Effect of Temperature on Equilibrium Constant

Which of the following curves between $\log \mathrm{K}$ and $\frac{1}{\mathrm{T}}$ is correct for an endothermic reaction?

The energy profile of the reaction, $\mathrm{A}+\mathrm{B}\rightleftharpoons \mathrm{C}$ is shown as,

The equilibrium constant for the said equilibrium

The equilibrium constants for the reaction ${\mathrm{Br}}_2\rightleftharpoons 2\mathrm{Br}$ at $500 \mathrm{K}$ and $700 \mathrm{K}$ are $1\times {10}^{-10}$ and $1\times {10}^{-5}$, respectively. The reaction is

For the gas phase reaction ${\mathrm{C}}_2{\mathrm{H}}_4+{\mathrm{H}}_2\rightleftharpoons {\mathrm{C}}_2{\mathrm{H}}_6; \mathrm{\Delta H}=-32.7 \mathrm{kcal}$, carried out in a vessel, the equilibrium concentration of ${\mathrm{C}}_2{\mathrm{H}}_4$ can be increased by

${\mathrm{COCl}}_2 \left(\mathrm{g}\right)\rightleftharpoons \mathrm{CO} \left(\mathrm{g}\right)+{\mathrm{Cl}}_{2 }\left(\mathrm{g}\right); {\mathrm{K}}_1=0.329$

$2\mathrm{CO} \left(\mathrm{g}\right)+{\mathrm{O}}_{2 }\left(\mathrm{g}\right)\rightleftharpoons 2{\mathrm{CO}}_2 \left(\mathrm{g}\right); {\mathrm{K}}_2=2.24\times {10}^{22}$

From the given data at $1000 \mathrm{K}$, calculate the equilibrium constant at $1000 \mathrm{K}$ for

$2{\mathrm{COCl}}_{2 }\left(\mathrm{g}\right)+{\mathrm{O}}_{2 }\left(\mathrm{g}\right)\rightleftharpoons 2{\mathrm{CO}}_{2 }\left(\mathrm{g}\right)+2{\mathrm{Cl}}_{2 }\left(\mathrm{g}\right); {\mathrm{K}}_3=\mathrm{x}\times {10}^{20}$. 

 Find the value of $\mathrm{x}$(Nearest integer).

Equilibrium constants $\left({\mathrm{K}}_{\mathrm{p}}\right)$ for the reaction $\frac{3}{2}{\mathrm{H}}_2 \left(\mathrm{g}\right)+\frac{1}{2}{\mathrm{N}}_2 \left(\mathrm{g}\right)\rightleftharpoons {\mathrm{NH}}_3 \left(\mathrm{g}\right)$ are $0.0266$ and $0.0129$ at $350°\mathrm{C}$ and $400°\mathrm{C}$ respectively. Calculate the heat of formation of gaseous ammonia.

Equilibrium constant $\left({\mathrm{K}}_{\mathrm{p}}\right)$ for the reaction, $2{\mathrm{H}}_2\mathrm{S} \left(\mathrm{g}\right)\rightleftharpoons 2{\mathrm{H}}_2 \left(\mathrm{g}\right)+{\mathrm{S}}_{ 2} \left(\mathrm{g}\right)$ is $0.0118$ at $1065°\mathrm{C}$ and the heat of dissociation is $42.4 \mathrm{kcal}$. Find equilibrium constant $\text{ at }1132°\mathrm{C}$.

Would you expect the equilibrium constant for the reaction ${\mathrm{I}}_2 \left(\mathrm{g}\right)\rightleftharpoons 2\mathrm{I} \left(\mathrm{g}\right)$ to increase or decrease as the temperature increases? Why?

[Hint: The forward reaction is endothermic as energy is required to break ${\mathrm{I}}_2$ into $\mathrm{I}$.]