2020-05-14
Le Chatelier’s Principle
Observe the shift in equilibrium between ferric ions (Fe³⁺) and thiocyanate ions (SCN⁻) based on Le Chatelier’s Principle by altering the concentration of either component. The reaction forms a blood-red complex, [FeSCN]²⁺, whose color intensity indicates the direction of the equilibrium shift.
Fe³⁺ (pale yellow) + SCN⁻ (colorless) ⇄ [FeSCN]²⁺ (red)
A deep red solution was prepared by mixing ferric chloride (FeCl₃) and potassium thiocyanate (KSCN). The solution was diluted and divided among test tubes for different treatments: adding water, Fe³⁺, SCN⁻, or oxalate (Na₂C₂O₄). The change in the color intensity of each test tube was compared to evaluate how the equilibrium shifted.
Test Tube A: Served as a reference, containing only the diluted red solution, to compare color changes.
Test Tube B: Addition of Fe³⁺ deepened the red color, indicating a shift towards the right as more [FeSCN]²⁺ complex formed.
Test Tube C: Adding SCN⁻ also intensified the red color, showing that equilibrium shifted right to produce more [FeSCN]²⁺.
Test Tube D: The addition of oxalate lightened the red color, suggesting that oxalate binds with Fe³⁺, shifting equilibrium to the left, reducing the complex.
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