At a definite temperature,the $K_c$ of the following reaction is $0.18$: $PCl_{5(g)} \rightleftharpoons PCl_{3(g)} + Cl_{2(g)}$. In a reaction mixture at the same temperature,$[PCl_3] = 0.042 \ M$,$[Cl_2] = 0.024 \ M$,and $[PCl_5] = 0.005 \ M$. Is this reaction in equilibrium? In which direction will the reaction proceed?

Consider the imaginary equilibrium $4A + 5B \rightleftharpoons 4X + 6Y$. The equilibrium constant $K_c$ has the unit:

$4$ moles of $A$ are mixed with $4$ moles of $B$. At equilibrium for the reaction $A + B \rightleftharpoons C + D$,$2$ moles of $C$ and $D$ are formed. The equilibrium constant for the reaction will be

Write the expression for the equilibrium constant,$K_{c}$ for each of the following reactions:
$(i)$ $2 NOCl_{(g)} \longleftrightarrow 2 NO_{(g)} + Cl_{2(g)}$
$(ii)$ $2 Cu(NO_{3})_{2(s)} \longleftrightarrow 2 CuO_{(s)} + 4 NO_{2(g)} + O_{2(g)}$
$(iii)$ $CH_{3}COOC_{2}H_{5(aq)} + H_{2}O_{(l)} \longleftrightarrow CH_{3}COOH_{(aq)} + C_{2}H_{5}OH_{(aq)}$
$(iv)$ $Fe^{3+}_{(aq)} + 3 OH^{-}_{(aq)} \longleftrightarrow Fe(OH)_{3(s)}$
$(v)$ $I_{2(s)} + 5 F_{2(g)} \longleftrightarrow 2 IF_{5(g)}$

The equilibrium constant for the equilibrium $2HX_{(g)} \rightleftharpoons H_{2(g)} + X_{2(g)}$ is $1 \times 10^{-5}$. What will be the equilibrium concentration of $HX$ if the equilibrium concentrations for $H_2$ and $X_2$ are $1.2 \times 10^{-3} \ M$ and $1.2 \times 10^{-4} \ M$ respectively?

The following concentrations were observed at $500 \ K$ for the formation of $NH_3$ from $N_2$ and $H_2$. At equilibrium: $[N_2] = 2 \times 10^{-2} \ M$,$[H_2] = 3 \times 10^{-2} \ M$ and $[NH_3] = 1.5 \times 10^{-2} \ M$. The equilibrium constant for the reaction is: