Building Electrochemical Cells

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1 Cautions Heavy metals, such as lead, and solutions of heavy metals may be toxic and an irritant. Purpose To determine the cell potential (E cell ) for various voltaic cells and compare the data with the calculated E cell values. Introduction An electrochemical cell is a device that may be used for converting chemical energy into electrical energy. An oxidation-reduction reaction may serve as the basis for designing an electrochemical cell. The tendencies of metals to be oxidized differently are often used to design an electrochemical cell based on their reaction differences. Oxidation-reduction or redox reactions involve the transfer of electrons from one reactant to another. Oxidation and reduction half-reactions each represent half of the overall reaction. When adding the halfequations to yield the overall equation, the electrons appearing in the two half cell reactions must cancel each other. In any oxidation-reduction reaction equation, the number of electrons released must be equal to the number of electrons consumed. For example, the oxidation of zinc metal and reduction of Cu 2+ for a spontaneous electrochemical reaction results in the transfer of 2 electrons. Zn (s) Zn 2+ (aq) + 2e - oxidation half-reaction Cu 2+ (aq) + 2e - Cu (s) reduction half-reaction Zn (s) + Cu 2+ (aq) Zn 2+ (aq) + Cu (s) Overall reaction This reaction can be observed by placing a strip of zinc metal into a solution of copper(ii) nitrate, one may observe metallic copper being deposited on the zinc surface. With time the blue color characteristic of Cu 2+ ions in solution will fade and the amount of metallic copper will visibly increase. The strip of zinc metal is gradually being consumed during this process, indicating zinc atoms are being oxidized to zinc (II) ions. This observation shows that zinc is more easily oxidized than copper. Alternatively, the Cu 2+ ion can be described as being more likely to accept electrons from Zn metal than Zn 2+ ion is to accept electrons from Cu metal; a Cu 2+ ion is easier to reduce than Zn 2+ and thus has a greater reduction potential. The standard reduction potentials, E o red, for metal ions Reduction Half E o red, V provide a quantitative measure of a metal ion s tendency to Al 3+ (aq) + 3 e- Al (s) accept electrons and a selective list is shown in Table 1. Zn 2+ (aq) + 2 e- Zn (s) These potentials are determined under standard conditions; Fe 2+ (aq) + 2 e- Fe (s) M concentrations, 25 C, and 1 atm. The reference point for Sn 2+ (aq) + 2 e- Sn (s) these potentials is the reduction potential of hydrogen ions Pb 2+ (aq) + 2 e- Pb (s) (H + ) in aqueous solution being converted to hydrogen gas (H 2 ). 2 H + (aq) + 2 e- H 2 (g) 0.00 The greater more positive the reduction potential, the greater Cu 2+ (aq) + 2 e- Cu (s) ease of reduction. Ag + (aq) + e- Ag (s) The oxidation half-reactions of these metals in the table are simply the reverse of the half-reactions. The Table 1: Standard Reduction Potentials standard oxidation potential, E 0 oxid, of an oxidation half-reaction has the same numerical value, but the opposite sign of the corresponding standard reduction potential. The larger positive voltage indicates that a metal is a more reactive (or active) metal, which will donate electrons easily in an oxidation-reduction reaction. Revision F10 IB Page 1 of 8

2 We can use tabulated reduction potentials to make predictions about the spontaneity for certain oxidationreduction reactions. Let s predict what will happen if we place a Zn strip in a solution of Cu(NO 3 ) 2 under standard conditions. We need to determine if the Zn atoms will donate electrons to thecu 2+ ion, as part of a spontaneous oxidation-reduction reaction. The reduction half-reaction for Cu 2+ ion is described by the standard reduction potential for Cu 2+ at +0.34V. The oxidation of Zn is described by the standard oxidation potential (or the negative of the standard reduction potential), +0.76V. To obtain the net equation and overall cell voltage, we add the equations of the two half-reactions and cancel out the electrons. Zn (s) Zn 2+ (aq) + 2e - E o ox = 0.76 V Cu 2+ (aq) + 2e - Cu (s) E o red = 0.34 V Zn (s) + Cu 2+ (aq) Zn 2+ (aq) + Cu (s) E o cell = 1.10 V Likewise, the sum of the standard reduction potential and the standard oxidation potential of the half-reactions equals the overall cell voltage. If the standard net potential, E 0 cell, or sum of the standard reduction and oxidation potentials, is positive then the reaction will occur spontaneously. E o ox + E o red = E o cell In writing a balanced net cell reaction, we may need to multiply the coefficients of the half-cell reactions to properly cancel the electrons. However, changing the coefficients of a half-cell reaction does not influence the reduction and oxidation potential. Reduction and oxidation potentials are intensive properties in that they are independent of the amount of substance. A voltaic cell, also known as a Galvanic cell, is an electrochemical cell where a spontaneous reaction generates an electrical current. A picture of a typical electrochemical cell is shown in Figure 1. A voltaic cell is comprised of two connected half-cells, one containing the anode and the other the cathode. The connection allows a path for electrons to flow from one metal electrode to another through an external circuit and an internal cell connection (or salt bridge). Oxidation, or loss of electrons, occurs at the anode, while reduction, or gain of electrons, occurs at the cathode. In this experiment you will build various electrochemical cells using Cu, Zn, and Pb and measure the voltages. Based upon your observed voltages, you will be able to rank the metals in order of their relative ease of oxidation and compare the measure cell voltages with those calculated from the standard reduction potentials. Figure 1: An Electrochemical Cell ( Revision F10 IB Page 2 of 8

3 Procedure Note: Metal strips should appear bright and shiny before beginning. Look carefully for changes either in the appearance of the metal surfaces or the solutions to indicate a reaction. A. The Zn-Cu Redox System A1. Clean the Zn metal strip with sandpaper. A2. Place the clean metal in a small test tube. A3. Fill the test tube with 0.1M CuSO 4 solution is sure that the Zn metal strip is completely submerged. A4. After 2.5 minutes, record your observations. A5. After 5 minutes, record your observations. A6. Clean the Zn metal strip with sandpaper and dry using paper towels. A7. Retain the CuSO 4 solution for use in Part B. B. The Pb-Cu Redox System B1. Clean the Pb metal strip with sandpaper. B2. Place the clean metal in a small test tube. B3. Fill the test tube with 0.1M CuSO 4 solution be sure that the Pb metal strip is completely submerged. B4. After 2.5 minutes, record your observations. B5. After 5 minutes, record your observations. B6. Clean the Pb metal strip with sandpaper and dry using paper towels and return to the container B7. Pour the CuSO 4 solution into the container labeled Discarded CuSO 4 Solution. C. The Zn-Pb Redox System C1. Clean the Zn metal strip with sandpaper. C2. Place the clean metal in a small test tube. C3. Fill the test tube with 0.1M Pb(NO 3 ) 2 solution be sure that the Zn metal strip is completely submerged. C4. After 2.5 minutes, record your observations. C5. After 5 minutes, record your observations. C6. Clean the Zn metal strip with sandpaper and dry using paper towels and return to the container C7. Discard the Pb(NO 3 ) 2 solution into the container labeled Discarded Pb(NO 3 ) 2 Solution. D. Electrochemical Cells D1. Place about 5 ml of solutions 0.1 M Cu(NO 3 ) 2, 0.1 M Zn(NO 3 ) 2, 0.1 M Pb(NO 3 ) 2, and 0.1 M KNO 3 into small labeled beakers. D2. Clean the copper, zinc, and lead electrodes using steel wool or sandpaper and rinse with deionized water. D3. Place each metal electrode in its corresponding ionic solution; e.g. copper strip goes into the Cu(NO 3 ) 2 solution. It is important that the correct metal is in the correct solution or your cell will not work properly. Revision F10 IB Page 3 of 8

4 D4. Obtain small strips of filter paper to be used as salt bridges. Completely wet one strip in the beaker containing 0.1 M KNO 3. Carefully remove the completely wet strip and place one end in the Cu(NO 3 ) 2 solution and the other in the Zn(NO 3 ) 2 solution. The salt bridge should not touch the electrodes. D5. Attach one alligator clip from the voltmeter to the Cu electrode and the second clip to the Zn electrode. If the voltmeter has a negative voltage, reverse the hookup so that each clip is now attached to the other metal in the pair. D6. Record the voltage of the electrochemical cell. D7. Repeat for the remaining cells, recording the positive cell voltages and using a new wet piece of filter paper as a salt bridge for each. D8. Clean the metal strips with sandpaper or steel wool. Dry each strip using paper towels Disposal Dispose of all solutions into the appropriate waste container. Return all metal pieces to their original container clean and dry. Clean-Up Clean and dry your work area with water. Wash your hands before leaving the laboratory. Wash all glassware with soap then rinse 3 times with tap water, and once with deionized water. Calculations Oxidation Reduction Reactions 1. Consider the metals and ions involved to write the oxidation and reduction half-reactions. 2. Obtain the potentials from Table 1 to calculate E 0 cell. Electrochemical Cells 1. Calculate the net cell voltage by using the tabulated potentials. 2. Calculate the percent deviation between the experimental and calculated cell potentials by the following equation: % Deviation = E cell Measured E cell Calculated E cell Calculated x 100 Revision F10 IB Page 4 of 8

5 Data Sheet Name: Oxidation Reduction Reactions A. Zn-Cu System Experimental Observations Oxidation Half Reduction Half E o ox = E o red = Net E o cell = B. Pb-Cu System Experimental Observations Oxidation Half Reduction Half E o ox = E o red = Net E o cell = Revision F10 IB Page 5 of 8

6 C. Zn-Pb System Experimental Observations Oxidation Half Reduction Half E o ox = E o red = Net E o cell = D. Electrochemical Cells Metals Used E cell Measured E cell Calculated from tabulated potentials % Deviation Cu-Zn Cu-Pb Pb-Zn Revision F10 IB Page 6 of 8

7 Post-lab Assignment Name: 1. Complete the following table with the observed reactions for the electrochemical cells. Write the correct oxidation and reduction half-reaction in the appropriate column for each. Anode Cathode Zn-Cu Cu-Pb Pb-Zn 2. Compare the measured and calculated potentials or your electrochemical cells. Provide an explanation of your percent errors. 3. Based on your voltage measurements for the electrochemical cells, give the order of activity of the metals for the metals used in this experiment. List the most active (easiest to oxidize) metal first. Revision F10 IB Page 7 of 8

8 Pre-lab Assignment Name: 1. Define voltaic cell. 2. Consider the following two reduction reactions and their standard electrode potentials: Al +3 (aq) + 3e - Al(s) Cd +2 (aq) + 2e - Cd(s) E = V E = V a. Write the cell reaction for a voltaic cell based on these two electrodes and calculate the standard cell potential. b. Which species is easier to oxidize and why? c. Which species if easier to reduce and why? 3. A strip of tin is placed in a CuSO 4 solution. Do you expect a spontaneous chemical reaction to occur? If so, write the net cell reaction. Revision F10 IB Page 8 of 8