Electrochemical reactions are vital to many processes in technology and the human body and involve the transfer of electrons from one chemical species to another. These reactions are called reduction-oxidation reactions or redox reactions.
When a species loses electrons, it undergoes oxidation, and we say that it is oxidized. When a species gains electrons, it has undergone reduction, and we say that it is reduced. In an electrochemical galvanic cell, redox reactions occur spontaneously, which generate electrical energy. However, in an electrolytic cell, electrical energy is applied, which causes the redox reactions to occur.
An electrolytic cell consists of the reaction solution called the electrolyte, which contains any chemical species that are oxidized or reduced as well as ions needed to enable electron flow. The electrolytic cell has two metal electrodes that are immersed in the electrolytic solution. An external circuit connects the two electrodes, which completes the circuit and enables us to apply voltage or current.
The applied current or voltage is what causes the electrochemical reaction to take place. The oxidation half-reaction occurs at the anode, and the reduction half-reaction occurs at the cathode. To help you differentiate between the two, remember the phrase 'red cat,' which stands for 'reduction at cathode'.
But how do we determine which electrode is which? We can use the standard reduction potential of the metal electrodes, which is a measure of the metal's tendency to lose electrons. The higher the reduction potential, the more likely it is for the metal to be reduced.
Let's say we place a silver electrode and a copper electrode in an electrolytic cell and apply current. The silver electrode has a higher reduction potential, so it is reduced and acts as the cathode. The copper electrode has a lower reduction potential, so it is oxidized and acts as the anode.
One use of electrolytic cells is to perform electroplating, which is a reaction where one metal is oxidized and then reduced onto the surface of another metal. Since the anode is the metal doing the plating, in our example of silver and copper, the silver electrode is plated with a thin layer of copper.
In this lab, you'll assemble an electrolytic cell and perform the electroplating process by plating a brass key with copper.
Source: Smaa Koraym at Johns Hopkins University, MD, USA
- Electrolytic Cells
In this experiment, you will be assembling an electrolytic cell to perform electroplating. An electrolytic cell contains an anode and cathode connected by a power source, which drives the nonspontaneous redox reaction in the cell. When the anode material is oxidized, the ions travel to the cathode, where they are reduced and plated.
In the case of this experiment, the electrolytic cell contains a copper anode and a brass key for the cathode. Oxidation of the anode creates copper(2+) ions, which are then reduced on the brass key, forming a thin plating of solid copper.
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Put on gloves, safety goggles, and a lab apron, which must be worn at all times.
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Obtain one 10-ohm resistor, one copper sheet electrode, one copper wire, and one piece of emery paper.
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Use the emery paper to polish the brass key and the copper sheet. Rinse the key and the copper electrode with water. The brass key will be the cathode, and the copper sheet will be the anode.
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Take one 50-mL beaker and a watch glass to the instructor's hood and obtain approximately 20 mL of acetone. Place the watch glass on top of the beaker to prevent evaporation and bring it back to your bench.
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Label the other 50-mL beaker as 'organic waste'.
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Use a disposable pipette to rinse the copper electrode and key with acetone. Collect excess in the waste beaker. Place the key and copper electrode on a paper towel to dry.
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Measure the initial masses of the key and copper electrode and record their values in your lab notebook.
Table 1: Moles of Cu Transferred
Click Here to download Table 1Mass of key (g) Change in mass (g) Moles Cu plated Mass of Cu electrode (g) Change in mass (g) Moles Cu lost 0 min 5 min 1st 30 min 2nd 30 min -
Set up the electrolytic cell. Use one electrical wire with alligator clips to connect the negative terminal on the battery to the current probe. Wrap the 10-ohm resistor around the other end of the current probe.
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Use another wire to connect the positive terminal on the battery to the copper electrode.
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Wrap the copper wire around the brass key and use an electrical wire to connect the resistor on the current probe to the copper wire on the brass key.
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Obtain about 200 mL of concentrated copper sulfate in a 250-mL beaker. The amount does not need to be exact. Bring the beaker back to your bench and set it on the stir plate. Carefully, place a magnetic stir bar in the beaker and turn the stir setting on to high.
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Connect the data acquisition system to the current probe and allow the data acquisition system to sense the probe. It should return a reading in amps. Change the rate to two scans per min and the duration to 5 min.
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Place the brass key in the beaker. Hang it off the side of the beaker so that the key is almost completely submerged but does not touch the magnetic stir bar. Note: The copper wire should not be in the solution.
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Place the copper electrode in the beaker, bending it so that it hangs over the edge of the beaker. Ensure that it is not touching the stir bar or the brass key.
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Start the scan and let it collect data for 5 min. This will act as a test run to make sure plating occurs on the brass key.
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After the scan ends, remove the key and copper electrode. The copper plating on the brass key should be evident. If it is not visible, check your circuit to make sure that the battery is connected in the correct orientation.
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Once you have completed a successful 5-min plating, disconnect the key and copper electrode from the wiring. Remove the copper wire from the key and rinse the key and copper electrode with water and then acetone. Lay them on a paper towel to dry.
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Measure the masses of the key and the copper electrode and record them in your lab notebook.
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Reconnect the key and copper electrode and place them back in the solution like before. Restart the run on the data acquisition system, making sure that the rate is set to two scans per min. With the duration set to 30 min, allow the electroplating process to run.
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After the scan is complete, remove the electrode and key from the solution, disconnect them from the wires, and rinse them with water and then acetone. Allow them to dry on a paper towel, then weigh them and record their masses in your notebook.
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Reassemble the setup and repeat the electroplating for a second 30-min trial. When the scan is complete, remove the key and copper electrode, rinse them like before, and record their mass.
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Save the recorded data as a text file on a flash drive.
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To clean up, turn the stir setting off and remove the stir bar from the solution using forceps. Dispose of the acetone in the organic waste container provided by your instructor. The copper sulfate solution can be reused for future labs, so pour the concentrated copper sulfate back into the carboy.
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Rinse all of your glassware with water and return it to the instructor.
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- Results
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First, calculate the total charge Q that flowed through the system during the different trials. One coulomb is the quantity of electricity flowing through a wire carrying 1 ampere of current in 1 s, thus Q = I x t.
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Review the current probe data for the first 5-min electroplating trial. Note: If the current dropped to 0 or close to 0 at the end of the trial, then the electrons stopped flowing from the copper electrode to the key. Only use the values recorded before the drop in current occurred. If the current drop to 0 early in the plating process, repeat the trial.
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The current varies over time, so calculate the average current, I.
Table 2: Moles of Electrons Transferred
Click Here to download Table 2Trial Time (s) Currentaverage (A) Total charge (C) Moles (e?) 5 min 1st 30 min 2nd 30 min Total -
Use the average current over the total time to determine the total charge transferred during electroplating, Q. About 20 coulombs were transferred during the 5-min trial, and between 110 and 120 coulombs were transferred during each 30-min trial. The total charge transferred over the entire 65-min electroplating process was about 250 coulombs.
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Use the total charge transferred to calculate the moles of electrons plated during each trial. The moles of electrons plated equals the total charge transferred divided by the Faraday constant. The Faraday constant represents the magnitude of electric charge per mole of electrons and is determined by multiplying the charge of an electron by Avogadro's constant.
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Determine the moles of electrons that were transferred for the 5-min and 30-min trials.
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Calculate the moles of copper that were plated on the key after each trial. Similarly, determine the moles of copper lost from the copper sheet electrode.
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Compare the moles of copper that were lost from the electrode to the amount that was plated to the brass key. The numbers should be close.
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Check to see if these values match what is expected based on the number of moles of electrons that were transferred. The reduction of 1 mole of copper(2+) ions requires 2 moles of electrons to produce 1 mole of solid copper.
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