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ELECTROCHEMISTRY PROBLEMS

ELECTROCHEMISTRY PROBLEMS. 1 . Using the table of standard reduction potentials, arrange the following substances in decreasing order of reduction: Al, Zn, Ni, Ag, H 2 , Na. Reduction potentials: Al = -1.66, Zn = -0.76, Ni = -0.25, Ag = + 0.80, H 2 = 0, Na = -2.71

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ELECTROCHEMISTRY PROBLEMS

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  1. ELECTROCHEMISTRY PROBLEMS

  2. 1. Using the table of standard reduction potentials, arrange the following substances in decreasing order of reduction: Al, Zn, Ni, Ag, H2, Na. Reduction potentials: Al = -1.66, Zn = -0.76, Ni = -0.25, Ag = +0.80, H2 = 0, Na = -2.71 Ag, H2, Ni, Zn, Al, Na

  3. 2. Using the table of standard reduction potentials, arrange the following substances in order of increasing order of reduction: Fe3+, F2, Pb2+, I2, Sn2+, O2 Reduction potentials: Fe3+ = 0.77, F2 = 2.87, Pb2+ = -0.13, I2 = 0.54, Sn2+ = -0.14, O2 = 1.23 Sn+2, Pb+2, I2, Fe+3, O2, F2,

  4. 3. Calculate Eocell for the following cells using this data:Reaction Eo voltsCd2+ + 2e-  Cd - 0.403 2H+ + 2e- H2 0.00 Ag+ + e- Ag 0.80 A. Cadmium and hydrogen 0.403 + 0 = 0.403V B. Silver and hydrogen 0.80 + 0 = 0.80V C. Cadmium and silver 0.403 + 0.80 = 1.20V

  5. 5. For the following voltaic cell, write the half-reactions, designating which is oxidation and which is reduction. Write the overall cell reaction and calculate the voltage of the cell made from standard electrodes. The cell has electrodes of solid cobalt and nickel, each immersed in a 1 M solution of their ions. ox Co  Co2+ + 2e- 0.28V red Ni2+ + 2e-Ni -0.25V Co + Ni2+ Co2+ + Ni + 0.03V

  6. 6. Determine the cell reaction and the standard cell potential, Eocell, for the voltaic cells composed of the following half-cells. a. Mg2+(aq) + 2e- Mg(s) Cl2(g) + 2e- 2Cl-(aq) Mg2+(aq) + 2e- Mg(s) -2.37V Cl2(g) + 2e- 2Cl-(aq) 1.36V Mg  Mg2+ + 2e- 2.37V Cl2 + 2e- 2Cl- 1.36V Mg + Cl2 Mg2+ + 2Cl- 3.73V

  7. b. Ni2+(aq) + 2e-  Ni(s) Ag+(aq) + e- Ag(s) Ni2+(aq) + 2e-  Ni(s) -0.25V Ag+(aq) + e- Ag(s) +0.80V Ni(s)  Ni2+(aq) + 2e- +0.25V Ag+(aq) + e- Ag(s) +0.80V Ni(s) + 2Ag+(aq)  Ni2+(aq) + 2Ag(s) +1.05V

  8. c. MnO4-(aq) + 8H+(aq) + 5e-  Mn2+(aq) + 4H2O(l) Cd2+(aq) + 2e- Cd(s) MnO4-(aq) + 8H+(aq) + 5e-  Mn2+(aq) + 4H2O(l) 1.51V Cd2+(aq) + 2e- Cd(s) -0.40V MnO4-(aq) + 8H+(aq) + 5e-  Mn2+(aq) + 4H2O(l) 1.51V Cd(s) Cd2+(aq) + 2e- 0.40V 2MnO4- + 16H+ + 5Cd  2Mn2+ + 8H2O + 5Cd2+ 1.91V

  9. d. Br2(l) + 2e- 2Br-(aq) Na+(aq) + e-  Na(s) Br2(l) + 2e- 2Br-(aq) 1.07V Na+(aq) + e-  Na(s) -2.71V Br2(l) + 2e- 2Br-(aq) 1.07V Na(s)  Na+(aq) + e- 2.71V Br2 + 2Na  2Br- + 2Na+ 3.78V

  10. 1. Which statement describes electrolysis? a. Reduction occurs at the anode b. Energy is produced c. Oxidation occurs at the cathode d. Positive ions move to the cathode d.

  11. 2. Which of these statements about rusting is true?I. Iron is oxidizedII. Oxygen is the oxidizing agentIII. Iron atoms gain electrons a. I and II only b. II and III only c. I only d. I, II, and III a

  12. 3. The energy source for an ordinary flashlight is a a. dry cell b. voltaic cell c. battery d. all of the above d.

  13. 4. Magnesium metal is prepared by the electrolysis of molten MgCl2. One half-reaction is Mg2+(l) + 2e- Mg(l). a. This half-reaction occurs at the cathode. b. Magnesium ions are oxidized. c. Chloride ions are reduced at the anode. d. Chloride ions gain electrons during this process. a.

  14. 5. If the cell potential for a redox reaction is positive, a. the redox reaction is spontaneous. b. the redox reaction is not spontaneous. c. the reaction only occurs during electrolysis. d. More than one statement is correct. a.

  15. Given the activity series of metals, answer questions 6-12.6. Which metal will more easily lose an electron, sodium or potassium? potassium

  16. 7. Which metal is more easily oxidized, copper or aluminum? aluminum

  17. 8. What is the relationship between ease of oxidation and the activity of a metal? The more active the metal, the more easily it is oxidized.

  18. 9. Describe what would happen if you placed a clean strip of aluminum in a solution of copper(II) sulfate. Explain your answer. The aluminum strip would start to dissolve and would become covered with copper. Aluminum would be oxidized and copper would be reduced.

  19. 10. Would a copper strip placed in a solution containing zinc ions react spontaneously with the zinc ions? No, Cu is less active than Zn.

  20. 11. Based on the positions of zinc and iron in the table, explain how attaching zinc blocks to a steel ship hull protects the steel from corrosion. Since Zn is more active than Fe, Zn is more easily oxidized. It will lose electrons instead of the iron. This block of Zn can be replaced more easily than the ship’s hull.

  21. 12. Write the half-reaction for the reduction of aluminum ions. Al3+ + 3e- Al

  22. 13. Write an equation for the decomposition for water by electrolysis. 2H2O  2H2 + O2

  23. 14. At which electrode, A or B, is hydrogen produced? (see picture) Hydrogen must be produced at electrode A because there is twice the volume of gas.

  24. 15. The equation for the electrolysis of brine is 2NaCl(aq) + 2H2O(l)Cl2(g) + H2(g) + 2NaOH(aq) How would you modify the electrolysis diagram to make it quantitatively represent the formation of hydrogen and chlorine. The volumes of the two gases would need to be equal.

  25. 16. What do voltaic cells and electrolytic cells have in common? How do they differ? Voltaic and electrolytic cells both involve redox reactions (the transfer of electrons). Voltaic cells use chemical energy to produce electrical energy while electrolytic cells use electrical energy to produce chemical energy. Voltaic cells use spontaneous redox reactions. Electrolytic cells do not.

  26. In both types of cells, oxidation occurs at the anode and reduction occurs at the cathode. In both types of cells, electrons travel from anode to cathode. In a voltaic cell, the anode is negative and the cathode is positive. In an electrolytic cell, the anode is positive and the cathode is negative.

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