le Châtelier's Principle: common ion effect

Copper (II) ions react with halide ions (Cl^-, Br^-, I^-) to form complex ions.  While Cu²⁺ (aq) ions are blue, the resulting product is green.  The following equation can be used to represent this process.

Cu²⁺ (aq) + 4Cl^- (aq) CuCl₄^2- (aq)

We can use le Châtelier's principle to predict what should happen as we change the concentration of the ions in solution.  Adding more Cl^- will shift the equilibrium to the right, so the solution will become more green.  One interesting result of this is that the color of copper(II) chloride (CuCl₂)solutions depends on concentration: when dilute there is a smaller amount of Cl^- so the color is blue.  However, as the concentration increases the solution becomes more of a green-blue color, rather than a true blue.

There are other ways to change the concentration of Cl^- ions without dissolving more CuCl₂

• you could change the amount of water (either add more, or evaporate some) since these are aqueous solutions

• you could add another substance that reacts to remove either Cl^- or Cu²⁺ ions
• you could add the ions from another chemical compound containing either Cl^- or Cu²⁺ ions

The last method, adding the same ion from a different source is called adding a "common ion".  This is a standard chemical procedure, and the change predicted by le Châtelier's principle when we mix the same ions from different sources, is called the "common ion effect".  For example, we would have a common ion of Cl^- if we dissolved some NaCl in the solution.  We would have a common ion of Cu²⁺ if we dissolved some Cu(NO₃)₂ (s) in the solution.

In this experiment you will examine the common ion effect on a simple solution equilibrium.

Materials:

• 20 mL of 0.5 M CuCl₂ solution
• about 3 g of NaCl and 1.5 g of NaNO₃
• small squares of weighing paper (about 2 cm x 2cm each)
• 6 18 x 150 mm test tubes
• calibrated pipet

Procedures:

Caution: Copper (II) chloride is moderately dangerous (harmful if swallowed, causes irritation to skin, eyes and respiratory tract), so be careful not to ingest it, spill on your skin, or get it into your eyes.  If it does come in contact with your skin, wash thoroughly with plenty of water.  The other substances used in this experiment are relatively safe; however, you need to observe the normal safety precautions you would for any lab while using them.  Wear a lab apron and safety goggles at all times.

1. Get about 20 mL (about 1/2 a test tube full) of a 0.5 M solution of CuCl₂ in a test tube.
2. Measure onto separate weighing papers 0.5 g, 1.0 g, and 1.5 g of NaCl.  Note: in order to save time while weighing, you can follow this optional method.
3. Measure 1.5 g of NaNO₃ onto a fourth weighing paper.
4. Using a calibrated pipet put 4 mL of the CuCl₂ solution into each test tube.
5. Leave the first test tube as is.  Add 0.5 g of NaCl to test tube 2.  Add 1.0 g of NaCl to test tube 3.  Add 1.5 g of NaCl to test tube 4.
6. Add 1.5 g of NaNO₃ to test tube 5.
7.  Stopper each tube with a rubber stopper, shake, and observe the contents.   Note: it may be hard to get all the solid to dissolve, as you are very near a saturated solution of the solid in the higher mass tubes.

Conclusion:

Observe the color of each tube, and explain the results in terms of le Châtelier's principle. 

Optional method of weighing NaCl:  On one piece of weighing paper, measure 3 g of NaCl.  Divide it into 6 equal piles.   Transfer one pile (0.5 g) to one piece of weighing paper.  Transfer two piles (1.0 g) to a second piece of paper.  The remaining three piles (1.5 g) can stay on the original paper.  This saves time if balances are limited.