Materials Included In Kit

Additional Materials Required

Prelab Preparation

Copper Plating Solution—Approximately 200 mL of copper plating solution is needed per group. To prepare 1 liter of copper plating solution: dissolve 250 grams of cupric sulfate in 600 mL of distilled or deionized water, add 50 mL of 3 M sulfuric acid, and dilute to 1 L with distilled or deionized water. Note: Enough cupric sulfate (750 g) and sulfuric acid solution (150 mL) are provided to make 3 L of copper plating solution (enough for 15 groups of students).

Copper Wire—Use wire cutters to cut the 3-foot length of copper wire into pieces that are approximately 5 cm long. Each group will need one piece.

Power Supply and Connecting Wires—A low voltage DC variable power supply, such as Catalog No. AP9279, or a similar unit will work. Be sure to use insulated connecting wires with alligator clips or banana plugs. Practice the electrolysis setup ahead of time to be sure the power supply, ammeter and wires are compatible.

Safety Precautions

Sodium hydroxide solution and sulfuric acid solution are corrosive to eyes, skin and other tissues. Copper plating solution is an acidic solution of cupric sulfate; it is moderately toxic by ingestion and inhalation and is a skin and respiratory irritant. Avoid skin contact with all chemicals. Do not operate a power supply with wet hands or in wet areas. Be sure the area is dry before turning on the power supply or closing the circuit. Follow additional safety precautions as appropriate to your power supply. 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. Excess copper plating solution may be saved for reuse or neutralized with base and then disposed of according to Flinn Suggested Disposal Method #26b.

Teacher Tips

• Enough materials are provided in this kit for 30 students working in pairs or for 15 groups of students. This laboratory activity can reasonably be completed in one 50-minute class period. The post-lab questions and calculations can be assigned for out-of-class work, or can be completed and discussed on Day 2 of the lab.
• Depending on equipment availability (power supply and ammeter), you may have students work in pairs or in larger groups of 3 or 4. Considerable bench space is needed, so larger groups will work fine.
• The keys provided in the kit are suitable for electroplating. Not all keys are suitable for this activity. If you wish to have students electroplate their own keys, a piece of jewelry, or some other chosen piece of metal, test the key or metal by adding a drop of copper plating solution onto the key or metal or by dipping the metal into the plating solution. If the key or metal reacts spontaneously with the plating solution (with no current added), then another key or metal must be found. Remember electroplating is not a spontaneous process.
• If a power supply with a built-in ammeter is available, the need for an external ammeter is eliminated.
• The Calculations and Post-Lab Questions have been written as a higher level, quantitative activity. The lab is also suitable, however, for general chemistry students as a qualitative introduction to the applications of electrochemistry.

Further Extensions

• Discuss what would happen if the wiring were reversed from the start. Or what would happen if the wiring were reversed after the initial run.
• Can a thin layer of metal be electroplated onto a surface other than a metal?

Answers to Prelab Questions

1. A thin layer of copper will deposit on the key during the electroplating process. What is the source of this copper?
   Copper ions from the electrode go into the solution; copper ions from the solution plate onto the key.
2. What is the purpose of the power supply in the electroplating reaction?
   The power supply serves as an electron pump.
3. Write the reduction half-reaction for the electroplating reaction. Where will this reaction occur?
   Cu²⁺(aq) + 2e^– → Cu(s)
   Reduction will occur at the cathode; that is, on the key.
4. What ions will form at the anode? Write the oxidation half-reaction for the electroplating reaction. Where will this reaction occur? Copper ions (Cu²⁺) form at the anode.
   Cu(s) → Cu²⁺(aq) + 2e^–
   Oxidation will occur at the anode; that is, on the copper electrode.
5. Explain the relevance of the Law of Conservation of Matter to this lab activity.
   The Law of Conservation of Matter states that matter cannot be created or destroyed. Any mass lost by the copper strip goes into solution. Copper ions from the solution deposit on the key.

Sample Data

Answers to Questions

Note: Sample calculations are shown using Trial I data.

1. What observations suggest that a chemical reaction occurred?
   The mass of the copper electrode decreased. Copper deposited onto the key.
2. Calculate the change in mass of the key and of the copper electrode. Record in the Data Table. Compare the change in mass of the two electrodes. Explain your observations.
   Change in mass of the key = 8.27 g – 8.14 g = +0.13 g
   Change in mass of the Cu electrode = 4.39 g – 4.53 g = –0.14 g
   The changes in masses are nearly equal—the mass that the key gained is approximately equal to the mass that the Cu electrode lost. This observation follows the Law of Conservation of Mass; slight differences may be accounted for by copper flaking off the key.
3. Calculate the number of coulombs (C), the number of Faradays (F), and the number of moles of electrons (moles e^–) transferred by the power supply. Record in the Calculations Table.
   0.75 C/sec × 600 sec = 450 C
   450 C × 1 F/96,500 C = 0.00466 F
   0.00466 F × 1 mole e^–/F = 0.00466 mole e^– 
4. Use the mass of copper deposited on the key (negative electrode) to determine the moles of copper ions deposited at this electrode (the cathode). Record in the Calculations Table.
   Moles of Cu ions deposited on the key = (0.13 g Cu) × (1 mole Cu ions/63.5 g Cu) = 0.00205 moles Cu ions
5. Use the mass of copper lost by the copper strip (positive electrode) to determine the moles of copper atoms lost at this electrode (the anode). Record in the Calculations Table.
   Moles of Cu atoms lost by the copper strip = (0.14 g Cu) × (1 mole Cu atoms/63.5 g Cu) = 0.00221 moles Cu 
6. 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)? Record in the Calculations Table.
   Ratio of moles of e^– transferred to moles of Cu ions deposited on the key = 0.00466 mole e^–/0.00205 moles Cu ions = 2.27
7. Use the result from Question 6 to determine the charge on the copper ion.
   The charge on the copper ion, based on Question 6, is 2.
8. Write the reduction half-reaction that occurred at the cathode. Relate this reaction to your answer for Question 6.
   Cu²⁺(aq) + 2e^– → Cu(s)
   The stoichiometric relationship shows that for every 2 moles of electrons transferred, there is one mole of copper formed.
9. What is the ratio between the moles of electrons transferred and the moles of copper atoms lost by the anode (copper strip)? Record in the Calculations Table.
   Ratio of moles of e– transferred to moles of Cu atoms lost by the copper strip = 0.00466 mole e^–/0.00221 moles Cu = 2.11
10. Use the result from Question 9 to write the oxidation half-reaction that occurred at the anode.
    Cu(s) → Cu²⁺(aq) + 2e^–
    The stoichiometric relationship shows that for every 2 moles of electrons transferred, one mole of copper ionizes.

 

Introduction

Coat your keys or another metal object with a thin layer of copper using an electrolytic process called electroplating. Electroplating is a technique used to deposit a layer of metal—such as copper, chromium, silver, gold or nickel—onto the surface of another metal. The deposit can provide both a protective and decorative coating for the metal which lies beneath it. Chrome plating (on car bumpers) or silver plating (on serving dishes) are common electroplating processes.

Concepts

• Electrolysis/electroplating
• Oxidation–reduction
• Faraday’s law

Background

Electrochemistry is the study of the interconversion of electrical and chemical energy. Two types of cell processes can result in the interconversion of these two energy sources—voltaic cells and electrolytic cells. Voltaic cells use a spontaneous chemical reaction to generate electrical energy. Electrolytic cells use electrical energy (e.g., from a battery) to make a nonspontaneous chemical reaction take place. A nonspontaneous process requires energy from an external source in order to drive the reaction to occur. The electroplating reaction performed in this lab activity is an example of an electrolytic nonspontaneous process.

The chemical reactions that take place in electrochemical cells are redox reactions. Reduction occurs at the cathode (negative [–] electrode) and oxidation occurs at the anode (positive [+] electrode). Within an electrochemical cell, positive ions (cations) move toward the cathode and negative ions (anions) move toward the anode.

Since the charge on an electron is 1.6022 x 10^–19 coulombs and there are 6.022 x 10²³ e^–/mole, then one mole of electrons carries a total charge of 96,485 coulombs (1 mole e^– = 96,485 coul or C). The constant 96,485 C/mol is called the Faraday constant in honor of Michael Faraday, who did extensive research on electrochemical cells. Michael Faraday observed that the amount of substance undergoing oxidation or reduction at each electrode during electrolysis is directly proportional to the amount of electricity that passes through the cell. This is known as Faraday’s Law of Electrolysis. The rate at which electrical charge moves through a circuit is most commonly measured in amperes (amps or A). One ampere of current equals one coulomb of charge passing a given point in one second, or 1 amp = C/sec. Hence, 1 Faraday = 6.022 x 10²³ e^– = 96,485 C.

In the electrolytic cells used in this lab activity, the nonspontaneous oxidation–reduction reactions occurring at the electrodes are the result of the transfer of electrons by the power supply. The power supply acts as an electron pump, pushing electrons into the cathode and removing them from the anode. Electrical neutrality, however, must be maintained; therefore, some process must consume electrons at the cathode and generate them at the anode.

The half-reaction occurring at the anode (the copper electrode) is the oxidation of copper to Cu²⁺. When copper is oxidized to copper(II) ions, two electrons are produced according to Equation 1.

The half-reaction occurring at the cathode (the metal key or object) is the reduction of copper(II) ions to copper. When copper(II) ions are reduced to copper metal, two electrons are consumed according to Equation 2.

When the combined oxidation–reduction reaction is carried out in an electrolytic (nonspontaneous) cell, the process is called electrolysis (see Figure 1). In this laboratory activity, a metallic object (a key) will be electroplated with a copper coating. The energy needed to perform the oxidation–reduction reaction will come from an external power source. The mass change at each electrode (the copper electrode and the metal key) will be investigated. The reactions occurring at each electrode will be observed and written. The charge on the copper ion will then be calculated by applying Faraday’s Law.

Materials

Safety Precautions

Sodium hydroxide solution and sulfuric acid solution are corrosive to eyes, skin and other tissues. Copper plating solution is an acidic solution consisting of cupric sulfate and sulfuric acid; it is moderately toxic by ingestion and inhalation and is a skin and respiratory irritant. Avoid skin contact with all chemicals. Do not operate a power supply with wet hands or in wet areas. Be sure the area is dry before turning on the power supply or closing the circuit. Follow additional safety precautions as appropriate to your power supply. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Part 1. Cleaning the Electrodes

1. Clean a key and a copper electrode by rubbing with fine steel wool. Wash the key and the copper electrode with detergent and rinse both with tap water.
2. Attach a 5-cm length of copper wire to the key. The wire will serve as a handle to remove it from the cleaning solutions and to support it during the electrolysis.
3. Soak the key and copper electrode for a few minutes in 30 mL of 3 M sodium hydroxide solution. Remove the key and rinse it with distilled or deionized water.
4. Soak the key and copper electrode for a few minutes in 30 mL of 3 M sulfuric acid solution. Remove the key and rinse it with distilled or deionized water.
5. Remove the wire from the key. Save the wire for use in step 9. Dry the key and the copper electrode thoroughly for Part 2.

Part 2. Copper Plating the Key

6. Determine the mass of the key to the nearest 0.01 g. Determine the mass of the copper electrode to the nearest 0.01 g. Record both initial masses in the Data Table.
7. Pour approximately 200 mL of copper plating solution into a 250-mL beaker. Note: Copper plating solution is an acidic solution consisting of cupric sulfate and sulfuric acid. Avoid all skin contact. Rinse skin immediately if any contact is made.
8. Place the copper electrode into the beaker. Bend the copper strip to fit over the top of the beaker (see Figure 2).
   {11950_Procedure_Figure_2}
9. Reattach the 5-cm length of wire to the key. Suspend the key in the beaker by wrapping the wire around a glass rod. Rest the rod on top of the beaker. Be sure the copper electrode and key are not touching.
10. Hook up the variable DC power supply. Wire the key directly to the negative (black) terminal of the power supply. Wire the positive (red) terminal of the power supply to the positive terminal of the ammeter. Wire the negative terminal of the ammeter to the copper electrode.
11. Note: Before turning on the power supply, be sure the area is dry. Read the Safety Precautions. Turn on the power supply and use a timer to begin timing. Adjust the current to approximately 0.50 amps (A). Monitor the current to maintain a constant reading throughout the electroplating. (Note: Current may vary as the copper is removed from the electrode and plated onto the key.) Record the exact current in A in the Data Table.
12. Turn off and unplug the power supply after approximately 10–15 minutes and stop the timer. Record the exact time for electrolysis in minutes in the Data Table. Convert the electrolysis time to seconds. Record the seconds in the Data Table.
13. Remove the key and copper electrode from the solution.
14. Carefully rinse the key with distilled water and gently pat it dry. Remove the wire from the key. Be careful to minimize the rubbing of the key or some copper may flake off and alter the final mass.
15. Determine the final mass of the key to the nearest 0.01 g. Record the final mass in the Data Table.
16. Rinse the copper electrode with distilled water, dry it off, and determine its final mass. Record the final mass of the copper electrode in the Data Table.
17. Disconnect the wires from the power supply and return the copper plating solution to the reagent table for reuse.
18. Complete the Calculations and Post-Lab Questions.
19. Optional Trial II—Repeat steps 1–18 using another metal object such as a coin or a nail.
20. Consult your instructor for appropriate disposal procedures.

Student Worksheet PDF

11950_Student1.pdf