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CHAPTER 16 (pages 776-792)

CHAPTER 16 (pages 776-792). Oxidation and Reduction Galvanic Cells, Half Reactions ( E° anode & E° cathode ) Standard Reduction Potential (E°) Nernst Equation, and the dependence of Potential on Concentration Relationship between Equilibrium Constant and Standard Potential

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CHAPTER 16 (pages 776-792)

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  1. CHAPTER 16 (pages 776-792) Oxidation and Reduction Galvanic Cells, Half Reactions (E°anode & E°cathode) Standard Reduction Potential (E°) Nernst Equation, and the dependence of Potential on Concentration Relationship between Equilibrium Constant and Standard Potential Driving Force, ΔG and ε

  2. REDOX REACTIONS 3 H2S + 2 NO3– + 2 H+⇋3S + 2 NO + 4 H2O MnO2 + 4 HBr⇋MnBr2+ Br2 + 2 H2O

  3. OBSERVED REDOX PROCESSES

  4. GALVANIC CELLS

  5. INERT ELECTRODES

  6. STANDARD REDUCTION POTENTIALS

  7. MEASURING STANDARD POTENTIALS

  8. CALCULATING STANDARD CELL POTENTIAL Al(s) + NO3−(aq) + 4 H+(aq) ⇋ Al3+(aq) + NO(g) + 2 H2O(l)

  9. ADDITIONAL EXAMPLE Fe(s)+ Mg2+(aq) ⇋ Fe2+(aq) + Mg(s)

  10. ox: Fe(s)  Fe2+(aq) + 2 e−E = +0.45 V red: Pb2+(aq) + 2 e−  Pb(s)E = −0.13 V tot: Pb2+(aq) + Fe(s)  Fe2+(aq) + Pb(s)E = +0.32 V

  11. ELECTROMOTIVE POTENTIAL

  12. E°CELL,ΔG° AND K Under standard state conditions, a reaction will spontaneously proceeds in the forward direction if: • ΔG° < 1 (negative) • E° > 1 (positive) • K > 1

  13. Design a voltaic cell with the following half cells and complete the calculations: Ag+ (aq) + 1e- Ag (s) Eo = 0.80 V Pb2+ (aq) + 2e- Pb (s) Eo = -0.13 V • Calculate the Eocell(potential at standard conditions) • Calculate Go. • Calculate  • Calculate the Ecell if [Ag+] = 2.0 M and [Pb2+] = 1.0 x 10-4 M.

  14. Williams, spring 2009stop here

  15. Design a voltaic cell with the following half cells and complete the calculations: Ag+ (aq) + 1e- Ag (s) Eo = 0.80 V Pb2+ (aq) + 2e- Pb (s) Eo = -0.13 V Calculate the Eocell(potential at standard conditions)

  16. Design a voltaic cell with the following half cells and complete the calculations: Ag+ (aq) + 1e- Ag (s) Eo = 0.80 V Pb2+ (aq) + 2e- Pb (s) Eo = -0.13 V Calculate Go.

  17. Design a voltaic cell with the following half cells and complete the calculations: Ag+ (aq) + 1e- Ag (s) Eo = 0.80 V Pb2+ (aq) + 2e- Pb (s) Eo = -0.13 V Calculate 

  18. Design a voltaic cell with the following half cells and complete the calculations: Ag+ (aq) + 1e- Ag (s) Eo = 0.80 V Pb2+ (aq) + 2e- Pb (s) Eo = -0.13 V Calculate the Ecell if [Ag+] = 2.0 M and Pb2+] = 1.0 x 10-4 M.

  19. OBJECTIVE 11.4: PROVIDE A THOROUGH OVERVIEW OF APPLICATIONS OF ELECTROCHEMICAL CELLS INCLUDING FUEL CELLS, CORROSION, AND OTHER TOPICS AS TIME PERMITS.

  20. CORROSION • corrosion is the spontaneous oxidation of a metal by chemicals in the environment • since many materials we use are active metals, corrosion can be a very big problem

  21. RUSTING • rust is hydrated iron(III) oxide • moisture must be present • electrolytes promote rusting • acids promote rusting • lower pH = lower E°red

  22. Dry Cell Batteries

  23. Lead – Acid Storage Battery

  24. Biological Electrochemistry

  25. Lithium Ion Battery

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