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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
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$\mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{SCN}^{-}(\mathrm{aq}) \rightleftharpoons[\mathrm{Fe}(\mathrm{SCN})]^{2+}(\mathrm{aq})$
yellow colourless deep red
Addition of aqueous oxalic acid solution to the above equilibrium
Solution:
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Verified Answer
The correct answer is:
Intensity of deep red color decreases
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.
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