1 / 31

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

agnes
Télécharger la présentation

CHAPTER 16 (pages 776-792)

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


  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

More Related