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Overview: |
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- Some chemical reactions are reversible; that is, not only do the reactants react to form products, but the products can in turn reform into the original reactants. When a reversible system reaches a point at which the rate of the forward reaction equals the rate of the reverse reaction, the system is said to be at equilibrium. At equilibrium, no observable changes in the system can be noted. It is important to understand that at equilibrium all reaction participants are present – all reactant particles, as well as all product particles. This does not mean, however, that all are present in equal amounts.
- Le Chatalier’s Principle tells us that if a system at equilibrium is subjected to a stress, the system will shift in order to minimize the effect of that stress. In this laboratory exercise you will examine how various stresses cause equilibrium systems to shift. These particular equilibrium systems undergo color changes as the equilibrium shifts between the reactants and products, allowing you to see which side of the reaction becomes favored.
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Purpose
- To observe the effect of various stresses (ion concentration; temperature) on equilibrium systems.
The Reactions
Part 1. Chromate – Dichromate Equilibrium
2 CrO42-(aq) + 2 H3O+(aq) |
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2 Cr2O72-(aq) + 3 H2O(l) |
yellow |
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orange |
Part 2. Iron(III) – Thiocyanate Ion Complex
Fe3+(aq) + SCN-(aq) |
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Fe(SCN)2+(aq) |
light brown |
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red |
Part 3: Cobalt(II) Chloride Complex;
Effect of Temperature
[CoCl4]2-(aq) + 6 H2O(aq) |
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[Co(H2O)6]2+(aq) + 4 Cl-(aq) |
blue |
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pink |
Equipment and Materials
Amounts needed per group are approximate
Part 1. Chromate – Dichromate Equilibrium
- 0.1 M K2CrO4 5 mL per group
- 0.1 M K2Cr2O7, 5 mL per group
- 1 M HCl, dropper bottle
- 1 M NaOH, dropper bottle
- 2 test tubes in a test tube rack
Part 2. Iron(III) – Thiocyanate Ion Complex
- 0.1 M FeCl3, 10 mL per group
- 0.1 M KSCN, 10 mL per group
- 0.1 M KCl, 5 mL per group
- 10 mL graduated cylinder
- 250 mL or larger beaker
- distilled water, approx 100 mL
- 4 test tubes in a test tube rack
- dropper pipette
Part 3. Cobalt(II) Chloride Complex;
Effect of Temperature
- 0.2 M acidified CoCl2 · 6 H2O, 15 mL per group
- hot water bath (approx. 90°C)
- ice water bath
- 3 test tubes in a test tube rack
Procedure
Part 1. Chromate – Dichromate Equilibrium
- Half-fill a test tube with potassium chromate, K2CrO4 (Tube 1).
- Half-fill a second test tube with potassium dichromate, K2Cr2O7 (Tube 2).
- To Test Tube 1 add several drops of HCl. HCl is an acid; adding HCl increases the concentration of H3O+ ions in the equilibrium system. Note the color change.
- After recording the color change in Test Tube 1, add several drops of NaOH. NaOH is a base; adding a base decreases the concentration of H3O+ ions in the equilibrium system. Record the color change.
- To Test Tube 2 add several drops of NaOH until a color change is observed.
- After recording the color change, add several drops of HCl to Test Tube 2. Again note the change in color.
Part 2. Iron(III) – Thiocyanate Ion Comlex
- Pour 5 mL of 0.1 M FeCl3 into the beaker.
- After rinsing the graduated cylinder, measure 5 mL of 0.1 M KSCN. Add to the beaker containing the FeCl3 . Note the color change.
- Add enough distilled water to the beaker to dilute the solution to a light brown color. Pour some into a test tube to check the color.
- Pour about 10 mL of this solution into each of the four numbered test tubes. The first test tube will serve as a control.
- To Test Tube 2 add several drops of FeCl3 until a color change is observed. Adding more FeCl3 increases the concentration of Fe3+ in solution. Record the color change.
- To Test Tube 3 add several drops of KSCN until a color change is observed. Adding more KSCN increases the concentration of SCN- in solution. Record the color change.
- To Test Tube 4 add several drops of KCl until a color change is observed. Adding KCl causes the concentration of Fe3+ to decrease because theFe3+ and Cl- react to form FeCl4-.
Part 3. Cobalt(II) Chloride Complex;
Effect of Temperature
- Fill three test tubes approximately half-full with the acidified CoCl2 · 6 H2O solution.
- Test tube 1 will serve as the control. Keep this test tube at room temperature. Record the initial colour of the solution.
- Place the second test tubes in the hot water bath. After a few minutes a colour change will occur. Record the colour.
- Place the third test tube in the cold water bath. Record any colour change.
- Reverse tubes 2 and 3. Observe any colour changes that occur.
Results
Part 1. Chromate – Dichromate Equilibrium
Solution |
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Colour Change |
K2CrO4 |
initial colour |
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HCl added |
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NaOH added |
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K2Cr2O7 |
initial colour |
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NaOH added |
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HCl added |
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Part 2. Iron(III) – Thiocyanate Ion Comlex
Test
Tube |
Stress
Applied |
Initial Color |
Final Color |
1 |
Control |
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-- |
2 |
Fe3+ added |
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3 |
SCN- added |
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4 |
Cl - added: decreases [Fe3+] |
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Part 3. Cobalt(II) Chloride Complex; Effect of Temperature
Temperature |
Solution Colour |
room temperature |
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hot water bath |
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cold water bath |
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Conclusions and Questions
- Use Le Châtelier’s Principle to explain the color changes observed in both test tubes with the addition of both HCl and NaOH.
- Use Le Châtelier’s Principle to explain the color changes observed in Test Tubes 2 – 3 upon the addition of FeCl3, KSCN, and KCl.
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Based on the colour changes observed in the hot water and cold water baths, determine whether the forward reaction is endothermic or exothermic.
Rewrite the equation with a simple energy term (“ + heat”) included on the appropriate side of the equation. You may find it easier if you begin by using only the terms “heat”, “pink”, and “blue” in your equation.
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