Initially $1\mathrm{.}50$ moles of $N_2(g)$ and $3\mathrm{.}50$ moles of $H_2(g)$ were added to a $1L$ container at $700C.$ a result of the reaction

$N_2(g)+3H_2(g)2N{H}_3$

the equilibrium concentration of $N{H}_3(g)$ became $0\mathrm{.}540M.$ What is the value of the equilibrium constant for this reaction at the given temperature of $700C.$

$8.$

Initially, $1\mathrm{.}0$mol of $NO(g)$ and $1$mol of $C{l}_2(g)$ were added to a $1L$ container. As a result of the reaction

$2NO(g)+C{l}_2(g)2$NOCl$(g)$

the equilibrium concentration of NOCl$(g)$ became $0\mathrm{.}96M.$ Using the RICE table methodology determine the value of the equilibrium constant $K_C$ for this reaction.