In the equilibrium reaction given:
\[ \text{FeSCN}^{2+ (aq)} \rightleftharpoons \text{Fe}^{3+ (aq)} + \text{SCN}^- \]
Initially, the solution appears orange due to the presence of the complex ion \(\text{FeSCN}^{2+}\). Here's how different changes affect the color:
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Addition of Fe³⁺: This would drive the equilibrium to the left according to Le Chatelier's principle, favoring the reverse reaction where the complex ion \(\text{FeSCN}^{2+}\) is formed. As this reaction proceeds, the concentration of \(\text{FeSCN}^{2+}\) will increase, leading to the solution becoming more orange.
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Removal of SCN⁻: This change would also shift the equilibrium to the left (reverse reaction) since the system would try to replace the removed SCN⁻ ions. As a result, the concentration of the orange complex \(\text{FeSCN}^{2+}\) decreases, leading the solution to become less orange (paler or yellowish).
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Removal of FeSCN²⁺: Lowering the concentration of the orange complex would also shift the equilibrium to the left, resulting in a similar color change to that observed upon removing SCN⁻. The color would become less intense and likely appear pale yellow as \(\text{Fe}^{3+}\) and \(\text{SCN}^-\) are predominant.
In summary, a shift from an orange color at equilibrium would indicate:
- If SCN⁻ is removed or if FeSCN²⁺ is removed, the color change would indicate the solution becoming pale yellow due to the increased formation of the reactants (\(\text{Fe}^{3+}\) and \(\text{SCN}^-\)).
- If Fe³⁺ is added, the solution would become more orange as the formation of the complex \(\text{FeSCN}^{2+}\) increases.
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