Equilibrium: Co^II Complexes

0.5 lab period; work in pairs. Complete the Preparation page before laboratory. The experiment will be done in two shifts of 1.5 hours each.
Goals

To explore the application of Le Chatelier's Principle to chemical equilibrium
To explore the reversibility of a chemical reaction
To enjoy the colors of transition metal ions in various environments

Background

Le Principe du Chatelier (Le Chatelier's Principle) is of great value in understanding and predicting the effects of external stresses on chemical equilibrium. It has been put to good use in the chemical industry to maximize the yields of desired product, and operates in biological systems in the regulation of coupled chemical reactions. The principle is perfectly general, applying to chemical reactions in the gas phase as well as to those taking place in solution.

We will briefly review Le Chatelier's Principle in the context of the generic reaction (1):

(1) aA(sol) + bB(sol) <===> dD(sol)

The equation indicates that the chemical substances A and B react to form substance D in a solvent, indicated by the notation (sol). In many cases the solvent is water, but chemists use a large variety of other solvents as well. Suppose that the reaction establishes equilibrium, which means that the concentrations of the chemical substances reach values that subsequently do not change as long as the system is undisturbed. Let us denote the initial equilibrium concentrations of the reagents as A_o, B_o, and D_o. We then subject the reaction to a series of stresses, allowing it to return to equilibrium following each stress, and use Le Principe du Chatelier to predict qualitatively the effect of the stress on the equilibrium.

Stress 1: Add some A to the solution. Le Chatelier predicts that if A is added to the solution, the equilibrium established prior to the addition is perturbed, and something must happen to minimize the effect of the stress. In this case, the stress is a surfeit of A, so the reaction will take place so as to use up A until equilibrium is reestablished. When the new equilibrium is attained, all three substances will have new concentrations related as follows to the original concentrations:

A₁ > A_o. To relieve the stress caused by adding A, the reaction shifts left to right to use up some of the added A. Of course, it cannot use up all of the added A, so the concentration of A at equilibrium will exceed that present before addition of A
B₁ < B_o. When the reaction shifts to the right to use up added A, B must also be used up. Thus its concentration once the new equilibrium is reached must be less than it was prior to addition of A.
D₁ > D_o. Shift of the reaction toward the right must cause an increase in the concentration of D.

Stress 2: Add some D to the solution. When D is added, the reaction will shift to offset the stress of the excess of D; it will shift toward the left, using up some (but not all) of the added D, and forming both A and B. A new equilibrium will be established with concentrations A₂, B₂, and D₂ such that

A₂ > A₁
B₂ > B₁
D₂ > D₁

Be sure you understand why all three concentrations are larger than before addition of D!

Stress 3: Remove some D from the solution. When D is removed, reaction will take place in an attempt to replace it; the equilibrium will shift to the right, resulting in new concentrations:

A₃ < A₂
B₃ < B₂
D₃ < D₃

Make sure you understand why all three concentrations are smaller than before removal of D!

Finally, we examine the effect of changing the temperature. To precict this we must know whether the reaction is endothermic or exothermic. However, we can turn this around and determine whether the reactioin is endo- or exothermic by watching which way it shifts when temperature is changed! Let's assume that we know the reaction is exothermic. It can thus be rewritten as follows:

(2) aA(sol) + bB(sol) <===> dD(sol) + heat

Stress 4: raise the temperature of the system. Stressing the system by raising the temperature will cause the reaction to shift in an attempt to offset the increase in T; the reaction will shift right to left to use up heat. When a new equilibrium is established, the following concentration conditions will prevail:

A₄ > A₃
B₄ > B₃
D₄ < D₄

In this experiment, we will examine a chemical equilibrium involving two Lewis adducts of the transition metal ion, Co²⁺, which functions as a Lewis acid. The two adducts are CoCl₄^2- and Co(H₂O)₆²⁺, which are related by the following incomplete reaction in ethanol solvent:

(3) CoCl₄^2- + ____ <==> Co(H₂O)₆²⁺ + ____

You will subject the ethanolic equilibrium system to a number of stresses and observe Le Chatelier in action. Based on your observations, you will fill in the missing species in (3), and in addition will determine whether the process is endo- or exothermic as written. Because the equilibrium is not as simple chemically as it appears visually, we will not be able to measure the value of the equilibrium constant for the oversimplified reaction that we will develop. However, your understanding of the qualitative aspects of equilibrium should be greatly increased by your study of this system.

Focus Questions

Write the simplest chemical equation consistent with all of your observations in the qualitative experiments.
Write a narrative explanation of the results of each qualitative experiment in terms of the equation you wrote in question 1.
(Optional)Use the quantitative spectral data to determine whether Co²⁺ has greater affinity for Cl^- or for H₂O.

Equipment and Materials

Labkit
10-mL volumetric flask
50-mL syringe or automatic pipettor
cobalt chloride hexahydrate, CoCl₂.6H₂O
Silver acetate
100-mL beakers (2)
ice
Bunsen burner
ring stand/ring/wire gauze
absolute ethanol
ethanolic 6 M HCl (dilute conc. HCl in an equal volume of absolute EtOH)
distilled water
glass or quartz spectrometer cell
UV-visible spectrometer

Note to instructor: Click here for recipes for preparation of solutions.

Safety

Safety goggles should be worn at all times in the laboratory. Ethanol is flammable, so keep it away from open flames. Avoid contact of 12 M HCl (concentrated) with the skin, and avoid breathing its fumes. In the event of skin contact, flush with water for at least 5 minutes.

Experimental

Record your observations and data in your notebook.

Qualitative Experiments
Carry out the following series of experiments: 1) Weigh 0.5 mmole of cobalt chloride hexahydrate to at least 3 decimal places, and transfer to a 50-mL Erlenmeyer flask. Add exactly 10 mL of absolute ethanol and swirl until the salt has completely dissolved. What do you observe about the solution? Obtain a stopper for the flask.
2) Add water dropwise with stirring or swirling to about 5 mL of the solution from 1). Observe. Save the solution for experiment 3.
3) Add ethanolic 6 M HCl dropwise to the solution from 2). DO NOT ADD EXCESS. Observe. Save the solution for experiment 4.
4) To the solution from 3), add a few drops of H₂O. Observe.
5) To 1 mL of the solution from 1), add a small amount of silver acetate. Observe. Save the solution for experiment 7.
6) Add water dropwise to 1 mL of the solution from 1) until the color looks purple.

Place the solution in a hot water bath and observe.
Now place the solution in an ice bath and observe.

7) What do you think would happen if you were to place the solution from 5) in the hot water bath? Try it and observe. Save the solution for experiment 8.
8) Add a few drops of ethanolic HCl to the hot solution from 7. Observe. Quantitative Experiments (Optional)

Transfer 0.5 mL of the solution from 1) above to a 10-mL volumetric flask, and add ethanol to the mark. Stopper and shake vigorously to homogenize the solution. Transfer 2.5 mL of this solution to a spectrometer cell and stopper. Record the absorbance of the solution at a wavelength of 660 nm.

Titrate the solution in the cell with ethanolic 6 M HCl, using aliquots of 20 or 30 microliters. Aliquots can be measured using either a 50-mL syringe or an automatic pipettor. Shake the cell and record the absorbance at 660 nm after each addition. Continue adding HCl until the absorbance no longer changes. If you use a syringe, thoroughly rinse it with distilled water when finished.

Titrate the solution in the cell with distilled water, using aliquots of 20 microliters. Shake the cell and record the absorbance at 660 nm after each addition. Continue adding water until the absorbance no longer changes.

Clean-up. When you have finished all of your work:

Clean Labkit glassware by recommended procedures, shake off excess water, and return to the Labkit.
Clean Pasteur pipets, graduated pipets, and vials by recommended procedures, shake off excess water, and place in the drying oven.
Rinse the spectrometer cell with water and aspirate dry.
If you used a microliter syringe, rinse it by repeatedly filling it with distilled water and ejecting into the sink. Then rinse similarly with acetone.
Clean and return all borrowed equipment to the instructor before leaving lab.
Clean up your work area before leaving lab.

Disposal Methods

Pour solutions of HCl down the sink with plenty of water. Pour all remaining solutions in the container marked Metal Waste.

References

Martins and da Costa, J. Chem. Ed. 1986, 63, 989.
Ophardt and Smith, J. Chem. Ed. 1980, 57, 453.

Preparation
Equilibrium: Cobalt^II Complexes

Preparation Questions

Equilibrium: CoII Complexes

Equilibrium: Co^II Complexes