Materials Included In Kit

Additional Materials Required

Safety Precautions

Copper(II) chloride solution is toxic by ingestion and is a body tissue irritant; avoid contact with body tissues. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please review current Safety Data Sheets for additional safety, handling and disposal information. Remind students to wash hands thoroughly with soap and water before leaving the laboratory.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures governing the disposal of laboratory waste. Copper(II) chloride solution may be washed down the drain with plenty of water. The aluminum wire may be reused. The copper may be disposed of in the solid waste disposal.

Teacher Tips

• This kit contains more than enough chemicals for five classes of 30 students working in pairs to perform two trials of the experiment.
• Two trials should be easily performed by students within a 40–50-minute lab period.
• The aluminum wire is reusable. Each lab group should be able to perform two trials with each piece of wire if they place the unreacted end of the aluminum wire into the solution.
• The test tubes can be washed and reused from class to class.
• It is difficult to prevent the copper from slightly oxidizing during the drying step. This is a source of error.

Sample Data

Data Table

Results Table

Answers to Questions

1. Write the balanced chemical equation for the reaction that occurred between the aluminum wire and the copper(II) chloride solution.

   3CuCl₂•2H₂O(aq) + 2Al(s) →2AlCl₃(aq) + 3Cu(s) + 6H₂O(l)
   or
   3CuCl₂(aq) + 2Al(s) →2AlCl₃(aq) + 3Cu(s)

2. Which starting material in the reaction is the limiting reactant and which material is present in excess?

   To determine which reactant is the limiting reactant, we must determine which reactant will limit the amount of product (Cu) that is formed:

   {12905_Answers_Equation_7}

   The CuCl₂•2H₂O thus limits the amount of product formed and is the limiting reactant while the Al is present in excess.

3. Complete the data table and the results table. Show all calculations on a separate sheet of paper, clearly indicating where each result came from for each box a through g. Sample calculations are shown for Trial I data only. Trial II calculations should be done the same way.
   1. Convert the starting mass of CuCl₂•2H₂O to moles of CuCl₂•2H₂O using its molecular mass.
      {12905_Answers_Equation_8}
   2. Convert the moles of CuCl₂•2H₂O to moles of Al using the balanced equation ratio.
      {12905_Answers_Equation_9}
   3. Convert the number moles of Al to grams of Al using its atomic mass.
      {12905_Answers_Equation_10}

      Therefore, if we start with 0.50 g of CuCl₂•2H₂O, then 0.053 g of Al should react.

   4. {12905_Answers_Equation_11}
   5. {12905_Answers_Equation_12}
   6. {12905_Answers_Equation_13}
   7. {12905_Answers_Equation_14}
4. What physical evidence do you have that shows that copper(II) chloride is, indeed, the limiting reactant?

   The evidence is that there is excess aluminum remaining at the end of the experiment. Also the green-blue CuCl₂ solution becomes completely colorless, indicating that all of the CuCl₂ is “used up.”

5. Discuss reasonable and potential sources of error in this experiment.

   Copper could be lost while scraping it from the wire into the evaporating dish. The copper may still be slightly wet. The copper may have darkened due to oxidation where a layer of CuO forms on the metal surface. The reactants may not have reacted completely.

6. Discuss potential reasons why the percent yield of recovered copper may be greater than 100%.

   The copper may still be slightly wet which would add weight to the final mass. The copper may have oxidized to copper(II) oxide, CuO, which has a greater mass than copper.

7. Why do you think that scientists add excess of one or more chemicals when performing a reaction rather than combine the exact stoichiometric ratio?

   An excess of one of the reactants speeds up the rate of the reaction and thus the rate of production of the product. Also an excess of one of the reactants will assure that all of the limiting reactant is completely consumed; that is, it will push the reaction equilibrium towards product and thus towards completion.

8. What factors do you think may lead scientists (in industry, for instance) to decide to use a certain starting material as the limiting reactant and another as the excess chemical?

   The limiting reactant will generally be the chemical with the higher cost or lower availability. The reactant present in excess must be relatively inexpensive, and easy to separate out from the products.

References

Special thanks to Jeff Bracken, Westerville North High School, Westerville, OH, for bringing this experiment to our attention.

Plumsky, R. J. Chem. Educ. 1996, 73, 451–454.

Zumdahl, S. S. Chemistry; D. C. Heath: Lexington, MA, 1993.

Introduction

The redox experiment between copper(II) chloride and aluminum metal consistently gives impressive results in just a matter of minutes! The progress of the reaction is easily followed by observing the disappearance of the green-blue copper color in the solution and the formation of solid copper on the aluminum wire. Use the balanced equation and stoichiometry calculations to predict the amount of aluminum that should react with the copper(II) chloride and compare this to the actual amount of “leftover” aluminum wire.

Concepts

• Stoichiometry
• Limiting reactant
• Yield calculations
• Oxidation–reducation reactions

Background

Stoichiometry is the branch of chemistry that deals with the numerical relationships and mathematical proportions of reactants and products in chemical reactions. Chemical reactions are represented by balanced chemical equations. Proper interpretation of an equation provides a great deal of information about the reaction it represents and about the substances involved in the reaction. For example, the coefficients in a balanced chemical equation indicate the number of moles of each substance in the reaction. Therefore the ratio of moles of one substance to moles of any other substance in the reaction can be easily determined.

Consider the following unbalanced chemical equation representing the reaction that will be performed in this lab:

CuCl₂(aq) + Al(s) →AlCl₃(aq) + Cu(s)

In the experiment, aluminum wire will be added to an aqueous solution of copper(II) chloride causing a single replacement oxidation–reduction reaction to take place. The oxidation of aluminum metal to aluminum(III) (Al⁰ to Al³⁺) will occur, which is apparent by the dissolving of the aluminum wire to form aluminum chloride (AlCl₃). The simultaneous reduction of copper(II) or copper(II) ions to copper metal (Cu²⁺ to Cu⁰) will occur and solid copper metal will precipitate from solution. As the copper(II) ions are reduced, the green-blue color will fade until the solution is completely colorless—the indication that the reaction is complete.

When performing an experiment such as this, the stoichiometry of the reaction and the actual experimental procedure should be examined to first determine which material is the limiting reactant. The limiting reactant (LR) is the reagent that is used up in the reaction and on which the overall yield of product depends. The quantities of copper(II) chloride and aluminum used in this lab are such that the aluminum is in excess and the copper(II) chloride is the limiting reagent in the reaction (as evident from the title of the lab). The limiting reactant in any reaction, however, can be determined by calculating the starting number of moles of each reactant. The balanced equation is then used to determine which starting material will “run out” first or, in other words, limit the amount of product formed. Consider, for example, the reaction between hydrogen gas and oxygen gas to produce water. The balanced equation is as follows:

2H₂(g) + O₂(g) →2H₂O(l)

If 10.0 grams of H₂ are mixed with 10.0 grams of O₂, which one will “run out” first and act as the limiting reactant? We must first determine the number of moles of each reactant that we are starting with: Then we must determine which of these reactants limits the amount of product formed. Consider H₂ first. The 4.95 moles of H₂ could theoretically produce 4.95 moles of H₂O. This is determined by looking at the balanced chemical equation to determine the stoichiometric ratio, which in this case is two to two. That is, for every two moles of hydrogen that react, two moles of water can be generated: Now consider O₂. The 0.313 mole of O₂ could theoretically produce 0.626 moles of H₂O. This is determined from the stoichiometric ratio, which in this case is one to two. That is, for every one mole of oxygen that reacts, two moles of water can be generated: Therefore, if all of the H₂ reacted, 4.95 moles of H₂O could theoretically form while only 0.626 moles of H₂O could form from the available O₂. The O₂ is therefore the limiting reactant in this example since O₂ limits the amount of H₂O produced. The O₂ will “run out” first while some of the H₂ will remain in excess.

Theoretical yield, then, is the maximum number of grams of product expected from the stoichiometric reaction when the limiting reagent (LR) is completely consumed, with side reactions, reversibility, losses, and the like ignored. The theoretical yield is calculated from the expression: The theoretical yield is the maximum amount of product that can be produced from the quantities of reactants used. However, the amount of product predicted by the theoretical yield is seldom actually obtained due to side reactions, losses, or other complications. The actual yield of product is often given as a percentage of the theoretical yield. This is called the percent yield, which describes the efficiency of the reaction and is calculated from the expression: The percent error for the reaction can be calculated as an indication of accuracy using the expression:

Materials

Safety Precautions

Copper(II) chloride solution is toxic by ingestion and is a body tissue irritant; avoid contact with body tissues. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

1. Set up a hot water bath in a 400-mL beaker using approximately 300 mL of tap water. Bring the water to a boil using a hot plate or Bunsen burner while performing steps 1–5.
2. Obtain approximately 0.50 grams of copper(II) chloride crystals. Use a centigram balance to determine the exact mass to the nearest hundredth of a gram. Record the mass in the data table.
3. Place the copper(II) chloride crystals in a 16 x 125 mm test tube. Add 10 mL of distilled water and stir to dissolve the crystals.
4. Cut a 16- to 18-cm piece of aluminum wire with wire cutters. Use a small piece of steel wool to shine the wire and to remove any surface impurities which may be present on the wire.
5. Measure and record the exact mass of the aluminum wire to the nearest hundredth of a gram in the data table.
6. Use tweezers to lower the aluminum wire into the test tube containing the copper(II) chloride solution. The wire will extend beyond the height of the tube. Place the test tube into the boiling water bath for approximately 5–10 minutes. The reaction is complete when the solution becomes colorless.
7. Use tweezers to remove the wire from the test tube. Use caution so as not to disturb the copper on the wire.
8. Measure the mass of a clean, dry evaporating dish and record the mass in the data table.
9. Use the spatula to scrape as much solid copper as possible from the aluminum wire into the evaporating dish.
10. Allow the solid copper in the evaporating dish to dry overnight or place the dish in a drying oven or under a heat lamp for several hours. (Note: The copper will slightly oxidize and change color to green or dark brown. Overheating will cause the copper to oxidize even faster.)
11. Once dry, measure the mass of the copper and the evaporating dish. Record this mass in the data table.
12. Rinse the aluminum wire with tap water to remove any impurities and dry the wire with a paper towel. Once dry, measure and record the actual “leftover” mass of the aluminum wire in the data table.
13. Repeat this experiment to obtain data for trial II using the other end of the aluminum wire. Record all trial II data in the data table.

Student Worksheet PDF

12905_Student1.pdf