THE COBALT CHLORIDE EQUILIBRIUM.

 

BACKGROUND INFORMATION The element cobalt can form compounds in two different oxidation states, +2 and +3. The +2 state is more common. The ion Co2+ (aq) is pink. Other compounds of cobalt(II), which include both anhydrous Co2+ and complex ions, are commonly blue.

If an aqueous solution contains both cobalt(II) and chloride ions, the blue ion CoCl42- forms, in equilibrium with the pink Co2+ (aq) ion.

CoCl42- (aq) <===========> Co2+ (aq) + 4Cl1- (aq)
Blue Pink

At relatively low concentrations of chloride, the equilibrium lies far to the right, and the solution is pink. If there is a large concentration of excess chloride, the equilibrium tends to the left, and the solution tends to be blue.

The equilibrium is sensitive to temperature as well as to concentration of solutes. At lower temperatures, the equilibrium tends to lie to the right, that is, to be more pink; at higher temperatures, it lies to the left and appears more blue.

 

AIM To investigate the effects of temperature and concentration upon the position of equilibrium in a solution of cobalt(II) containing excess chloride ions.

 

PLANNING
By trial and error, it is known that a solution in which [Co2+ (aq)] = 0.5M and [Cl-(aq)] = 5M will be blue at room temperature but pink if refrigerated.

If concentrations of ions are changed by addition of solid cobalt(II) chloride, concentrated hydrochloric acid, or water, then predictable changes in colour are likely to be observed. For example, if a solution that is pink (equilibrium to the right) increases in its concentration of either cobalt(II) or of chloride, then its colour should change towards blue.

If water is added to a blue solution, its colour should change towards pink.

For an experiment involving dilution with water, a solution that is blue at room temperature is needed. If such a solution is diluted just sufficiently for a distinct change towards pink to be observed, then it may be possible to reverse the change and restore blueness by addition of either concentrated hydrochloric acid or solid cobalt(II) chloride.

 

MATERIALS REQUIRED
To make 25 mL of solution in which [Co2+ (aq)] = 0.5M and [Cl-(aq)] = 5M:

balance, CoCl2.6H2O = 3.0 g, concentrated HCl = 12.0 mL, 25 mL volumetric flask, 25 mL graduated cylinder, balance, small funnel, deionised water, dropper. (This amount should be sufficient for about six students, working individually.)

To test effects of changes of concentration of solutes:
white glazed tile, white glazed well-plate, 1 mL graduated pipette with filler bulb (or automatic delivery pipette), pasteur pipettes, small spatula, four small glass stirring rods, small stock supplies of concentrated hydrochloric acid and of solid cobalt(II) chloride, deionised water.

 

INSTRUCTIONS
1. Weigh 3.00 g of solid cobalt(II) chloride and put it in a small funnel placed in the mouth of a 25 mL volumetric flask.

2. Measure 12 mL of concentrated hydrochloric acid in a graduated cylinder, then pour it slowly over the solid cobalt chloride in the funnel, so that the crystals dissolve and fall into the flask. Rinse the cylinder with a little deionised water and pour the rinse water slowly over any remaining crystals in the funnel. Continue to drop water slowly until all the crystals have fallen or dissolved into the flask, then finish filling the flask to the 25 mL graduation, and shake. The solution should be dark blue-purple in colour.

3. To show the reversibility of the reaction with changing temperature, place the flask, or a sample of solution taken from the flask into a small closed tube or bottle, into the coldest part of a refrigerator. Observe how the colour changes. Then remove the flask or sample from the refrigerator and watch the change of colour as the temperature of the solution rises.

Steps 4 to 10, require the use of a 4 x 3 glazed well-plate, or a suitable alternative.

4. Label the wells of a white glazed well-plate 1 to 12. Transfer 0.40 mL of the blue solution, at room temperature, into ten wells, leaving wells 5 and 9 empty.

At each of steps 5 to 9, all observed changes of colour should be recorded.

5. Using a pasteur pipette, add one drop of water to each of wells 2, 3 and 4, two drops of water to each of wells 6, 7 and 8, and three drops of water to each of wells 10, 11 and 12. (Well 1 should contain a sample of the initial solution, as a control.) Stir each sample gently with a small glass rod.

6. To each of wells 3, 7 and 11, add one drop of concentrated hydrochloric acid.

7. To each of wells 4, 8, and 12, add equal small amounts of crystals of cobalt(II) chloride.

8. With a small stirring rod, stir wells 3, 7, and 11; rinse and dry the rod after each use.

9. With a small stirring rod, stir wells 4, 8, and 12.; rinse and dry the rod after each use.

10. Record the final pattern of colours seen in the wells 1 to 12.

 

OBSERVATIONS
Design and construct tables to record the colours observed.

(Colours might be recorded as "blue", "pinkish-blue", "bluish pink", or "pink". Variations, comparisons, or trends of colour should also be recorded.)

 

Table one: effect of addition of water to solutions.

Sample number 1 2,3,4 6, 7, 8 10, 11, 12
Water added None 1 drop 2 drops 3 drops
Final colour

Table two: effects of addition of a drop of concentrated HCl to solutions 3, 7 and 11.

Sample number 1 3 7 11
Water added none 1 drop 2 drops 3 drops
Initial colour
Colour before stirring no acid
added
Colour after stirring

Table three: effects of addition of solid cobalt(II) chloride to solutions 4, 8, and 12.

Sample number 1 4 8 12
Water added none 1 drop 2 drops 3 drops
Initial colour
Colour before stirring no CoCl2
added
Colour after stirring

EVALUATION
1. How well do the observed changes in the colours of the solutions agree with the expected changes?

  1. with respect to the effect of dilution with water upon the position of the equilibrium?
  2. with respect to the effect of an increase in the concentration of chloride ions upon the position of the equilibrium?
  3. (with respect to the effect of an increase in the concentration of cobalt(II) ions upon the position of the equilibrium?

2. How can any changes of colour observed in the solutions as a result of stirring be explained?


Questions? Comments??
Mike Clark