The value of $K_p$ for the following reaction $2H_2S_{(g)} \rightleftharpoons 2H_{2(g)} + S_{2(g)}$ is $1.2 \times 10^{-2}$ at $106.5 ^oC$. The value of $K_c$ for this reaction is

$1.1 \ mol$ of $A$ are mixed with $2.2 \ mol$ of $B$ and the mixture is kept in a $1 \ L$ vessel till the equilibrium is established in the reaction. If the molar concentration of $C$ at the equilibrium point is $0.2 \ M$,then find the value of the equilibrium constant $(K_c)$ for the reaction: $A + 2B \rightleftharpoons 2C + D$

For the reaction $A + B \rightleftharpoons C + D$,the initial concentrations of $A$ and $B$ are equal. If the equilibrium concentration of $D$ is twice the equilibrium concentration of $A$,then the equilibrium constant $K_c$ is:

For a gaseous reversible reaction,the enthalpy of reaction at constant pressure is $1.8 \ Kcal/mol$ greater than that at constant volume at $300 \ K$. The value of $\left( \frac{K_P}{K_C} \right)$ for the reaction at $T = \left( \frac{1}{0.00821} \right) \ K$ is:

For the reaction ${N_{2(g)}} + {O_{2(g)}} \rightleftharpoons 2NO_{(g)}$,the value of $K_c$ at $400 \ K$ is $4.0 \times 10^{-6}$. The value of $K_p$ for this reaction is .....

$9.2 \ g$ of $N_2O_{4(g)}$ is taken in a $1 \ L$ closed vessel and heated until the following equilibrium is attained:
${N_2}{O_{4(g)}} \rightleftharpoons 2N{O_{2(g)}}$
If $50\%$ of $N_2O_{4(g)}$ dissociates at equilibrium,what will be the equilibrium constant (in $mol \ L^{-1}$)? (Mol. wt. of $N_2O_4 = 92$)