The enthalpy changes for the following processes are listed below

$\begin{array}{l}{\mathrm{Cl}}_2\left(g\right)=2\mathrm{Cl}\left(g\right);\hspace{0.25em}\hspace{0.25em}\hspace{0.25em}\hspace{0.25em} 242.3\mathrm{kJ}{\mathrm{mol}}^{-1}\\ {\mathrm{I}}_2\left(g\right)=2\mathrm{I}\left(g\right);\hspace{0.25em}\hspace{0.25em}\hspace{0.25em}\hspace{0.25em} 151.0\mathrm{kJ}{\mathrm{mol}}^{-1}\\ \mathrm{ICl}\left(g\right)=\mathrm{I}\left(g\right)+\mathrm{Cl}\left(g\right);\hspace{0.25em} \hspace{0.25em}\hspace{0.25em}\hspace{0.25em}211.3\mathrm{kJ}{\mathrm{mol}}^{-1}\\ {\mathrm{I}}_2\left(g\right)={\mathrm{I}}_2\left(g\right); \hspace{0.25em}\hspace{0.25em}\hspace{0.25em}\hspace{0.25em}62.76\mathrm{kJ}{\mathrm{mol}}^{-1}\end{array}$

Given that the standard states for iodine and chlorine are ${\mathrm{I}}_2\left(\mathrm{s}\right)$ and ${\text{Cl}}_2\left(g\right)$, the standard enthalpy of formation for $\mathrm{ICl}\left(\mathrm{g}\right)$is :