $-20 \ J \ K^{-1} \ mol^{-1}$ respectively. If $\Delta_r C_p = 20 \ J \ K^{-1} \ mol^{-1}$,find $\Delta_r H$ at $400 \ K$ in $kJ \ mol^{-1}$.

Which of the following is not a correct statement?

For the complete combustion of ethanol,$C_2H_5OH_{(l)} + 3O_{2(g)} \rightarrow 2CO_{2(g)} + 3H_2O_{(l)}$,the amount of heat produced as measured in a bomb calorimeter is $1364.47 \ kJ \ mol^{-1}$ at $25 \ ^oC$. Assuming ideality,the enthalpy of combustion,$\Delta_cH$,for the reaction will be: $(R = 8.314 \ J \ K^{-1} \ mol^{-1})$ .....$kJ \ mol^{-1}$

The molar heat capacity for an ideal gas at constant pressure is $20.785 \ J \ K^{-1} \ mol^{-1}$. The change in internal energy is $5000 \ J$ upon heating it from $300 \ K$ to $500 \ K$. The number of moles of the gas is [Nearest integer] (Given: $R = 8.314 \ J \ K^{-1} \ mol^{-1}$)

Find the enthalpy of neutralisation in $kJ/mol$ for the reaction between $NH_4OH$ and $HCN$ in aqueous solution,given that the enthalpies of ionisation of $NH_4OH$ and $HCN$ are $7 \ kJ/mol$ and $8 \ kJ/mol$ respectively,and the enthalpy of neutralisation of a strong acid and a strong base is $-57.3 \ kJ/mol$.

State $1 \longleftarrow$ State $2 \longleftarrow$ State $3$
$\left(\begin{array}{c} T=300 \ K \\ P=15 \ bar \\ 1 \ mole \end{array}\right) \left(\begin{array}{c} T=300 \ K \\ P=10 \ bar \\ 1 \ mole \end{array}\right) \left(\begin{array}{c} T=300 \ K \\ P=5 \ bar \\ 1 \ mole \end{array}\right)$
The above shows a cyclic process. Calculate the total work done during one complete cycle. (Assume a single step to reach the next state).