Materials Included In Kit

Additional Materials Required

Safety Precautions

Nitric acid is severely corrosive, a strong oxidant and toxic by ingestion and inhalation. Hydrochloric acid solution is corrosive to skin and eyes and toxic by ingestion and inhalation. Sodium hydroxide solution is a corrosive liquid, can cause skin burns and is very dangerous to eyes. Sulfuric acid solution is corrosive to eyes, skin and other tissue. Avoid contact of all acids and bases with eyes and all body tissue. Clean up all spills immediately; neutralize any acid spills with a weak base; neutralize any base spills with a weak acid; rinse with water. 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.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The final copper product can be disposed of in the solid waste disposal according to Flinn Suggested Disposal Method #26a. Remaining amounts of reagents included in the kit can be saved for later use or disposed of according to the appropriate Flinn Suggested Disposal Methods.

Teacher Tips

• Part A—Caution students about the hazards of the poisonous red-brown NO₂ gas that is released from the reaction. Be sure to perform this step in an operating fume hood. Also caution students about the hazards of using nitric acid.
• Part B—Have students turn on the hot plates at the start of Part B. Have them place a beaker of water on the hot plate. This will save time when they get to Part C.
• Part C—Before the lab, it is a good idea to demonstrate the correct procedure for folding filter paper, and for setting up a properly-working funnel. Warn students to always add and wash with a minimum amount of water. The more water, the longer the filtering process will take.
• Part E—It will take approximately 80 drops of 8 M NaOH to produce a slight murkiness in step 20. Students can add these drops quickly at first to reduce the time of this step.
• Part G—If time is a factor, the evaporating dishes may be dried faster by placing them in a drying oven set on low, or by placing them under a heating lamp. However, too much heat will oxidize the copper to copper oxide as indicated by darkened copper or a black coating on the copper. 15 mL of 2 M H₂SO₄ may be needed to completely dissolve the excess zinc in step 36. The gas evolved in step 36 may be very irritating. Run step 36 in an operating fume hood.

  General Tips

• Enough materials are provided for a class of 30 students working in pairs.
• This lab will most likely take four to five 50-minute lab periods. It will be necessary to preserve solutions or precipitates from one period to the next. To do this, cover the beaker of solution or filter containing the precipitate by using another large beaker, Parafilm M^®, a watchglass, or a piece of filter paper. Material should be labeled to show its identity and class period. For day one, a good stopping point is step 7 of Part B; for day two, step 18 of Part D, and for day three, either step 28 of Part E or step 33 of Part F.
• We found that it took the following lengths of time:
  • Part A: 15 minutes
  • Part B: 10 minutes
  • Part C: 40–50 minutes depending on the filtration stop
  • Part D: 30 minutes
  • Part E: 40–60 minutes
  • Part F: 30 minutes
  • Part G: 40–50 minutes
• The Background section is provided as optional material. This section is printed on its own separate sheet and contains the formulas, equations, and reaction types of the complete copper reaction cycle. The Background actually answers many of the post-lab questions for the student. You may consider copying the background for student use before the lab for introductory chemistry or physical science students. For honors or higher-level chemistry students, you may consider providing this information after the lab or not at all.
• A prelaboratory assignment sheet is provided. It is recommended that the lab handout and prelab sheet be given to students at least one day prior to beginning actual lab work. The prelab should then be checked and/or discussed before lab.
• Explain to students the necessity of recording detailed lab observations in the data table as they do the lab. Students should label their pipets to avoid accidental contamination.
• Make sure the students have completed their final rinses of precipitates in the filtering steps before completing any lab session. If filter paper then dries out and “unseats” from the filter, the student can then carefully use a stirring rod and wash bottle to reset the filter paper. Caution the students to avoid puncturing the filter paper.

Answers to Prelab Questions

Read the entire lab and then complete the Prelaborabory Assignment. This must be done and checked before beginning work in the laboratory.

1. Write the correct chemical formula for each substance listed in the first column. Then read through the laboratory procedure to determine where each material is found in the procedure, either as a product or as a reactant. List the corresponding part(s) of the procedure in the third column.
   {11860_PreLabAnswers_Table_1}
2. Define each of the following terms:
   1. reagent—stock chemical which must be kept pure
   2. decantation—process of carefully pouring off the liquid while leaving the precipitate behind
   3. filtration—process used to separate an insoluble solid from a liquid
   4. precipitate—a solid that makes a sudden appearance in a liquid or a gas
   5. filtrate—the liquid that has passed through the filter paper
   6. supernatant liquid—the liquid floating on top of a precipitate
3. What does the M in 8 M HNO₃ and in 6 M HCl stand for?

   The M stands for molarity in moles per liter, or solution concentration. The higher the molarity, the more concentrated is the solution.

4. Strong acids and bases are very damaging to the eyes as well as to body tissue in general. For this reason, what must be worn at all times while in the laboratory (even if you are not personally working with chemicals)?

   Chemical splash goggles and a chemical-resistant apron must be worn at all times. Chemical-resistant gloves are recommended.

5. NO₂ is a red-brown, reactive, poisonous gas. It is evolved in the reaction in Part A. For this reason, reaction A must be done in a fume hood.
6. Imagine that two beakers containing colored, but clear, solutions were mixed. How could you determine if precipitation had occurred? What would you observe?

   To determine if precipitation had occurred, you would look for formation of a solid material forming in the solution. You would observe this as solid particles, cloudiness, and/or murkiness.

7. If precipitation occurs, how could you separate the solid from the solution?

   To separate the solid from a solution, you could filter the solution using a filtration setup or you could decant off the liquid, leaving the solid in the beaker.

8. Copper compounds are frequently colored. Imagine a student is doing Part E of this lab. This student is filtering the Cu₃(PO₄)₂. The filtrate is slightly blue.
   1. Why might the filtrate be blue?

      The problem could be that some solid went through the filter paper, or there may be incomplete precipitation.

   2. Should the filtrate be discarded? Why or why not?

      No, the filtrate should not be discarded because some copper would be lost.

   3. If the filtrate is discarded, how will this affect the final mass of the copper that was recovered?

      If the filtrate is discarded, the final mass of copper will be less.

   4. How might a blue color in the filtrate have been avoided?

      Be sure precipitation is complete. Make sure the filter paper is securely in the funnel and that no solid went through.

Sample Data

Answers to Questions

1. For each reaction, write a balanced stoichiometric equation. Note: Parts B, E and G include a second reaction in which excess reagents are used up. Include these reactions as well.
   1. Include all physical states, using the following abbreviations: s = solid, l = liquid, g = gas, aq = aqueous
   2. Write the color of each material under the formula in the balanced equations.
   3. Label the reaction type for each equation (whether it is a double replacement, single replacement, decomposition, or synthesis reaction).
   {11860_Answers_Reaction_1}
2. What is the maximum temperature to which the evaporating dish containing copper will be heated in Part G? How is this known?

   The maximum temperature to which the copper in the evaporating dish will be heated is 100 °C. This is known because it is heated over a beaker of boiling water.

3. What will happen if the copper is overheated during the drying process? How will this affect the percent recovery?

   If the copper is overheated during the drying process, some may oxidize into copper oxide, CuO. This will cause the percent recovery to be more than it actually is because copper oxide is heavier than copper.

4. Comment on the physical appearance of the reclaimed copper. How could it be made to look like the original copper?

   The reclaimed copper is brownish in color, clumpy in texture, and not shiny like the starting copper. In order to make it look like the original copper, it would first need to be pure and dry. It would then have to be melted down at its melting temperature of 1063 °C and drawn into copper turnings.

5. Calculate the percent recovery of the copper by means of the formula for percent yield. (Percent yield = mass of copper recovered/mass of copper used x 100%)

   Example calculation: If 0.42 grams of copper were recovered from the experiment and 0.51 grams were used to begin the experiment, then the percent yield would be 82%, as shown below:

   {11860_Answers_Equation_2}
6. Calculate the amount of copper lost (or gained) during the reaction series.

   Example calculation: If 0.42 grams of copper were recovered from the experiment and 0.51 grams were used to begin the experiment, then the amount of copper lost during the series of reactions is 0.09 grams, as shown:

   0.51 g at start – 0.42 g at end = 0.09 g Cu lost

7. List some possible experimental errors that might lead to a mass of reclaimed copper less than that originally used.

   Student answers will vary. Possible answers include (a) some copper material lost on the stirring rod during filtering steps (b) some copper material left in filter paper during filtering steps (c) copper lost during the decanting step (d) not enough zinc used in step G; thus some copper still left in solution, as seen by the solution still being slightly blue.

8. List some possible experimental errors that might lead to a mass of reclaimed copper greater than that originally used.

   Student answers will vary. Possible answers include (a) copper is still wet when measuring the final mass (b) some of the final copper has oxidized to copper oxide, as seen by a dark coating on the copper (c) some zinc metal is still left in the evaporating dish with the copper (d) other impurities are mixed in with the copper adding to its weight.

Introduction

Start with copper—end with copper! Perform a series of chemical reactions to convert copper metal to cupric nitrate, cupric hydroxide, cupric oxide, cupric chloride, cupric phosphate, cupric sulfate and finally... back to copper. An amazing chemical feat which incorporates many standard laboratory procedures!

Concepts

• Chemical reactions
• Law of conservation of matter

Background

The purpose of this laboratory experiment is to: (a) illustrate different types of chemical reactions (b) show how a quantity of an element can be carried through a series of chemical transformations without significant loss of mass, thereby illustrating the law of conservation of matter and (c) provide experience in fundamental laboratory procedures such as transferring a reagent from a reagent bottle, transferring a solution or a solid from one vessel to another, decanting, filtering, washing and dissolving a precipitate.

Experiment Overview

A weighed quantity of copper will be carried through the following transformations:

Cu → Cu(NO₃)₂ → Cu(OH)₂ → CuO → CuCl₂ → Cu₃(PO₄)₂ → CuSO₄ → Cu

In Part A, copper metal is oxidized by nitric acid to produce a blue solution containing cupric nitrate, Cu(NO₃)₂, and a brown poisonous gas of nitrogen dioxide, NO₂, as shown in Equation A: In Part B, the blue copper(II) nitrate solution, Cu(NO₃)₂, reacts with sodium hydroxide, NaOH, in a double replacement reaction to produce a blue precipitate of copper(II) hydroxide, Cu(OH)₂ according to Equation B: An acid–base neutralization reaction also occurs in Part B between sodium hydroxide, NaOH, and the excess nitric acid, HNO₃, from Part A according to Equation B2: In Part C, the blue copper(II) hydroxide solid, Cu(OH)₂, is decomposed with heating into the black copper(II) oxide solid, CuO, according to the decomposition reaction shown in Equation C: In Part D, the black copper(II) oxide, CuO, undergoes an acid–base neutralization reaction with hydrochloric acid, HCl, to form a green copper(II) chloride salt solution, CuCl₂, and water according to Equation D: Part E involves a double replacement reaction between the green copper(II) chloride, CuCl₂, and sodium phosphate, Na₃PO₄, to produce a blue precipitate of copper(II) phosphate, Cu₃(PO₄)₂, and sodium chloride, NaCl, as shown in Equation E: An acid–base neutralization reaction also occurs in Part E between sodium hydroxide, NaOH, and the excess hydrochloric acid, HCl, from Part D to Equation E2: Part F involves another acid–base neutralization reaction as the sulfuric acid, H₂SO₄, dissolves the blue copper(II) phosphate, Cu₃(PO₄)₂, to produce a blue copper(II) sulfate solution, CuSO₄, and phosphoric acid, H₃PO₄, as shown in Equation F: The final reaction in Part G involves the replacement of a less active metal (copper) by a more active metal (zinc). The zinc, Zn, replaces the copper in copper(II) sulfate, CuSO₄, forming zinc sulfate, ZnSO₄, and solid copper, Cu, according to the single replacement redox reaction shown in Equation G: The excess zinc in Part G, over and above that required to replace all of the copper, is dissolved by sulfuric acid, H₂SO₄, according to another single replacement redox reaction shown in Equation G2:

Materials

Prelab Questions

Read the entire lab and then complete the pre-laboratory assignment. This must be done and checked before beginning work in the laboratory.

1. Write the correct chemical formula for each substance listed in the first column. Then read through the laboratory procedure to determine where each material is found in the procedure, either as a product or as a reactant. List the corresponding part(s) of the procedure in the third column.
   {11860_PreLab_Table_1}
2. Define each of the following terms:
   1. reagent
   2. decantation
   3. filtration
   4. precipitate
   5. filtrate
   6. supernatant liquid
3. What does the M in 8 M HNO₃ and in 6 M HCl stand for?
4. Strong acids and bases are very damaging to the eyes as well as to body tissue in general. For this reason, what must be worn at all times while in the laboratory (even if you are not personally working with these chemicals)?
5. NO₂ is a red-brown, reactive, poisonous gas. It is evolved in the reaction in Part ____. For this reason, reaction ____ must be done ____________________________.
6. Imagine that two beakers containing colored, but clear, solutions were mixed. How could you determine if precipitation had occurred? What would you observe?
7. If precipitation occurs, how could you separate the solid from the solution?
8. Copper compounds are frequently colored. Imagine a student is doing Part E of this lab. This student is filtering the Cu₃(PO₄)₂. The filtrate is slightly blue.
   1. Why might the filtrate be blue?
   2. Should the filtrate be discarded? Why or why not?
   3. If the filtrate is discarded, how will this affect the final mass of the copper that was recovered?
   4. How might a blue color in the filtrate have been avoided?

Safety Precautions

Nitric acid is severely corrosive, a strong oxidant and toxic by ingestion and inhalation. Hydrochloric acid solution is corrosive to skin and eyes and toxic by ingestion and inhalation. Sodium hydroxide solution is a corrosive liquid, can cause skin burns and is very dangerous to eyes. Sulfuric acid solution is corrosive to eyes, skin and other tissue. Avoid contact of all acids and bases with eyes and all body tissue. Clean up all spills immediately; neutralize any acid spills with a weak base; neutralize any base spills with a weak acid; wipe up with water. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron.

Procedure

Part A. Preparation of Copper(II) Nitrate From Copper by Oxidation with Nitric Acid

1. In a 250-mL beaker, weigh out about 0.50 grams of pure copper turnings. Record the exact mass of copper to the nearest hundredth of a gram in the data table.
2. Place the beaker containing the copper in a fume hood. Use a Beral-type pipet to carefully add 5 mL of 8 M HNO₃, dropwise, to the beaker. Caution: Nitric acid is severely corrosive. Be sure to wear chemical splash goggles, chemical-resistant gloves, and a chemical-resistant apron.
3. Warm the beaker gently under a heating lamp or on a hot plate set on low until the copper is completely dissolved. Record detailed observations in the data table. If necessary, add a few more drops of 8 M HNO₃ to dissolve any remaining copper. Caution: The brown fumes released are poisonous; do not remove the beaker from the fume hood until all of the copper is reacted and no more gas is released.
4. Dilute the solution with 10 mL of distilled or deionized water.

Part B. Preparation of Copper(II) Hydroxide From Copper(II) Nitrate

5. To the solution from Part A, add 5 mL of 8 M NaOH, a little at a time, with constant stirring. 6. Allow the precipitate to settle a bit. Test for complete precipitation by adding 1 more drop of 8 M NaOH to the clear supernatant liquid. If no additional precipitate forms, the reaction is complete. If additional precipitate does form, add more 8 MNaOH dropwise until precipitation is complete.
6. Test the acidity of the liquid by touching a stirring rod first to the solution and then to a piece of 1–12 pH paper. The liquid should be basic, turning the pH paper to blue. Since the solution itself is blue, be sure to look at the spread of the liquid asit touches the paper. If the solution is not basic, add 8 M NaOH dropwise until it is basic.

Part C. Preparation of Copper(II) Oxide From Copper(II) Hydroxide by Decomposition

8. To the mixture from Part B, add enough distilled or deionized water to give a total volume of about 100 mL.
9. Using a hot plate, boil the mixture gently for about 3 minutes, stirring constantly, until the mixture turns a different color.
10. Prepare a filtration setup as illustrated in Figure 1. Carefully fold a piece of filter paper and fit it into a clean funnel. Moisten the filter paper so that it fits snugly into the funnel. Test the filtration system with a small amount of water to be sure the water flows freely and rapidly. This will save a great deal of time in the filtering process.
    {11860_Procedure_Figure_1_Filtration setup}
11. Filter the mixture. Do not add so much solution into the filter that it rises over the sides of the filter paper. Wash any traces of solid from the beaker using a stream of water from a wash bottle.
12. Wash the solid precipitate in the filter paper twice with hot distilled or deionized water. To do this, add enough water to completely cover the precipitate and allow it to run through.
13. Discard the filtrate; leave the product in the filter paper for Part D.

Part D. Preparation of Copper(II) Chloride From Copper(II) Oxide

14. Place a clean, 250-mL beaker below the funnel containing the precipitate from Part C.
15. Dissolve the solid by carefully pouring 8 mL of 6 M HCl directly into the funnel. Allow the resulting solution to run into the beaker.
16. If the solid does not completely dissolve when the HCl runs through the filter the first time, the liquid that has run through should be poured back through the filter. Be sure to use a second clean beaker below the funnel. This process should be repeated until the solid completely dissolves.
17. If the dissolving of the solid seems to have stopped, add more 6 M HCl in 1-mL portions.
18. When the solid is completely dissolved, the filter should be washed twice with a minimum amount of cold distilled or deionized water and the washings allowed to run into the beaker with the solution. Be sure to rinse the second beaker with water and add the washings to the solution.

Part E. Preparation of Copper(II) Phosphate From Copper(II) Chloride

19. Cool the solution from Part D to about room temperature by placing the beaker containing the solution into a large beaker of cold water.
20. Neutralize the solution by first adding 8 M NaOH until the liquid in the beaker acquires a slight murkiness or cloudiness. Add about 2 mL of NaOH, then add it drop by drop, with constant stirring, until it becomes cloudy.
21. Then add 6 M HCl, drop by drop with constant stirring, until the cloudiness just disappears. The solution should be neutral at this point and have a pH of about 7. Test the pH by touching a stirring rod first to the solution and then to a piece of 1–12 pH paper. Add more NaOH or HCl as necessary to adjust the pH to 7.
22. Add 18 mL of 0.3 M Na₃PO₄, slowly and with constant stirring. Observe precipitate formation.
23. Allow the precipitate to settle a bit. Test for complete precipitation by adding 1 more drop of 0.3 M Na₃PO₄ to the liquid. If no additional precipitate forms, the reaction is complete. If additional precipitation does occur, add more 0.3 M Na₃PO₄ dropwise until precipitation is complete. If the quantity of precipitate is so large that a thick slurry is formed, add a small amount of water to give a suspension in which the precipitate is free to settle.
24. Using a hot plate, very carefully heat the solution until it just begins to show signs of boiling. This will make filtration easier by helping to consolidate the precipitate. Do not allow the mixture to boil vigorously.
25. Prepare a filtration setup as described in step 10 and illustrated in Figure 1.
26. Filter the mixture. Do not add so much solution into the filter that it rises over the sides of the filter paper. Wash any traces of solid from the beaker using a stream of water from a wash bottle. Clean beaker for use in step 29.
27. Wash the solid precipitate in the filter paper twice with hot distilled or deionized water. To do this, add enough water to completely cover the precipitate and allow it to run through.
28. Discard the colorless filtrate; leave the product in the filter paper for Part F.

Part F. Preparation of Copper(II) Sulfate From Copper(II) Phosphate

29. Place a clean, 250-mL beaker below the funnel from Part E containing the precipitate.
30. Dissolve the solid by carefully pouring 12 mL of 2 M H₂SO₄ directly into the funnel. Allow the resulting solution to run into the beaker.
31. If the solid does not completely dissolve when the H₂SO₄ runs through the filter the first time, the liquid that has run through should be poured back through the filter. Be sure to use a second clean beaker below the funnel. This should be repeated until the solid completely dissolves.
32. If the dissolving of the solid seems to have stopped, add more 2 M H₂SO₄ in 1-mL portions.
33. When the solid is completely dissolved, the filter should be washed twice with a minimum amount of cold distilled or deionized water and the washings allowed to run into the beaker with the solution. Be sure to rinse the second beaker with water and add the washings to the solution.

Part G. Preparation of Copper From Copper(II) Sulfate

34. Use a weighing dish to measure out about 2 g of mossy zinc. Add the zinc to the solution from Part F.
35. Allow the solution to stand, with occasional stirring, until it is entirely colorless. This indicates that all of the copper has been removed from solution. If absolutely necessary, add a little more zinc to the solution.
36. Once the solution is colorless, the excess zinc must be dissolved. To do this, place the beaker in an operating hood and add 2 M H₂SO₄ with stirring until the excess zinc is completely dissolved. It may require 10–18 mL of H₂SO₄, so add the acid in 2–3 mL portions at first and then 1-mL portions. Warm the solution to speed the rate of solution. Complete dissolution of zinc will be indicated by the fact that the precipitate is uniformly copper-colored and no more hydrogen gas bubbles are produced.
37. Allow the copper to settle. Decant off and discard the supernatant liquid. In the process of decanting, the liquid is carefully poured off while the solid remains behind. To avoid loss of solid, the liquid is not drained off completely; some is allowed to remain behind with the solid.
38. Wash the copper in the beaker three times with cold distilled or deionized water. To do this, add about 20 mL of water to the precipitate in the beaker. Stir thoroughly with a stirring rod, allow the precipitate to settle, and decant off the supernatant clear liquid. Repeat this two more times.
39. Weigh a clean, dry evaporating dish and record the mass in the data table.
40. Use a small amount of water to flush the copper from the beaker into the evaporating dish. Decant the supernatant liquid from the dish.
41. Dry the copper by heating the evaporating dish on top of a boiling water bath over a Bunsen burner or on a hot plate set on low, as illustrated in Figure 2. The heat of the escaping steam causes the water in the dish to evaporate and dries the contents of the dish without overheating the copper.
    {11860_Procedure_Figure_2_Boiling water method for drying the copper}
42. When the copper in the evaporating dish is dry, wipe the outside of the dish dry with a clean towel. Allow the dish to cool to room temperature and weigh the dish and the copper. Record this mass in the data table.
43. Save your recovered copper. Place it in a vial labeled with your name and turn it in to your instructor.
44. Consult your instructor for appropriate disposal procedures.

Student Worksheet PDF

11860_Student1.pdf