Observe the following equilibrium $\mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{SCN}^{-}(\mathrm{aq})…
Observe the following equilibrium
$\mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{SCN}^{-}(\mathrm{aq}) \rightleftharpoons[\mathrm{Fe}(\mathrm{SCN})]^{2+}(\mathrm{aq})$
yellow colourless deep red
Addition of aqueous oxalic acid solution to the above equilibrium
Shifts the equilibrium towards the formation of $[\mathrm{Fe}(\mathrm{SCN})]^{2+}$
Deep red color increases
Intensity of deep red color decreases
No change in equilibrium
Solution
Addition of oxalic acid $\mathrm{H}_2 \mathrm{C}_2 \mathrm{O}_4$ causes the oxalate ions $\mathrm{C}_2 \mathrm{O}_4^{2-}$ to react with $\mathrm{Fe}^{3+}$ ions and form a complex that decreases the concentration of $\mathrm{Fe}^{3+}$ ions. Thus, the equilibrium would shift towards the direction where the concentration of $\mathrm{Fe}^{3+}$ ions would increase which is towards left and this decreases the intensity of deep red colour.