Upper secondary · 12 min

Shifting an equilibrium by concentration: the iron(III) thiocyanate complex

Use the blood-red complex $\ce{Fe^{3+} + SCN^- <=> [FeSCN]^{2+}}$ as a visible sensor of its own equilibrium: each reagent added or removed pushes the colour one way or the other.

Goal

Apply Le Chatelier’s principle to an ionic equilibrium in solution and predict colour changes for six different perturbations, including removal of a reactant.

Apparatus and reagents

Solutions of FeCl₃ 0.1 mol/L and KSCN 0.1 mol/L, AgNO₃, NaF, distilled water, test tubes and droppers.

Procedure

Select the reference equilibrium and note the red intensity of the complex vessel.
Add Fe³⁺, then SCN⁻; in each case predict the shift from $Q<K$ before looking at the colour.
Add Ag⁺ (precipitates AgSCN) and F⁻ (forms colourless $\ce{[FeF6]^{3-}}$ ); explain why removing a partner also shifts the equilibrium.
Dilute the mixture with water and justify the fading using the reaction quotient for a dissolved equilibrium.

What to observe

Adding Fe³⁺ or SCN⁻ deepens the red; adding Ag⁺ or F⁻, which remove free ions, fades it toward yellow.
Dilution pales the colour even though no reagent was consumed: the equilibrium moves toward the side with more dissolved particles.

Explanation

For $\ce{Fe^{3+} + SCN^- <=> [FeSCN]^{2+}}$ ( $K_c \approx 140$ ), comparing $Q$ with $K$ predicts the shift: raising a reactant makes $Q<K$ and drives complex formation; removing Fe³⁺ (as $\ce{[FeF6]^{3-}}$ ) or SCN⁻ (as AgSCN) makes $Q>K$ and dissociates the complex. Dilution lowers all concentrations, favouring the side with more species (2 against 1). $K$ itself stays constant throughout.

History of the experiment

The

\ce{Fe^{3+}/SCN^-}

colour test dates to the 19th century for detecting iron; its use as a Le Chatelier demonstration came from the fact that, unlike

\ce{N2O4 <=> 2NO2}

, every perturbation is a simple drop from a pipette.

Chemists behind it

Jacobus Henricus van ’t Hoff Alfred Werner