Electroplating can be used to deposit a layer of metal such as chromium, copper, gold, nickel, or zinc on another metal. This deposit can provide a protective and decorative coating for the metal which lies beneath it. This laboratory activity will show you how to electroplate a copper coating on a car key or other suitable metallic object.

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To investigate the mass changes, if any, at each electrode during electroplating and to calculate the charge on the copper ion through application of Faraday's Laws.

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Wear protective glasses and an apron at all times. Avoid skin contact with solids and solutions. Dispose of the solutions as designated by your teacher. Wash your hands before leaving the laboratory.

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Part I: Cleaning Key and Electrode

Prepare a data table to record changes in mass of the key and copper electrode and to record time and current if directed by your teacher.

1. Clean a key and copper electrode by rubbing with fine steel wool. Wash the key and copper electrode with a detergent solution and rinse both with tap water.
2. Attach a 5-cm length of #20 - 22 bare copper wire to your car key. The wire serves as a "handle" to remove the key from the cleaning solutions.
3. Soak the key and copper electrode for a few minutes in 30 mL of 3 M sodium hydroxide solution, NaOH. (Caution: Avoid contact of skin with sodium hydroxide.) Remove with tweezers and rinse in distilled water.
4. Soak the key and copper electrode for a few minutes in 30 mL of 3 M sulfuric acid, H₂SO₄. (Caution: Avoid contact of skin with sulfuric acid.) Remove with tweezers and rinse with distilled water.

Part II: Conditioning the Electrodes (Optional)

1. Place a 1 cm x 10 cm copper strip into a 250 mL beaker. Bend the strip to fit over the top of the beaker (Fig. 1).
2. Support the key in the solution by wrapping the copper wire around a small glass rod. Rest the rod on top of the beaker.
3. Hook up the variable DC power supply (Fig. 1). The key should be wired directly to the negative (black terminal) electrode. The copper electrode should be wired to the positive (red terminal) electrode through the ammeter.

   

4. Have your teacher check the set-up before proceeding with Step 5. (Data Check: Obtain your teacher's initials.)
5. Turn on the power supply and adjust the current to 0.25 A. Allow the current to flow for about five minutes.
6. Turn off the power supply. Remove the key and copper strip.
7. Rinse with distilled water and blot dry with a paper towel.

Part III: Copper Plating the Key

1. Determine the mass of the key and copper electrode separately to the nearest 0.01 g.
2. Add 200 mL of copper plating solution to a 250 mL beaker.
3. Place a 1 cm x 10 cm copper strip into the beaker. Bend the strip to fit over the top of the beaker (Fig. 1).
4. Support the key in the solution by wrapping the copper wire around a small glass rod. Rest the rod on top of the beaker.
5. Hook up the variable DC power supply (Fig. 1). The key should be wired directly to the negative (black terminal) electrode. The copper electrode should be wired to the positive (red terminal) electrode through the ammeter. Have your teacher check the set-up before proceeding with Step 6. (Data Check: Obtain your teacher's initials.)
6. Turn on the power supply and begin timing. Adjust the current to 0.25 A. Monitor the current to maintain a constant reading during the electroplating.
7. After 30 minutes, turn off the power supply and timer.
8. Remove the key. Carefully rinse the key with distilled water and allow to dry.
9. Determine the mass of the key to the nearest 0.01 g.
10. Repeat Steps 8 and 9 for the copper electrode.
11. Disconnect the wires, pour the plating solution into storage bottle, and rinse and dry the beaker.
12. Wash hands thoroughly before leaving the laboratory.

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1. What observations suggest that chemical reactions occurred?
2. Where in the plating cell do these reactions occur?
3. Compare the change in mass of the negative electrode with the mass change of the positive electrode. Explain any differences or similarities.
4. How many (a) coulombs, (b) faradays and (c) moles of electrons were transferred by the power supply?
5. 1. Use the mass of copper deposited on the key (negative electrode) to determine the moles of copper ions deposited at this electrode (the cathode).
   2. Use the mass of copper lost by the copper strip (positive electrode) to calculate the moles of copper atoms lost at this electrode (the anode).
6. From the results of Calculations 4c and 5a, what is the ratio between the moles of electrons transferred by the power supply and the moles of copper ions deposited at the cathode (key)?
7. Use the result from Calculation 6 to determine the charge on a copper ion.
8. Write the reduction half-reaction that occurred at the cathode.
9. From the results of Calculations 4c and 5b, what is the ratio between the moles of electrons transferred and moles of copper atoms lost by anode (copper strip)?
10. Use the result obtained in Calculation 9 to write the oxidation half-reaction that occurred at the anode.
11. Bonus: How many copper ions per second were reduced at the cathode?

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1. Must the surface to be electroplated be a metal? Explain in terms of the processes you observed.
2. What would happen if the wiring were reversed?
3. Would it be possible to use carbon electrodes for this laboratory activity? (Carbon electrodes are non-reactive and conductive.) What would happen at the cathode? Explain.

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Teachers Guide

Preparing for the Laboratory Activity

• Major Chemical Concepts
• Level
• Expected Student Background
• Time
• Safety
• Materials
• Advance Preparation

Conducting the Laboratory Activity

• Pre-lab Discussion
• Student/Teacher Interaction
• Anticipated Student Results
• Answers to Questions in Student Version
• Post-Lab Discussion
• Possible Extensions

Assessing the Laboratory Learning

• Laboratory Practical
• Pencil/Paper Items
• Answers to Pencil/Paper Items

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Preparing for the Laboratory Activity

Major Chemical Concepts

The chemical changes occurring at electrodes in an electrolytic cell result from the transfer of electrons by the power supply. These changes are oxidation-reduction reactions. Faraday's Laws allow quantitative predictions regarding these electrode reactions.

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Advanced chemistry. May be used for general chemistry as a non-quantitative activity.

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Students should be able to:

• Explain the operation of electrolytic cells
• Write simple half-reactions
• Solve Faraday's Law problems

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Day 1: 40 min

Day 2: 15 min

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• Protective glasses and lab coat/apron must be worn at all times.
• The 3 M sodium hydroxide is a strong base and is very dangerous. If spilled, the sodium hydroxide should be flushed with large amount of water and then neutralized with a dilute vinegar. For sodium hydroxide splashed into the eyes, the eyes must be washed in an eye wash for 15 min, and then a doctor should be consulted.
• Sulfuric acid, H₂SO₄, is a strong acid reagent. If spilled, sulfuric acid should be flushed with water, and then neutralized with sodium hydrogen carbonate, NaHCO₃. For sulfuric acid splashed into the eyes, wash in an eyewash for 15 min, then consult a doctor.

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Non-Consumables (per lab team)

• 1 Balance, 0.01 g precision
• 3 Beakers, 2 100 mL, 1 250 mL
• Emery paper or steel wool
• 1 Milliampere meter, DC
• 1 Power supply, variable, 12 V DC
• Wires for the circuit

Consumables (per lab team)

• 3 M H₂SO₄, 30 mL
• 3 M NaOH, 30 mL
• Plating solution, 200 mL *
• Distilled water

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Dissolve 250 g of CuSO₄•5H₂O in 500 mL distilled water, add 50 mL of 3 M sulfuric acid, H₂SO₄, and dilute to 1 L.

Preparation of 3 M sodium hydroxide, NaOH (for 12 stations):

While stirring and cooling, dissolve 120 g sodium hydroxide into 700 mL of distilled water. Dilute to one liter when cool.

Preparation of 3 M sulfuric acid, H₂SO₄ (for 12 stations):

While stirring and cooling, slowly add 160 mL of concentrated sulfuric acid, H₂SO₄, to 700 mL of distilled water. Dilute to one liter when cool.

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1. Students may work individually, in pairs, or in groups of three depending on the available apparatus. The activity uses considerable bench space. The activity may take one or more lab periods depending on the ability of the students. Students may set up the apparatus with help from the instructor.
2. Current Control: The current must be low enough so an even-adhering deposit of copper metal is made on the cathode. A current of 0.25 A is recommended. Students should try to maintain a constant current. The ammeter is the major limiting factor to the accuracy of this experiment. A digital voltmeter with at least 1 A DC range is recommended if this activity is done with current and time measurements.
3. You will need at least one hour to prepare the solutions and assemble and check out the procedure. The solutions are stable and can be used indefinitely if care is taken to prevent contamination.
4. Not all keys are suitable for this activity. Before use, test the key by adding a drop of copper plating solution to the key. If the key reacts with the plating solution, then another key must be found.

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Conducting the Laboratory Activity

Pre-lab Discussion

1. Review the operation of electrolytic cells.
2. Emphasize the differences between voltaic and electrolytical cells.
3. Review Faraday's Laws.
4. Review the proper wiring of the plating cell.

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1. Monitor the cleaning of the key and copper electrode and the wiring of the electrolytic cell.
2. The teacher should check the wiring setup. Initial the proper space on the student activity sheet.

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The mass lost or gained at the electrodes will be around 0.120 g for a 0.25 A current over 30 min.

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1. Copper was deposited on the key. The mass of the copper electrode decreased.
2. The reactions take place at the surface of each electrode.
3. 3. The changes in masses are about the same. Differences may be accounted for by copper flaking off the key.
4. 1. The number of coulombs (C) transferred by the power supply

      (0.250 C/s) x (1500 s ) = 375 C

   2. The number of faradays (F) transferred (375 C) x (1 F/ 96500 C)= 0.00389 F
   3. The number of moles of electrons transferred

      (0.00389 F) x (1 mol e-/F) = 0.00389 mol e-

5. 1. Moles of copper ions deposited on the key (negative electrode) =

      (0.118 g Cu) x (1 mol Cu ions/ 63.5 g Cu) = 0.00186 mol Cu ions

   2. Moles of copper atoms lost by the copper strip (positive electrode) =

      0.122 g Cu x 1 mol Cu atoms/63.5 g Cu = 0.00192 mol Cu

6. Ratio of moles of electrons transferred to moles of copper ions deposited on the key = 0.00389 mol e-/ 0.00186 mol Cu ions = 2.09/1.00
7. Copper ions; the charge on the copper is ²⁺.
8. Write the reduction half-reaction taking place at the cathode:

   Cu²⁺(aq) + 2e^- --> Cu(s)

9. Ratio of moles of electrons transferred to moles of copper atoms lost by the copper strip =

   0.00389 mol e-/ 0.00192 mol Cu atoms = 2.03/1.00

10. Write the oxidation half-reaction taking place at the anode (copper strip). Cu(s) --> Cu2+(aq) + 2e^-
11. Bonus (.00186 mol Cu²⁺/1500 sec)x(6.02 x 1023 Cu²⁺ ions/1 mol) = 4.01 x 1020 Cu²⁺/sec

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Have students post their data to compare the mass of copper gained by the key with the mass lost by the copper anode. Discuss the results. Ask for evidence of chemical changes taking place at the electrodes. Discuss why some keys are not suitable for this activity. Use an overhead to project a diagram of the circuit and electrolytic cell. Discuss the reactions taking place at each electrode.

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1. Set up a small industrial electroplating tank to make jewelry. Commercial electroplaters may be willing to sell the brighteners and used phosphorized copper anodes required for bright acid copper plating. The references below will provide the needed information for construction of a small industrial plating tank to make jewelry by electroplating or electroforming.
2. Visit an electroplating plant that produces chrome and nickel plate. After talking to the technical personnel, design an electroplating tank for nickel plating.
3. This activity deposits one metal on another metal. Plastics may be electroplated with proper conditioning with copper, nickel and chromium to form car grills and knobs for your stereo. Investigate this process and give a report to the class.

A commercial company manufactures a small electroplating tank for art and chemistry teachers. For detailed directions and information write to Hach Company, P.O. Box 389, Loveland, CO 80539.

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Assessing the Laboratory Learning

Laboratory Practical

1. Bring a quarter or other suitable metallic object to class.
2. Design an experiment to copper plate the metal object. Your teacher will provide a copper electroplating station and a copy of Fig. 1 from the student activity. Complete your experiment.
3. The lab report will include a procedure, observations, redox half-reactions, and a polished copper-plated object.

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Choose from among the following questions.

1. What is the source of copper deposited on the key?
2. Write the reduction half-equation for this reaction.
3. What is the source of electrons?
4. What is the purpose of the power supply?
5. 1. What ions are forming at the anode (the site of oxidation)?
   2. Write an oxidation half-reaction for this reaction.
6. If you continued the electrolysis for a longer period of time, you would observe no change in the intensity of the blue solution. Explain your observations in terms of the reactions taking place at the electrodes in the cell.
7. Explain the maxim "you cannot get something for nothing" in terms of what you have observed during this laboratory activity.
8. A current of 0.250 A deposits 0.237 g of a metal in 1808 s.
   1. What mass of metal will be deposited by 1 faraday?
   2. If the metal has an molar mass of 152 g/mol what is the charge on the metal ion?
   3. Write the reduction half-reaction for this reaction using Ae as the chemical symbol for the metal.

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1. The copper ions in the plating solution.
2. 2 e^- + Cu²⁺(aq) --> Cu(s)
3. The source of electrons is the oxidation reaction taken place at the anode.
4. The power supply serves as an electron pump.
5. 1. Copper ions, Cu²⁺(aq), are forming at the anode.
   2. Cu(s) --> Cu²⁺(aq) + 2 e^-
6. The blue color is due to the presence of Cu²⁺(aq). Since the color of the solution did not change, the copper electrode released one hydrated copper ion for each hydrated copper ion deposited on the key.
7. The Law of Conservation of Matter is obeyed. The mass lost by copper strip was deposited on the key.
8. 1. (1808 s)(.250 A) = 452 C

      (.237 g/ 452 C) x (96,5000 C/ F) = 50.6 g/ F

   2. (152 g metal/mol)(1 mol e-/50.6 g metal) = 3.00 mol e-/mol metal. Therefore charge on metal ion = 3.00
   3. 3 e^- + Ae³⁺(aq) --> Ae(s)

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