For a given exothermic reaction,$K_p$ and $K'_p$ are the equilibrium constants at temperatures $T_1$ and $T_2,$ respectively. Assuming that the heat of reaction is constant in the temperature range between $T_1$ and $T_2,$ where $T_2 > T_1,$ it is readily observed that:

Consider the reaction $N_{2}O_{4(g)} \rightleftharpoons 2NO_{2(g)}$. The temperature at which $K_{C}=20.4$ and $K_{P}=600.1$ is ....... $K$. (Round off to the nearest integer).
[Assume all gases are ideal and $R=0.0831 \, L \, bar \, K^{-1} \, mol^{-1}$]

For the equilibrium,$2 \ NOCl \ (g) \rightleftharpoons 2 \ NO \ (g) + Cl_{2} \ (g)$,the value of the equilibrium constant,$K_{c}$ is $3.75 \times 10^{-6}$ at $1069 \ K$. Calculate the $K_{p}$ for the reaction at this temperature?

For the reaction,$C_{(s)} + CO_{2(g)} \rightleftharpoons 2CO_{(g)}$,the partial pressures of $CO_2$ and $CO$ are $2.0 \ atm$ and $4.0 \ atm$,respectively,at equilibrium. The $K_p$ of the reaction is

The equilibrium constant for the reaction $N_{2(g)} + O_{2(g)} \rightleftharpoons 2NO_{(g)}$ at temperature $T$ is $4 \times 10^{-4}$. The value of $K_c$ for the reaction $NO_{(g)} \rightleftharpoons \frac{1}{2}N_{2(g)} + \frac{1}{2}O_{2(g)}$ at the same temperature is

The reaction rate for the reaction $[PtCl_4]^{2-} + H_2O \rightleftharpoons [Pt(H_2O)Cl_3]^- + Cl^-$ was measured as a function of concentrations of different species. It was observed that $\frac{-d[[PtCl_4]^{2-}]}{dt} = 4.8 \times 10^{-5} [[PtCl_4]^{2-}] - 2.4 \times 10^{-3} [[Pt(H_2O)Cl_3]^-] [Cl^-]$,where square brackets are used to denote molar concentrations. The equilibrium constant $K_c = ...$. (Nearest integer)