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By: Maggie Dang Electrochemical Cells https://egmanual.poly.edu/index.php?title=Lemon_Car.

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Presentation on theme: "By: Maggie Dang Electrochemical Cells https://egmanual.poly.edu/index.php?title=Lemon_Car."— Presentation transcript:

1 By: Maggie Dang Electrochemical Cells https://egmanual.poly.edu/index.php?title=Lemon_Car

2 Background Information A chemical reaction that involves the transfer of electrons from one substance to another is an oxidation-reduction (redox) reaction. A chemical reaction that involves the transfer of electrons from one substance to another is an oxidation-reduction (redox) reaction. Experimentally, when copper wire is placed into a silver ion solution copper atoms spontaneously lose electrons (copper atoms are oxidized) to silver ions ( which are reduced). Silver ions migrate to the copper atoms to pick up electrons and form silver atoms at the copper metal/solution interface; the copper ions that form then move into the solution away from the interface. The overall reaction that occurs at the interface is Experimentally, when copper wire is placed into a silver ion solution copper atoms spontaneously lose electrons (copper atoms are oxidized) to silver ions ( which are reduced). Silver ions migrate to the copper atoms to pick up electrons and form silver atoms at the copper metal/solution interface; the copper ions that form then move into the solution away from the interface. The overall reaction that occurs at the interface is Cu (s) + Ag + (aq) -> 2Ag(s) + Cu 2+ (aq) Cu (s) + Ag + (aq) -> 2Ag(s) + Cu 2+ (aq) This redox reaction can be divided into an oxidation and a reduction half-reaction. Each half-reaction, called a redox couple, consists of the reduced state and the oxidized state of the substance. This redox reaction can be divided into an oxidation and a reduction half-reaction. Each half-reaction, called a redox couple, consists of the reduced state and the oxidized state of the substance. Cu(s)  Cu 2+ (aq) + 2e - (Oxidation half-reaction) Cu(s)  Cu 2+ (aq) + 2e - (Oxidation half-reaction) 2Ag + (aq) + 2e -  2Ag (s) (Reduction half-reaction) 2Ag + (aq) + 2e -  2Ag (s) (Reduction half-reaction) es/en_oxidation_reduction.jpg

3 A voltaic cell (galvanic cell) is designed to take advantage of this spontaneous transfer of electrons. A voltaic cell separates the copper metal from the silver ions and forces the electrons to pass externally through a wire, an external circuit. A voltaic cell (galvanic cell) is designed to take advantage of this spontaneous transfer of electrons. A voltaic cell separates the copper metal from the silver ions and forces the electrons to pass externally through a wire, an external circuit. The two redox couples are placed in separate compartments, called half-cells. Each half-cell consists of an electrode, usually the metal (reduced state) of the redox couple, and a solution containing the corresponding cation (oxidized state) of the redox couple. The electrodes of the half-cells are connected by a wire; this is where the electrons flow, providing current for the external circuit. The two redox couples are placed in separate compartments, called half-cells. Each half-cell consists of an electrode, usually the metal (reduced state) of the redox couple, and a solution containing the corresponding cation (oxidized state) of the redox couple. The electrodes of the half-cells are connected by a wire; this is where the electrons flow, providing current for the external circuit. A salt bridge, which connects the two half-cells, completes the construction of the voltaic cell (and the circuit). The salt bridge permits limited movement of ions from one half-cell to the other, the internal circuit, so that when the cell operates, electrical neutrality is maintained in each half-cell. A salt bridge, which connects the two half-cells, completes the construction of the voltaic cell (and the circuit). The salt bridge permits limited movement of ions from one half-cell to the other, the internal circuit, so that when the cell operates, electrical neutrality is maintained in each half-cell.

4 The electrode at which reduction occurs is called the cathode; the electrode at which oxidation occurs is called the anode. The electrode at which reduction occurs is called the cathode; the electrode at which oxidation occurs is called the anode. Because oxidation releases electrons to the electrode to provide a current in the external circuit, the anode is designated the negative electrode in a galvanic cell. The reduction process draws electrons from the circuit and supplies them to the ions in solution; the cathode is the positive electrode. Because oxidation releases electrons to the electrode to provide a current in the external circuit, the anode is designated the negative electrode in a galvanic cell. The reduction process draws electrons from the circuit and supplies them to the ions in solution; the cathode is the positive electrode. The cell potential of a galvanic cell is due to the difference in tendencies of the two metals to oxidize or of their ions to reduce. The cell potential of a galvanic cell is due to the difference in tendencies of the two metals to oxidize or of their ions to reduce.

5 A measured reduction potential, the tendency for a substance to gain electrons, is the value used to identify the relative ease of reduction for a half-reaction. A measured reduction potential, the tendency for a substance to gain electrons, is the value used to identify the relative ease of reduction for a half-reaction. A voltmeter, placed in the external circuit between the two electrodes, measures the cell potential, E o cell, a value that represents the difference between the tendencies of the metal ions in their respective half-cells to undergo reduction. A voltmeter, placed in the external circuit between the two electrodes, measures the cell potential, E o cell, a value that represents the difference between the tendencies of the metal ions in their respective half-cells to undergo reduction.

6 Purpose To determine how closely the voltages found in the experiment compare to the E cell o voltages calculated using the standard reduction potential chart To determine how closely the voltages found in the experiment compare to the E cell o voltages calculated using the standard reduction potential chart ons/industrial/2763/images/ch954_2.gif

7 Materials Copper(II) nitrate Cu(NO 3 ) 2 1.0M Copper(II) nitrate Cu(NO 3 ) 2 1.0M Iron(III) nitrate Fe(NO 3 ) 3, 1.0M Iron(III) nitrate Fe(NO 3 ) 3, 1.0M Lead(II) nitrate Pb(NO 3 ) 2 1.0M Lead(II) nitrate Pb(NO 3 ) 2 1.0M Magnesium nitrate Mg (NO 3 ) 2 1.0M Magnesium nitrate Mg (NO 3 ) 2 1.0M Silver nitrate AgNO 3 1.0M Silver nitrate AgNO 3 1.0M Zinc nitrate Zn(NO 3 ) 2 1.0M Zinc nitrate Zn(NO 3 ) 2 1.0M Sodium chloride NaCl 1.0M Sodium chloride NaCl 1.0M 2 beakers (50 mL ) 2 beakers (50 mL ) Voltmeter Voltmeter 2 Alligator Clips 2 Alligator Clips A salt bridge of 1.0M KNO3 A salt bridge of 1.0M KNO3 Strips of Zn, Cu, and Mg Strips of Zn, Cu, and Mg Iron nails Iron nails Lead Piping Lead Piping Strips of filter paper Strips of filter paper b/bukm2040.jpg

8 Setup

9 Procedures Prepare a test cell to measure the voltage of the copper and zinc half-cells. Prepare a test cell to measure the voltage of the copper and zinc half-cells. Put approximately 2 mL of 1.0 M Cu(NO 3 ) 2 in one beaker and 2 mL of 1.0 M Zn(NO 3 ) 2 in the other beaker. Put approximately 2 mL of 1.0 M Cu(NO 3 ) 2 in one beaker and 2 mL of 1.0 M Zn(NO 3 ) 2 in the other beaker. Polish small strips of zinc and copper metal, and place the metal in the appropriate beaker containing the solution of the ions of that metal. Polish small strips of zinc and copper metal, and place the metal in the appropriate beaker containing the solution of the ions of that metal. copper_image.jpg Copper m/images/products/H- 4911beakers.jpg nuclear.com/zincs tick.jpg ZINC 

10 Take a small strip of filter paper that has been soaked in KNO3 solution, and drape it across the wells so that one end dips in the solution in each well. This will act as the salt bridge.Use a voltmeter to measure the voltages between the two half-cells. Take a small strip of filter paper that has been soaked in KNO3 solution, and drape it across the wells so that one end dips in the solution in each well. This will act as the salt bridge.Use a voltmeter to measure the voltages between the two half-cells. Connect the meter so the voltage reading is positive. When the voltmeter reads a positive voltage, the electrode connected to the positive terminal is the cathode and is undergoing reduction, while oxidation is occurring at the electrode connected to the negative terminal, the anode. Connect the meter so the voltage reading is positive. When the voltmeter reads a positive voltage, the electrode connected to the positive terminal is the cathode and is undergoing reduction, while oxidation is occurring at the electrode connected to the negative terminal, the anode. Record data. Record data. mlab/chm1046lmanual/electr ochem/boilkno3.JPG KNO3

11 Rinse out the beakers and measure the voltages of the other electrodes the same way with zinc since it is designated to be the standard electrode. Rinse out the beakers and measure the voltages of the other electrodes the same way with zinc since it is designated to be the standard electrode. The electrodes to be tested are : Ag, Ag + ; Cu, Cu + ; Fe, Fe 3+ ; Mg,Mg 2+ ; Pb, Pb 2+ ; Zn, Zn 2+ The electrodes to be tested are : Ag, Ag + ; Cu, Cu + ; Fe, Fe 3+ ; Mg,Mg 2+ ; Pb, Pb 2+ ; Zn, Zn 2+ Record data. Record data. Voltmeter content/uploads/2009/03/voltmeter-car- battery-hull.jpg

12 Voltage of each half-cell versus the zinc electrode VoltageAnodeCathode Zn versus Ag 1.40ZnAg Zn versus Cu.99ZnCu Zn versus Fe.55ZnFe Zn versus Mg.60MgZn Zn versus Pb.48ZnPb

13 Write reduction equations for each metal ion. Write reduction equations for each metal ion. Record the standard potentials for all of the electrodes using the standard reduction potentials chart, and calculate the potential energy of the entire cell by using the equation E o cell=E o red (cathode) – E o red (anode) Record the standard potentials for all of the electrodes using the standard reduction potentials chart, and calculate the potential energy of the entire cell by using the equation E o cell=E o red (cathode) – E o red (anode)

14 Calculations Standard Electrode  Zinc with E o = -.76 Standard Electrode  Zinc with E o = -.76 Ag+ + e-  Ag E o =.80 Ag+ + e-  Ag E o = =1.56 Cu2+2e-  Cu E o =.34 Cu2+2e-  Cu E o = =1.1 Fe  Fe E o = -.44 Fe  Fe E o = =.32 Mg2+ 2e-  Mg E o = -.76 Mg2+ 2e-  Mg E o = =1.61 Pb2+ 2e-  Pb E o = -.13 Pb2+ 2e-  Pb E o = =.63 Zn2+ 2e-  Zn E o = -.76 Zn2+ 2e-  Zn E o = = 0

15 Reduction Equations for Each Ion Arranged in Decreasing Order of Potential Reduction Reaction Voltages using Zinc as the Standard, E Zn Eored Voltages Using Standard Reduction Potential Chart Ag + + e -  Ag Cu e -  Cu Fe e -  Fe Mg e -  Mg Pb e -  Pb Zn e -  Zn

16 Conclusion Comparing to the E o red voltages calculated from using the standard reduction potential chart, the voltages measured using Zinc as the standard, E Zn, in the experiment are relatively close. Comparing to the E o red voltages calculated from using the standard reduction potential chart, the voltages measured using Zinc as the standard, E Zn, in the experiment are relatively close. Therefore, this is an accurate experiment that can be used to measure potential energies among different types of electrodes. Therefore, this is an accurate experiment that can be used to measure potential energies among different types of electrodes.

17 Sources Vonderbrink, Sally Ann. Laboratory Experiments for Advanced Placement Chemistry Student Edition. Flinn Scientific, Inc. Batavia, IL Vonderbrink, Sally Ann. Laboratory Experiments for Advanced Placement Chemistry Student Edition. Flinn Scientific, Inc. Batavia, IL Beran, J.A. Laboratory Manual for Principles of General Chemistry. John Wiley & Sons, Inc. Hoboken, NJ Beran, J.A. Laboratory Manual for Principles of General Chemistry. John Wiley & Sons, Inc. Hoboken, NJ

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