Aim
The purpose of this experiment is to investigate how a system in equilibrium responds to a change in concentration of components in the mixture and to understand the chemical reactions involved in equilibrium shifts.
Introduction
Equilibrium is a state in a chemical reaction where the rates of the forward and reverse reactions are equal, resulting in constant concentrations of reactants and products. In a dynamic equilibrium, the reaction doesn’t stop, but the concentrations of the reactants and products remain constant over time.
Iron(III) ions and thiocyanate ions react in solution to produce thiocyanatoiron(III), a complex ion, according to the equation:
Fe³⁺(aq) + SCN⁻(aq) ⇌ Fe(SCN)²⁺(aq)
Pale yellow + colourless ⇌ blood-red
The colour produced by the complex ion can indicate the position of equilibrium. When the concentration of one of the reactants or products changes, the equilibrium position shifts according to Le Chatelier’s principle, which states that when a stress is applied to a system in equilibrium, the system will adjust to counteract the stress and restore equilibrium.
Requirements
- safety glasses
- 4 test-tubes and test-tube rack
- 2 teat-pipettes
- distilled water
- potassium thiocyanate solution, 0.5 M KSCN
- iron(III) chloride solution, 0.5 M FeCl₃
- ammonium chloride, NH₄Cl
- spatula
- glass stirring rod
Procedure
- Mix together one drop of 0.5 M iron(III) chloride solution and one drop of 0.5 M potassium thiocyanate solution in a test-tube and add about 5 cm³ of distilled water to form a pale orange-brown solution.
- Divide this solution into four equal parts in four test-tubes.
- Add one drop of 0.5 M iron(III) chloride to one test-tube. Add one drop of 0.5 M potassium thiocyanate to a second.
- Compare the colours of these solutions with the original samples. Record your observations.
- Add a spatula-full of solid ammonium chloride to a third test-tube and stir well. Compare the colour of this solution with the remaining tube and note your observation. Ammonium chloride removes iron(III) ions from the equilibrium by forming complex ions such as FeCl₄⁻. A possible reaction is: Fe³⁺(aq) + 4Cl⁻(aq) → FeCl₄⁻(aq)
Interpretation of Results
Having made three observations, suggest a cause for each colour change (in terms of the concentrations of the coloured species) and then suggest what can be inferred about a shift in the position of equilibrium. If a pattern has emerged, then you can make a prediction based on the results of the experiment.
Results Table
| Observation | Explanation |
|---|---|
| [Fe³⁺] increased | |
| [SCN⁻] increased | |
| [Fe³⁺] decreased |
Conclusion
The results of this experiment demonstrate that changes in the concentration of reactants or products in a system at equilibrium cause a shift in the equilibrium position according to Le Chatelier’s principle. In this case, increasing the concentration of either Fe³⁺ or SCN⁻ ions caused the equilibrium to shift to the right, forming more Fe(SCN)²⁺ complex ions and resulting in a darker red colour. Conversely, decreasing the concentration of Fe³⁺ ions by adding NH₄Cl caused the equilibrium to shift to the left, forming fewer Fe(SCN)²⁺ complex ions and leading to a lighter red or even yellow colour.
These shifts in equilibrium position help maintain the balance between the forward and reverse reactions when the system experiences changes in concentration, restoring the system to a new equilibrium state.
Questions
- How would the position of equilibrium be affected by increasing the concentration of FeSCN²⁺?
- For each imposed change, show how the shift in equilibrium position conforms to Le Chatelier’s principle.
- Increasing the concentration of FeSCN²⁺ would cause the equilibrium to shift to the left, as the system would counteract the stress by reducing the amount of FeSCN²⁺ and producing more Fe³⁺ and SCN⁻ ions.
- a. Increasing [Fe³⁺]: The equilibrium shifts to the right to reduce the excess Fe³⁺ ions by producing more Fe(SCN)²⁺ complex ions, thus restoring equilibrium. b. Increasing [SCN⁻]: The equilibrium shifts to the right to reduce the excess SCN⁻ ions by forming more Fe(SCN)²⁺ complex ions, thereby restoring equilibrium. c. Decreasing [Fe³⁺]: The equilibrium shifts to the left to compensate for the reduced Fe³⁺ ions by breaking down Fe(SCN)²⁺ complex ions into Fe³⁺ and SCN⁻ ions, reestablishing equilibrium.
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