Explain the effect of adding $(i)$ Oxalic acid $(H_2C_2O_4)$,$(ii)$ $HgCl_2$,and $(iii)$ Potassium thiocyanate $(KSCN)$ on the equilibrium reaction: $Fe^{3+}(aq) + SCN^-(aq) \rightleftharpoons [Fe(SCN)]^{2+}(aq)$ (deep red color).

(N/A) The given equilibrium is: $Fe^{3+}(aq) + SCN^-(aq) \rightleftharpoons [Fe(SCN)]^{2+}(aq)$.
$(i)$ Addition of $H_2C_2O_4$: Oxalic acid provides $C_2O_4^{2-}$ ions which react with $Fe^{3+}$ to form a stable complex $[Fe(C_2O_4)_3]^{3-}$. This decreases the concentration of free $Fe^{3+}$ ions. According to Le Chatelier's principle,the equilibrium shifts to the left,causing the deep red color to fade.
$(ii)$ Addition of $HgCl_2$: $Hg^{2+}$ ions react with $SCN^-$ to form a very stable complex $[Hg(SCN)_4]^{2-}$. This decreases the concentration of $SCN^-$ ions. The equilibrium shifts to the left,and the deep red color fades.
$(iii)$ Addition of $KSCN$: This increases the concentration of $SCN^-$ ions. According to Le Chatelier's principle,the equilibrium shifts to the right to consume the excess $SCN^-$,resulting in an increase in the intensity of the deep red color.