Cobalt Chloride Equilibrium: Influence of Concentration and Temperature

An equilibrium exists between a hydrated cobalt species and anhydrous cobalt chloride, both Co ions have an oxidation state of 2+. The equilibrium equation representing the system is

[Co(H2O)6]2+(aq) + 4 Cl- <–> [CoCl4]2-(aq) + 6H2O  Keq = 1.7x10-3
    
    (pink)                                    (blue)

Initially, a beaker contains a red-pink solution of cobalt (II) chloride, present as [Co(H2O)6]2+ ions and chloride ions.  If the chloride or cobalt concentrations increase, the equilibrium will also shift towards the blue anhydrous cobalt chloride. For example, when hydrochloric acid is added, the added chloride ions shift the equilibrium position in favour of blue [CoCl4]2- ions and water.   Add water, however, and the equilibrium will shift back towards the pink hydrated species.

Test tubes containing a pink solution of cobalt and chloride ions are placed in hot water and cold water. The tube placed in hot water will turn blue. The tube placed in cold water will turn more pink.  A change in temperature or concentration of the ions will shift the equilibrium. If heat is added, the equilibrium will shift towards the cobalt chloride complex, which is blue in color.  Cooling will shift the products towards the hydrated complex, which is more pink.  Is the reaction, as written left to right, endothermic or exothermic

The instructor will have to interpret and narrate the demonstration as the changing equilibrium of cobalt ion complexes in solution is difficult to follow.

With a Keq of 1.7 x 10-3  the equilibrium favors the reactants but some products are also present. The reactants are of a different color(pink) than the products(blue). When reactants predominate, the solution looks pink and when products predominate, the solution looks blue.

For temperature changes

[Co(H2O)6]2+(aq) + 4 Cl- <–> [CoCl4]2-(aq) + 6H2O  + Heat
    
    (pink)                                          (blue)

The addition of heat favors the production of the blue complex, whereas removing heat favors the production of the pink complex to restore the lost energy.

What causes the pink and blue colors? Cobalt ions in salts are positive ions with a 2+ charge. Positive ions attract negative particles such as chloride ions (Cl-) and the oxygen end of water molecules (H2O) sine the oxygen in in water has a partial negative charge. When most of the negative species around the cobalt ion are water molecules, the ion absorbs light so that it appears pink. When the water molecules bonded to the Cobalt ions are replaced, the negative chloride form an IMF to the positive cobalt ions, and the cobalt appears blue. The water molecules form an octahedral arrangement around the Co(II) ion.  The the chloride ions from a tetrahedral arrangement around the Co(II) ion.  Teh different arrangment of ligands cause the electrons in the cobalt system to absorb different energies. These different energies result in different colors absorbed by the cobalt.

1. L. Summerlin, J. Ealy; Chemical Demonstrations: A Sourcebook for Teachers. Volume 1; 1985; p. 53.
2. B.Z. Shakhashiri.  Chemical Demonstrations: A Handbook for Teachers of Chemistry; Wisconsin; 1983; Vol. 1; p 280-285.

Research Questions

Predict what effect removing chloride ions will have on the equilibrium.

Predict what effect adding chloride ions will have on the equilibrium.

Is the reaction as written in the forward direction endothermic or exothermic?

Curriculum Notes 

Students have difficulty understanding this demostration for several reasons.  First, at this point in their study of chemistry students do not relaize that the label only provides how the initial concentrations were prepared, i.e. 0.10 M CoCl2.  The label does not indicate what is in the solution.  Second, students often think colors of solutions are due to acid-base indicators.  Prior to doing this demonstration, point out there are no acid-base indicators in the solution. If reminded, students will acknowledge Cu(NO3)2(aq) solution has a blue color, but they often fail to realize that the blue color is due to [Cu(H2O)6]2+(aq).  Third, before doing this demonstrations, students need help in understanding complex ions.

LeChatelier’s Principle: If a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change.

Lead Time 
Two days of lead time is required for this project.
Discussion 

Demonstration showing the changing equilibrium of cobalt complexes in solution. Initially, the beaker contains a red-pink solution of cobalt (II) chloride, present as        [Co (H2O)6] 2+ions and chloride ions. When hydrochloric acid is added, the added chloride ions shift the equilibrium position in favour of blue [CoCl4]2- ions and water. As more chloride ions are added, formation of the [CoCl4]2- ions is favored, as this changes the concentration of chloride ions in solution.The equilibrium can shifted to the left by adding water, turning the solution pink.  As water is added, the system moves to remove water from the soluiton by complexing it with the cobalt ions. This shifting of the equilibirum system can be repeated about three times.

https://www.sciencephoto.com/media/600219/view/cobalt-chloride-equilibri...

Materials 
  • 2 Pre-prepared test tubes with cobalt chloride and HCl (pink solutions) in beaker
  • 2 Pre-prepared test tubes with aqueous cobalt chloride (purple solutions) in beaker
  • hot plate
  • hot water bath
  • ice water bath
  • Reagents
  • 20 mL of 0.2 M Cobalt chloride solution: 26 g of CoCl2 per liter of water.
  • ~20 mL of 0.1 M Silver nitrate solution: 1.7g of AgNO3 per 100 mL of water.
  • ~40 mL of 6M hydrochloric acid
  • ~20 mL of Distilled water
Procedure 
•Temperature Influence
    Select two test tubes that are the same color (indicating the same concentration of Co2+ ions)
•Heat one, observe the color change
•Determine whether the reaction in the forward direction is exothermic or endothermic.
 
Concentration influence
•Select two test tubes that are the same color (indicating the same concentration of Co2+ ions)
•Add Cl- to one, observe the color change
•Determine whether the equilibrium system shifts to the right or left.

 

Safety Precautions 

Use caution working with acidic solutions (HCl is corrosive). Wear proper protective equipment.

Footnotes 

1. L. Summerlin, J. Ealy; Chemical Demonstrations: A Sourcebook for Teachers. Volume 1; 1985; p. 53
2. B.Z. Shakhashiri.  Chemical Demonstrations: A Handbook for Teachers of Chemistry; Wisconsin; 1983; Vol. 1; p 280-285

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© Copyright 2012 Email: Randy Sullivan, University of Oregon Chemistry Department and UO Libraries Interactive Media Group