THE CHEMISTRY OF COPPER PLATING
This experiment demonstrates the process of electroplating and a commercial method used to purify copper.
This experiment is most appropriate for a first-year college prep or AP class if done quantitatively. If done qualitatively, it would be appropriate for a general class. One of the most important applications of electrolytic cells is the process of electroplating, in which a thin layer of metal is deposited on an electrically conducting surface. In electroplating, the metal to be plated is used as the anode and the electrolytic solution contains an ion derived from that metal. In this experiment, a copper anode (US penny) will be used in a solution of copper sulfate. Copper will be plated out onto a second penny at the cathode.
30 minutes to set up; one period to complete.
Sulfuric acid can cause severe burns; handle with care. Goggles must be worn throughout this experiment. Although the power source is relatively weak, the electrodes and connecting wires should not be handled when the cell is operating. If a 9-V battery is used as the power source, it will become quite hot during use; caution should be exercised.
- electrolyte solution (200 g CuSO4· 5H2O + 25.0 mL concentrated H2SO4 solution in enough distilled or deionized water to make l.00 L of solution)*
- pre-1983 pennies
- power supply (6.0-9.0 volts, 0.60-1.0 amps)*
- connecting wires with alligator clips
- 16-18 gauge copper wire
- 250-mL beaker*
- cardboard square (approx. 15 cm on a side)
- ammeter (optional)
- Copper(II) sulfate pentahydrate is available from garden supply stores as root eater.
- Sulfuric acid is available from auto supply stores as battery acid. Substitute 95 mL of battery acid for 25 mL of concentrated sulfuric acid.
- A battery charger or 9-V battery may be substituted for the power supply. If a 9-V battery is used, it will be nearly "dead" after completing the experiment.
- A 16-oz plastic glass may be substituted for the beaker.
Pennies used in this experiment should not be reused as currency. Solids may be discarded with other solid waste. Electrolyte solution should be stored for re-use; if it becomes necessary to dispose of the solution, it should be flushed down the drain with plenty of water.
As copper is plated out at the cathode (negative electrode), copper goes into solution at the anode (positive electrode) as copper(II) ions, maintaining a constant concentration of copper(II) ions in the electrolytic solution.
- Pour 200 mL of the electrolyte solution into the beaker.
- Attach connecting wires with alligator clips to the terminals of the power supply.
- Clean the pennies with a mixture of 3 g NaCl and 15 mL vinegar; rinse and dry.
- Tightly wrap one end of a 10-cm length of copper around each penny, leaving 5-6 cm of wire free.
- Mass each penny-copper wire assembly and record the masses.
- Push the free end of each wire through the cardboard square and place the square over the beaker so that the penny "electrodes" are immersed in the electrolyte solution as illustrated below. Note: the two electrode assemblies must not touch.
- Attach the connecting wires to the top of the copper wire assemblies.
- Allow the electroplating cell to operate for 30-60 minutes. Record the exact time the cell was operating (optional).
- Record the ammeter reading from the meter or power supply (optional).
- Remove each "electrode;" dry, being careful not to lose any of the copper plating; mass each and record.
- If an ammeter reading is taken, calculate:
- the number of coulombs of charge passed through the electrolytic cell,
- the theoretical number of moles of copper that should have plated out,
- the actual number of moles of copper that plated out, and
- the % yield of copper.
- If an ammeter is not used, calculate:
- the number of moles of copper removed from the anode,
- the number of moles of copper added to the cathode,
- the % of copper conserved,
- the number of coulombs of current necessary to plate out the moles of copper calculated in step 12b, and
- the average current that must have passed through the cell.
cathode: Cu2+(aq) + 2 e- Cu(s)
Commercial plating is done very slowly in order to obtain a smooth, even coating of the plated metal. Although this experiment does not produce plating of commercial quality, it gives students the opportunity to study the chemistry of an important commercial process. This general method is also used in purifying copper. A small cathode of pure copper is used with a larger anode of impure copper. As the electrolytic cell operates, pure copper is transferred to the cathode.
anode: Cu(s) Cu2+(aq) + 2 e-
Students should be introduced to Faraday's law before doing this experiment. From this law, students will note that 2 x 96,485 coulombs of charge are required to produce one mole of copper from copper(II) ion. If an ammeter reading is taken, the number of coulombs that actually passed through the electrolytic cell can be calculated by using the formula; q = It, where q is the charge in coulombs, I is the current in amperes, and t is the time in seconds. From the coulombs of charge that pass through the cell, students can calculate the theoretical number of moles of copper that should have plated out and compare this to the actual number of moles that were plated out. If one assumes that the theoretical yield is equal to the actual yield, the atomic mass of copper can be calculated.
If an ammeter reading is not taken, students can compare the changes in mass of the two electrodes and from the number of moles of copper plated, calculate the number of coulombs of charge passed through the cell and the average current through the cell.
Masterton, W.L., Slowinski, E.J., and Stanitski, C.L., Chemical Principles, Saunders College Publishing, 1985, p. 701.
- If power supplies or batteries are not available in sufficient supply to allow students to do this as an experiment, it may be done as a demonstration. In such a case, it may be desirable to carry out the electrolysis in a 1-L beaker or large, wide-mouth jar and use small pieces of copper pipe or small copper plumbing fittings (available from a hardware store) to make the demonstration more visible to students. The reaction, then, should be run longer to get a larger mass change in the electrodes; this would require a power supply rather than a 9-V battery.
- Students might be encouraged to try plating copper onto other metals.
- This work describes the chemistry of electroplating and specific commercial electroplating solutions.
Submitted by Joe Baron, Elna Clevenger, Carolyn Lucas, Patti Ruff, and Bill Vitori
Woodrow Wilson Leadership Program in Chemistry
The Woodrow Wilson National Fellowship Foundation
CN 5281, Princeton NJ 08543-5281