Galvanic Cells

Source: Smaa Koraym at Johns Hopkins University, MD, USA

1. Constructing a Lead-Copper Galvanic CellExpand

   In this experiment, you will construct a lead-copper galvanic cell and then measure the voltage that is generated during the redox reactions occurring at the anode and cathode.

   • To begin, put on the necessary personal protective equipment, including a lab coat, gloves, and chemical splash goggles.
   • Prepare 100 mL of a 0.05 M copper sulfate solution. First, calculate the amount of copper sulfate that you will need to weigh out using the molar mass of copper sulfate pentahydrate.
   • Measure the amount you calculated and transfer it to a 100-mL volumetric flask.
   • Measure 50 mL of deionized water, pour the water into the flask, and seal the flask with paraffin film. Swirl the flask as many times as needed until the copper sulfate dissolves.
   • Measure another 50 mL of deionized water and fill the flask until the water reaches the 100-mL measurement line. Cover the flask with paraffin film and invert and swirl the flask until the solution appears homogeneous. Label the flask with the exact concentration of copper sulfate created.
   • Prepare a 20-mL solution of 0.05 M lead nitrate. Calculate the amount of lead nitrate you need, then weigh it out and add it to a 20-mL volumetric flask.
   • Measure 10 mL of deionized water and add it to the lead nitrate. Seal the flask and mix the solution until the lead nitrate is completely dissolved.
   • Measure another 10 mL of water and transfer it to the flask until the level of the solution reaches the fill line. Seal the flask and mix the solution thoroughly. Label the solution with the exact concentration prepared.
   • Now, set up the galvanic cell. First, retrieve two pieces of copper sheet, a lead strip, and a piece of emery paper.
   • Take your 6-well reaction plate and pour some of the copper sulfate solution into one well, filling it 3/4 full. Fill an adjacent well with lead nitrate solution. Be sure to label the two wells.
   • Label your 600-mL beaker as ‘waste’.
   • Use emery paper to polish the copper and lead strips. Rinse each strip with deionized water after polishing.
   • Take your multimeter and connect the red and black leads to the appropriate plus and minus connections. Use alligator clips to connect the copper strip to one lead and the lead (Pb) strip to the other lead.
   • Place the copper strip inside the copper sulfate solution and place the lead (Pb) strip inside the lead nitrate solution. Adjust both electrodes so that half of the electrode is in the solution, bending them if necessary to keep them in place.
   • Obtain a string that has been soaking in the potassium nitrate solution.
   • Use forceps to place one end of the string in the copper sulfate and the other end in the lead nitrate, making sure it does not come in contact with either electrode. This will act as the salt bridge.
   • Turn on your multimeter to the voltage setting, and check the reading. If it is negative, you need to reverse the leads attached to the electrodes. Record the voltage reading in your lab notebook.

     Table 1: Lead-copper Galvanic Cell

     ΔE measured (volts)
     Anode
     Cathode

     Click Here to download Table 1
   • Use a glass thermometer to measure the temperature of one of the wells and record it in your lab notebook.
   • When you're finished, remove the electrodes from the wells, pour the solutions into the waste beaker, and throw the salt bridge in the trash. Rinse the copper and lead strips with deionized water, and dry them with a paper towel.
2. Galvanic Cell with Unknown ElectrodesExpand

   In this part of the experiment, you will construct a galvanic cell with an unknown electrode and its corresponding solution. You'll then use the reduction potential to identify the electrode material. Your instructor will assign you the unknown that you will test.

   • Select the corresponding unknown electrode and use your 10-mL graduated cylinder to retrieve 8 mL of the unknown solution.
   • Pour the unknown solution into one of the unused wells. Label the well containing the unknown solution and pour your copper sulfate solution into the adjacent well.
   • Remove the lead (Pb) strip from the multimeter connection, and replace it with your unknown strip.
   • Leave the copper strip connected. Place the electrodes in their corresponding wells, bending them if necessary to keep them in place.
   • Retrieve another salt bridge and place one end in each well like before, making sure that it is not touching either electrode.
   • Turn on the multimeter and record the voltage reading.

     Table 2: Unknown Galvanic Cells

     Metal 1	Solution 1	Metal 2	Solution 2	ΔE measured (volts)	Eº copper (volts)	Eº unknown (volts)	ΔEº standard conditions (volts)	Temperature (K)	n	[Red] (M)	[Ox] (M)
     Cu	Copper sulfate	Unknown #1	Unknown #1
     Cu	Copper sulfate	Unknown #2	Unknown #2

     Click Here to download Table 2
   • Measure the temperature of the unknown well using the glass thermometer and record the value.
   • When finished, remove the electrodes from the wells and rinse them with water. Dry them with a paper towel and set them aside.
   • Pour the solutions into the waste container. If the instructor has assigned you a second unknown, perform the experiment again exactly the same way.
3. Creating a Concentration CellExpand

   In the next experiment, you will construct a copper concentration cell. A concentration cell is a galvanic cell where the two electrodes are the same material. One half-cell contains a concentrated solution, while the other contains a dilute solution. As oxidation occurs in the dilute cell, ions from the electrode enter the solution making it more concentrated.

   Reduction occurs in the concentrated cell as metal ions are reduced and plated onto the electrode, thereby lowering the concentration of the solution. The two half-reactions drive a voltage difference between the two electrodes until equilibrium is reached.

   You will construct a concentration cell using copper electrodes, concentrated copper sulfate, and dilute copper oxalate. You'll measure the voltage difference between the two cells caused by the difference in copper concentration. Because copper oxalate is only slightly soluble, you'll then use the voltage measurement to determine the true concentration of copper in the copper oxalate solution.

   • To begin, make copper oxalate. Obtain 0.1 M oxalic acid from your instructor, then label a 100-mL beaker with the chemical formula for copper oxalate.
   • Measure 50 mL of the oxalic acid and pour it into the beaker. Use the volumetric pipette to measure 2 mL of the 0.05 M copper sulfate solution and dispense it into the oxalic acid.
   • Use a glass stirring rod to stir the solution for about 5 min. This gives the reaction enough time to reach equilibrium.
   • Now, label a clean 100-mL beaker as ‘copper sulfate’ and pour the 0.05 M solution into it. Place the copper sulfate beaker and the copper oxalate beaker close together.
   • Obtain 0.1 M potassium nitrate in a labeled 100-mL beaker, and immerse a pre-cut piece of filter paper into it, letting it soak for 20 s. This salt-soaked filter paper will act as a salt bridge.
   • Set the filter paper across the beakers so that it connects the two half-cells.
   • Measure the temperature of the solution in one of the beakers and record it in your notebook.
   • Use the emery paper to polish your two copper strips. Rinse them both with water after polishing.
   • Connect the copper strips to the leads of the multimeter and place one electrode in each beaker.
   • Turn on the multimeter and measure the voltage. Record this value in your lab notebook.

     Table 3: Copper concentration cell

     ΔE measured (volts)
     Temperature (K)
     [Cu²+] (M)
     R (J/mol·K)	8.314
     F (C/mol)	96,485

     Click Here to download Table 3
   • When you are finished, remove the electrodes from the beaker and throw the filter paper away. Disconnect the electrodes and rinse them with deionized water. Then, dry them with a paper towel.
   • Finally, pour the solutions into your waste beaker. Neutralize the liquid waste with baking soda, swirling the solution and adding more baking soda until the solution stops bubbling. Pour the neutralized liquid into the heavy metal waste container provided by your instructor.
   • Rinse all glassware with water and return all equipment, including the electrodes, to your instructor.
4. ResultsExpand
   • First, determine the half-reactions happening at each electrode. Since the oxidation and reduction reactions in galvanic cells are spontaneous, the reaction with the highest reduction potential will be the reduction reaction. In the case of the lead-copper cell, the copper half-reaction has the higher reduction potential; thus, copper is reduced in our galvanic cell.
   • In turn, oxidation occurs at the lead electrode. Thus, the half-reaction occurring at the lead electrode is an oxidation reaction. Oxidation always occurs at the anode; therefore, the lead electrode is the anode. Reduction always occurs at the cathode; thus, the copper electrode is the cathode. Your multimeter reading should equal the voltage difference between the two half-reactions, 0.463 volts.
   • Determine the direction of electron flow. Since electrons always flow from the anode to the cathode, electrons will flow from the lead electrode to the copper electrode.
   • Next, look at your first unknown galvanic cell. Remember that the multimeter measured the difference in potentials between the two half-reactions, which corresponds to the voltage you measured.
   • We know that one electrode is copper. So, if our voltage difference is 1.057 volts, the reduction potential of the other reaction is negative 0.720 volts. This corresponds to the zinc half-reaction. Therefore, the unknown electrode is zinc and the unknown solution is zinc sulfate. The reduction potential of copper is higher than that of zinc, so the copper cell is the reduction cell and the zinc cell is the oxidation cell.
   • Calculate the concentration of the zinc sulfate solution using the Nernst equation. The Nernst equation relates the measured voltage, E, to the reduction potential expected at standard conditions, E naught. R is the universal gas constant, T is the temperature, F is Faraday's constant, n is the number of electrons transferred in the half-reaction, and Q_h is the reaction quotient.
   • The reaction quotient is equal to the concentration of the chemical species being oxidized divided by the concentration of the species being reduced. For our galvanic cell, copper sulfate is reduced and zinc sulfate is oxidized. Since we know the concentration of copper sulfate, 0.05 M, we can calculate the concentration of the oxidation cell, which was zinc sulfate in our experiment.
   • Finally, look at the copper concentration cell. Here we created a galvanic cell with half-cell containing dilute copper oxalate and the other containing copper sulfate. Since the copper concentration in the dilute half-cell changes as the experiment proceeds, we can calculate the concentration of dissolved copper ions using the Nernst equation.
   • The two electrodes are the same in a concentration cell, so the standard electrode potential is 0. The potential difference is the voltage that we measured, and it is due to the difference in copper concentration between the two cells. After plugging in the known values, we can then solve for the concentration of the copper ion in the dilute cell. Here, the concentration of copper ions was 7.9 x 10^-7 M.