Question

What is the correct volume of equilibrium constant for the following reaction at 400 K if the values of $\Delta H^{\circ}$ is 77.5 kJ $mol^{-1}$ and ${\Delta S^{\circ} = 135 \; JK^{-1} \; mol^{-1}}$ ${2NOCl{(g)} <=> 2NO{(g)} + Cl_2{(g)}}$

• A. $8.545 \times 10^{-4}$
• B. $8.545 \times 10^{-2}$
• C. $8.314$
• D. $135$

Answer 1

Answer: (A)

$8.545 \times 10^{-4}$

Answer 2

Given reaction,
${2NOCl(g) <=> 2NO(g) +Cl2(g)}$
Enthalpy $(?H^?? = 77.5kJ\, mol^{-1} = 77500\, J \,mol^{-1}$
Entropy $(?S^?? = 135\,J\,K^{-1}\,mol^{-1}$
Temperature $(T) = 400\, K$
We know that, Gibbs free energy ;
$?G^??= ?H^??- T?S^??$?G^??= 77500 ,J ,mol^{-1} - 400 \times 135$$?G = 23500J ,mol^{-1}$Further,$?G^??= - 2.303,RT, log, K$$23500 = - 2.303 ??8.314 ??400 ,log,K$$log,K=-\frac{23500}{2.303\times8.31\times400}$$log,K = - 3.06$$K =$Antilog$(- 3.06)$$K= 8.545 \times 10^{-4}$