$1.1 \ mol$ of $A$ and $2.2 \ mol$ of $B$ are mixed in a $1 \ L$ flask until equilibrium is reached. At equilibrium,$0.2 \ mol$ of $C$ is formed. If the equilibrium reaction is $A + 2B \rightleftharpoons 2C + D$,the value of the equilibrium constant $K_C$ is:

The gas phase reaction $2A_{(g)} \rightleftharpoons A_{2(g)}$ at $400 \ K$ has $\Delta G^{\circ} = +25.2 \ kJ \ mol^{-1}$. The equilibrium constant $K_{C}$ for this reaction is $...... \times 10^{-2}$. (Round off to the nearest integer) $[$Use: $R = 8.3 \ J \ mol^{-1} \ K^{-1}$,$\ln 10 = 2.3$,$\log_{10} 2 = 0.30$,$1 \ atm = 1 \ bar]$ $[$antilog $(-3.3) = 5.01 \times 10^{-4}]$

At $717 \ K$,$1.50 \ mol$ each of hydrogen and iodine are placed in a $10 \ L$ closed vessel. At equilibrium,$1.25 \ mol$ of each hydrogen and iodine remain. The value of the equilibrium constant $K_c$ for the reaction $H_2(g) + I_2(g) \rightleftharpoons 2HI(g)$ at $717 \ K$ is .......

For the reaction $2HI \rightleftharpoons H_2 + I_2$,which of the following is true?

For the reaction: $HI_{(g)} \rightleftharpoons \frac{1}{2}H_{2(g)} + \frac{1}{2}I_{2(g)}$,the equilibrium constant is $8$. What is the equilibrium constant for the reaction: $H_{2(g)} + I_{2(g)} \rightleftharpoons 2HI_{(g)}$?

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: