In this experiment you will investigate two important concepts. First you will see an example of an exchange reaction via the redox reaction, Eq. 1 and 2, and secondly you will calculate a percent yield to determine the effectiveness of your experimentation.

A redox reaction involves an exchange of electrons. One species must be oxidized which means it loses electrons so its oxidation state increases (Al0 → Al3+ + 3e-) while the other species must be reduced which means it gains electrons and its oxidation state decreases (Cu2+ + 2e- → Cu0).

The molecular equation for the chemical reaction being carried out today is:

3CuCl2(aq) + 2Al(s) → 3Cu(s) + 2AlCl3(aq)Eq. 1

The chloride ions are spectator ions so the net ionic equation for the reaction is written:

3Cu+2(aq) + 2Al 0(s) → 2Al+3(aq) + 3Cu 0(s) Eq. 2

where the superscripts indicate the oxidation states of the species, positive for the ions and zero for the elements. The changes in oxidation states of the species are more obvious in the net ionic equation although the student should learn to recognize that the form of Equation 1 represents a redox reaction.

The law of conservation of mass states that matter can neither be created nor destroyed; therefore, if you start your experiment with a certain amount of substance and convert this to another substance, then you should end up with the same total amount of substance as you started with.

You will obtain a sample of copper II chloride (CuCl2) and record an exact mass. You will then dissolve this compound in solution and convert the copper II ions (Cu+2) into metallic copper (Cu0) in a reaction with aluminum metal. In theory, you can calculate the number of moles of copper ions with which you started and you can assume that the same number of moles of metallic copper will be obtained. Therefore you will be able to calculate the number of grams of metallic copper that you should recover.

To carry out this chemical reaction you will be using aluminum wire. The aluminum atomsin the wire lose electrons and are considered to be oxidized (Eq. 3). The copper II ions gain electrons and hence are reduced (Eq. 4). At the same time that you are producing metallic copper, the metallic aluminum will be changing from the metallic state Al0 (aluminum with the oxidation state of zero) to dissolved aluminum ion (Al+3). You may assume that you have an excess of aluminum so you need only be concerned with the copper in your calculations for theoretical yield.

Al0 (s) → Al3+ (aq) + 3e- Eq. 3

Cu2+ (aq) + 2e- → Cu0 (s) Eq. 4

If you know how much of the limiting reagent you start with, CuCl2 in this experiment, you can calculate the amount of product that you would expect to produce. But everything does not always go according to plan. There is no flaw with the chemistry but perhaps there are various experimental techniques that will prevent you from obtaining the total amount of copper that you predict. Most often this is due to incomplete reaction, which can easily be seen in this reaction if there is a greenish color still persisting in solution. Other errors in measuring the mass of the final product or perhaps excess moisture will cause the mass that you obtain not to agree with the amount that you expect from your calculation. There is an equation that will let you know how well your experiment was performed by looking at a ratio of what you recovered compared to what you expected to collect. This term is called percent yield and can be calculated according to equation 5.

Eq. 5

Procedure

  1. Obtain a known amount of copper II chloride (slightly >3 g ) and record its mass to the precision ofthe balance. Use a small weigh boat so that you will be able to quantitatively transfer the sample to your reaction beaker.
  1. Place the sample in a 250-mL beaker.
  1. Rinse any residual copper II chloride from the weigh boat into the beaker with deionized water, adding approximately 60 ml of water to the beaker. Stir until the copper II chloride has completely dissolved. The resulting solution is blue or blue green. This color is characteristic of Cu+2 ions in water.
  1. Obtain a piece of aluminum wire with a mass of about 1.5 grams. These are precut so you may assume that the portion is large enough.
  2. Twist one end of the aluminum wire into a coil; bend the other end into a hook and hang the wire over the side of the beaker so that the coil is immersed in the copper solution.
  1. Observe the reaction that occurs. Note that with time the blue color of the solution will fade as the aluminum reacts, and copper metal will accumulate on the surface of the aluminum. Shake the wire frequently (preferably continuously) to loosen the copper that accumulates on the surface of the wire. The reaction has reached completion when the blue color in the solution has disappeared and the aluminum wire has no more copper forming on it (~ 20 minutes).
  1. When the reaction is complete, remove the aluminum wire and shake the aluminum so that all of the copper falls back into the beaker. Note that if aluminum accidentally falls or the wire breaks into the beaker, you will be isolating impure copper (copper mixed with aluminum), so be sure that only the copper remains in the beaker. Use a wash bottle of deionized water to rinse the aluminum wire over the beaker to remove any residual copper. You might find it helpful to scrape the copper off the aluminum with the rubber tip on a glass stirring rod.
  1. Observe any change in appearance of the wire and then discard the aluminum in the appropriate waste container, which is separate from the liquid waste.
  1. Obtain a piece of filter paper that fits into a Buchner funnel. Record the mass of this filter paper along with a watch glass to the precision ofthe balance.
  1. Assemble a vacuum filtration apparatus as described by your instructor.
  1. Isolate the copper by vacuum filtration. Carefully rinse all the copper in the beaker into the Buchner funnel using a deionized water wash bottle.
  1. When the copper has been transferred, wash the copper with about 20 mL of acetone. The acetone will wash the water away and evaporate more quickly than water. This will hasten the drying of the copper. Allow the vacuum to dry the copper in the funnel for five minutes.
  1. Carefully transfer the filter paper and copper to the preweighed watchglass using a scoopula or spatula.
  1. Place the watch glass and contents into the oven to dry for 10 minutes. Allow it to cool to room temperature for a minimum of 10 minutes. Record its mass.

Pre-Lab:Where appropriate show calculations below. Include correct significant figures and units on the lines.

  1. Calculate the Mass of copper (II) chloride CuCl2
  2. What is the molar mass of CuCl2?
  3. How many moles of CuCl2 were dissolved in solution?
  4. Write a balance equation for the dissociation of CuCl2 in water.
  5. Given the number of moles of CuCl2 dissolved in solution and the balanced equation above, how many moles of copper II ions (Cu2+) were initially in solution?
  6. How many moles of copper metal (Cu0) do you expect to be formed (See Eq. 2)?
  7. What is the molar mass of Cu0?
  8. What is the theoretical yield of copper metal in grams?

Experimental Data and Calculations

Mass of copper (II) chloride CuCl2

Where appropriate show calculations in the boxes below. Include correct significant figures and units on the lines.

Mass of filter paper and watchglass

Mass of filter paper, watchglass, and copper

Mass of copper recovered

Discussion and Post Lab Analysis:

  1. Describe any changes to the aluminum wire during the reaction.
  2. Consider equation 1 and the mass of copper you recovered. Show a calculation to determine the mass of aluminum that reacted.
  3. Did you collect less copper than expected or apparently more?
  4. What experimental errors occurred during your experiment that may have caused something other than 100% recovery? Use complete sentences and cite more than one example. (Think about your experimental technique)