Question

Observe the following equilibrium $$\text{Fe}^{3+}(aq) + \text{SCN}^-(aq) \rightleftharpoons [\text{Fe(SCN)}]^{2+}(aq)$$ yellow \hspace{1cm colourless \hspace{1cm} deep red}
Addition of aqueous oxalic acid solution to the above equilibrium

• A. Shifts the equilibrium towards the formation of $[\text{Fe(SCN)}]^{2+}$
• B. Deep red color increases
• C. Intensity of deep red color decreases
• D. No change in equilibrium

Hint:

Le Chatelier's principle is crucial for predicting the direction of equilibrium shifts when conditions like concentration, temperature, or pressure are changed. Remember that removing a reactant or product will shift the equilibrium towards the side where that species is present.

Answer 1

The Correct Option is C

Step 1: Understand the reaction of oxalic acid with Fe$^{3+$ ions.}
Oxalic acid (H$_2$C$_2$O$_4$) reacts with Fe$^{3+}$ ions to form stable complex ions:
$$\text{Fe}^{3+}(aq) + \text{C}_2\text{O}_4^{2-}(aq) \rightleftharpoons [\text{Fe(C}_2\text{O}_4)]^{1+}(aq)$$ This reaction removes Fe$^{3+}$ ions from the equilibrium system.
Step 2: Apply Le Chatelier's principle.
The given equilibrium is:
$$\text{Fe}^{3+}(aq) + \text{SCN}^-(aq) \rightleftharpoons [\text{Fe(SCN)}]^{2+}(aq)$$
The removal of a reactant (Fe$^{3+}$) will shift the equilibrium to the left to counteract this change and restore equilibrium.

Step 3: Analyze the effect on the concentration of $[\text{Fe(SCN)}]^{2+$.}
The shift to the left means the reverse reaction is favored, consuming $[\text{Fe(SCN)}]^{2+}$ ions and decreasing their concentration in the solution.
Step 4: Relate the concentration of $[\text{Fe(SCN)}]^{2+$ to the color intensity.}
The deep red color of the solution is due to the presence of $[\text{Fe(SCN)}]^{2+}$ ions. A decrease in the concentration of these ions will lead to a decrease in the intensity of the deep red color. Final Answer: The final answer is $\boxed{Intensity of deep red color decreases}$