To determine which change would cause the orange solution (resulting from the simultaneous presence of FeSCN²⁺ and its dissociation products, Fe³⁺ and SCN⁻) to become pale yellow, let’s consider the equilibrium established by the dissociation of iron thiocyanate:
\[ \text{FeSCN}^{2+}(aq) \rightleftharpoons \text{Fe}^{3+}(aq) + \text{SCN}^-(aq) \]
-
In this equilibrium, FeSCN²⁺ contributes to the red color, while Fe³⁺ and SCN⁻ contribute to a pale yellow and colorless solution, respectively.
-
If the solution is currently orange, it indicates that both FeSCN²⁺ and its dissociation products are present in sufficient concentration.
To shift the equilibrium to the left and produce more of the uncolored FeSCN²⁺, thus reducing the concentrations of Fe³⁺ and SCN⁻—which results in a more pale yellow color—we can consider the options given:
-
Adding KSCN (which dissociates into K⁺ and SCN⁻): This would increase the concentration of SCN⁻, shifting the equilibrium to the right and promoting further formation of FeSCN²⁺, giving a darker red solution, which is not what we want.
-
Adding an ion that causes SCN⁻ to precipitate: This would decrease the concentration of SCN⁻, causing the equilibrium to shift to the left, producing more FeSCN²⁺ and potentially resulting in a lighter color.
-
Adding FeCl₃ (which dissociates into Fe³⁺ and Cl⁻): This would increase the concentration of Fe³⁺, shifting the equilibrium to the right and leading to more FeSCN²⁺, which is not desired as it could also deepen the red color.
-
Adding an ion that causes Fe³⁺ to precipitate: This would decrease the concentration of Fe³⁺, shifting the equilibrium to the left towards the formation of more FeSCN²⁺, producing a lighter solution.
Given this analysis, the best response would be:
Adding an ion that causes SCN⁻ to precipitate.
This option effectively reduces the concentration of SCN⁻ in the solution, allowing the equilibrium to shift to the left and favor the formation of FeSCN²⁺, thus resulting in a pale yellow solution.