Observe the following equilibrium:Fe^{3+}_{(a | vedclass

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(C) Addition of oxalic acid $(H_2C_2O_4)$ provides oxalate ions $(C_2O_4^{2-})$,which react with $Fe^{3+}$ ions to form a stable complex $[Fe(C_2O_4)_

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Intensity of deep red color decreases

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(C) Addition of oxalic acid $(H_2C_2O_4)$ provides oxalate ions $(C_2O_4^{2-})$,which react with $Fe^{3+}$ ions to form a stable complex $[Fe(C_2O_4)_3]^{3-}$.This reaction significantly decreases the concentration of free $Fe^{3+}$ ions in the solution.According to Le Chatelier's principle,the equilibrium will shift to the left to compensate for the loss of $Fe^{3+}$ ions.As the equilibrium shifts to the left,the concentration of the deep red complex $[Fe(SCN)]^{2+}$ decreases,leading to a decrease in the intensity of the deep red color.

Observe the following equilibrium:
$Fe^{3+}_{(aq)} + SCN^{-}_{(aq)} \rightleftharpoons [Fe(SCN)]^{2+}_{(aq)}$
(yellow) (colorless) (deep red)
Addition of aqueous oxalic acid solution to the above equilibrium:

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Observe the following equilibrium:$Fe^{3+}_{(aq)} + SCN^{-}_{(aq)} \rightleftharpoons [Fe(SCN)]^{2+}_{(aq)}$(yellow) (colorless) (deep red)Addition of aqueous oxalic acid solution to the above equilibrium: