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CHEM 160 General Chemistry II Lecture Presentation Electrochemistry

CHEM 160 General Chemistry II Lecture Presentation Electrochemistry. December 1, 2004 Chapter 20. Electrochemistry. Electrochemistry deals with interconversion between chemical and electrical energy. Electrochemistry. Electrochemistry

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CHEM 160 General Chemistry II Lecture Presentation Electrochemistry

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  1. CHEM 160 General Chemistry IILecture PresentationElectrochemistry December 1, 2004 Chapter 20

  2. Electrochemistry • Electrochemistry • deals with interconversion between chemical and electrical energy

  3. Electrochemistry • Electrochemistry • deals with the interconversion between chemical and electrical energy • involves redox reactions

  4. Electrochemistry • Electrochemistry • deals with interconversion between chemical and electrical energy • involves redox reactions • electron transfer reactions • Oh No! They’re back!

  5. Redox reactions (quick review) • Oxidation • Reduction • Reducing agent • Oxidizing agent

  6. Redox reactions (quick review) • Oxidation • loss of electrons • Reduction • Reducing agent • Oxidizing agent

  7. Redox reactions (quick review) • Oxidation • loss of electrons • Reduction • gain of electrons • Reducing agent • Oxidizing agent

  8. Redox reactions (quick review) • Oxidation • loss of electrons • Reduction • gain of electrons • Reducing agent • donates the electrons and is oxidized • Oxidizing agent

  9. Redox reactions (quick review) • Oxidation • loss of electrons • Reduction • gain of electrons • Reducing agent • donates the electrons and is oxidized • Oxidizing agent • accepts electrons and is reduced

  10. Redox Reactions • Direct redox reaction

  11. Redox Reactions • Direct redox reaction • Oxidizing and reducing agents are mixed together

  12. Direct Redox Reaction Zn rod CuSO4(aq) (Cu2+)

  13. Direct Redox Reaction Zn rod CuSO4(aq) (Cu2+) Deposit of Cu metal forms

  14. Redox Reactions • Direct redox reaction • Oxidizing and reducing agents are mixed together • Indirect redox reaction • Oxidizing and reducing agents are separated but connected electrically • Example • Zn and Cu2+ can be reacted indirectly • Basis for electrochemistry • Electrochemical cell

  15. Electrochemical Cells

  16. Electrochemical Cells • Voltaic Cell • cell in which a spontaneous redox reaction generates electricity • chemical energy  electrical energy

  17. Electrochemical Cells

  18. Electrochemical Cells Voltaic Cell

  19. Electrochemical Cells • Electrolytic Cell • electrochemical cell in which an electric current drives a nonspontaneous redox reaction • electrical energy  chemical energy

  20. Cell Potential

  21. Cell Potential • Cell Potential (electromotive force), Ecell (V) • electrical potential difference between the two electrodes or half-cells • Depends on specific half-reactions, concentrations, and temperature • Under standard state conditions ([solutes] = 1 M, Psolutes = 1 atm), emf = standard cell potential, Ecell • 1 V = 1 J/C • driving force of the redox reaction

  22. Cell Potential low electrical potential high electrical potential

  23. Cell Potential Ecell = Ecathode - Eanode = Eredn - Eox E°cell = E°cathode - E°anode = E°redn - E°ox (Ecathode and Eanode are reduction potentials by definition.)

  24. Cell Potential • E°cell = E°cathode - E°anode = E°redn - E°ox • Ecell can be measured • Absolute Ecathode and Eanode values cannot • Reference electrode • has arbitrarily assigned E • used to measure relative Ecathode and Eanode for half-cell reactions • Standard hydrogen electrode (S.H.E.) • conventional reference electrode

  25. Standard Hydrogen Electrode • E = 0 V (by definition; arbitrarily selected) • 2H+ + 2e-  H2

  26. Example 1 A voltaic cell is made by connecting a standard Cu/Cu2+ electrode to a S.H.E. The cell potential is 0.34 V. The Cu electrode is the cathode. What is the standard reduction potential of the Cu/Cu2+ electrode?

  27. Example 2 A voltaic cell is made by connecting a standard Zn/Zn2+ electrode to a S.H.E. The cell potential is 0.76 V. The Zn electrode is the anode of the cell. What is the standard reduction potential of the Zn/Zn2+ electrode?

  28. Standard Electrode Potentials • Standard Reduction Potentials, E° • E°cell measured relative to S.H.E. (0 V) • electrode of interest = cathode • If E° < 0 V: • Oxidizing agent is harder to reduce than H+ • If E° > 0 V: • Oxidizing agent is easier to reduce than H+

  29. Reduction Half-Reaction E(V) F2(g) + 2e- 2F-(aq) 2.87 Au3+(aq) + 3e- Au(s) 1.50 Cl2(g) + 2 e- 2Cl-(aq) 1.36 Cr2O72-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O 1.33 O2(g) + 4H+ + 4e- 2H2O(l) 1.23 Ag+(aq) + e- Ag(s) 0.80 Fe3+(aq) + e- Fe2+(aq) 0.77 Cu2+(aq) + 2e- Cu(s) 0.34 Sn4+(aq) + 2e- Sn2+(aq) 0.15 2H+(aq) + 2e- H2(g) 0.00 Sn2+(aq) + 2e- Sn(s) -0.14 Ni2+(aq) + 2e- Ni(s) -0.23 Fe2+(aq) + 2e- Fe(s) -0.44 Zn2+(aq) + 2e- Zn(s) -0.76 Al3+(aq) + 3e- Al(s) -1.66 Mg2+(aq) + 2e- Mg(s) -2.37 Li+(aq) + e- Li(s) -3.04 Standard Reduction Potentials

  30. Uses of Standard Reduction Potentials • Compare strengths of reducing/oxidizing agents. • the more - E°, stronger the red. agent • the more + E°, stronger the ox. agent

  31. Reduction Half-Reaction E(V) F2(g) + 2e- 2F-(aq) 2.87 Au3+(aq) + 3e- Au(s) 1.50 Cl2(g) + 2 e- 2Cl-(aq) 1.36 Cr2O72-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O 1.33 O2(g) + 4H+ + 4e- 2H2O(l) 1.23 Ag+(aq) + e- Ag(s) 0.80 Fe3+(aq) + e- Fe2+(aq) 0.77 Cu2+(aq) + 2e- Cu(s) 0.34 Sn4+(aq) + 2e- Sn2+(aq) 0.15 2H+(aq) + 2e- H2(g) 0.00 Sn2+(aq) + 2e- Sn(s) -0.14 Ni2+(aq) + 2e- Ni(s) -0.23 Fe2+(aq) + 2e- Fe(s) -0.44 Zn2+(aq) + 2e- Zn(s) -0.76 Al3+(aq) + 3e- Al(s) -1.66 Mg2+(aq) + 2e- Mg(s) -2.37 Li+(aq) + e- Li(s) -3.04 Standard Reduction Potentials Ox. agent strength increases Red. agent strength increases

  32. Uses of Standard Reduction Potentials • Determine if oxidizing and reducing agent react spontaneously • diagonal rule ox. agent spontaneous red. agent

  33. Uses of Standard Reduction Potentials • Determine if oxidizing and reducing agent react spontaneously more + Spontaneous rxn if E°cathode > E°anode Cathode (reduction) E°redn (cathode) E°redn (V) Anode (oxidation) E°redn (anode) more -

  34. Reduction Half-Reaction E(V) F2(g) + 2e- 2F-(aq) 2.87 Au3+(aq) + 3e- Au(s) 1.50 Cl2(g) + 2 e- 2Cl-(aq) 1.36 Cr2O72-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O 1.33 O2(g) + 4H+ + 4e- 2H2O(l) 1.23 Ag+(aq) + e- Ag(s) 0.80 Fe3+(aq) + e- Fe2+(aq) 0.77 Cu2+(aq) + 2e- Cu(s) 0.34 Sn4+(aq) + 2e- Sn2+(aq) 0.15 2H+(aq) + 2e- H2(g) 0.00 Sn2+(aq) + 2e- Sn(s) -0.14 Ni2+(aq) + 2e- Ni(s) -0.23 Fe2+(aq) + 2e- Fe(s) -0.44 Zn2+(aq) + 2e- Zn(s) -0.76 Al3+(aq) + 3e- Al(s) -1.66 Mg2+(aq) + 2e- Mg(s) -2.37 Li+(aq) + e- Li(s) -3.04 Standard Reduction Potentials

  35. Uses of Standard Reduction Potentials • Calculate E°cell • E°cell = E°cathode - E°anode • Greater E°cell, greater the driving force • E°cell > 0 : spontaneous redox reactions • E°cell < 0 : nonspontaeous redox reactions

  36. Example 3 A voltaic cell consists of a Ag electrode in 1.0 M AgNO3 and a Cu electrode in 1 M Cu(NO3)2. Calculate E°cell for the spontaneous cell reaction at 25°C.

  37. Reduction Half-Reaction E(V) F2(g) + 2e- 2F-(aq) 2.87 Au3+(aq) + 3e- Au(s) 1.50 Cl2(g) + 2 e- 2Cl-(aq) 1.36 Cr2O72-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O 1.33 O2(g) + 4H+ + 4e- 2H2O(l) 1.23 Ag+(aq) + e- Ag(s) 0.80 Fe3+(aq) + e- Fe2+(aq) 0.77 Cu2+(aq) + 2e- Cu(s) 0.34 Sn4+(aq) + 2e- Sn2+(aq) 0.15 2H+(aq) + 2e- H2(g) 0.00 Sn2+(aq) + 2e- Sn(s) -0.14 Ni2+(aq) + 2e- Ni(s) -0.23 Fe2+(aq) + 2e- Fe(s) -0.44 Zn2+(aq) + 2e- Zn(s) -0.76 Al3+(aq) + 3e- Al(s) -1.66 Mg2+(aq) + 2e- Mg(s) -2.37 Li+(aq) + e- Li(s) -3.04 Standard Reduction Potentials

  38. Example 4 A voltaic cell consists of a Ni electrode in 1.0 M Ni(NO3)2 and an Fe electrode in 1 M Fe(NO3)2. Calculate E°cell for the spontaneous cell reaction at 25°C.

  39. Reduction Half-Reaction E(V) F2(g) + 2e- 2F-(aq) 2.87 Au3+(aq) + 3e- Au(s) 1.50 Cl2(g) + 2 e- 2Cl-(aq) 1.36 Cr2O72-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O 1.33 O2(g) + 4H+ + 4e- 2H2O(l) 1.23 Ag+(aq) + e- Ag(s) 0.80 Fe3+(aq) + e- Fe2+(aq) 0.77 Cu2+(aq) + 2e- Cu(s) 0.34 Sn4+(aq) + 2e- Sn2+(aq) 0.15 2H+(aq) + 2e- H2(g) 0.00 Sn2+(aq) + 2e- Sn(s) -0.14 Ni2+(aq) + 2e- Ni(s) -0.23 Fe2+(aq) + 2e- Fe(s) -0.44 Zn2+(aq) + 2e- Zn(s) -0.76 Al3+(aq) + 3e- Al(s) -1.66 Mg2+(aq) + 2e- Mg(s) -2.37 Li+(aq) + e- Li(s) -3.04 Standard Reduction Potentials

  40. Cell Potential • Is there a relationship between Ecell and DG for a redox reaction?

  41. Cell Potential • Relationship between Ecell and DG: • DG = -nFEcell • F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn.

  42. Cell Potential • Relationship between Ecell and DG: • DG = -nFEcell • F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn. • 1 J = CV • G < 0, Ecell > 0 = spontaneous

  43. Equilibrium Constants from Ecell • Relationship between Ecell and DG: • DG = -nFEcell • F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn • 1 J = CV • G < 0, Ecell > 0 = spontaneous • Under standard state conditions: • DG° = -nFE°cell

  44. Equilibrium Constants from Ecell • Relationship between Ecell and DG: • DG = -nFEcell • F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn • 1 J = CV • G < 0, Ecell > 0 = spontaneous • Under standard state conditions: • DG° = -nFE°cell

  45. Equilibrium Constants from Ecell • Relationship between Ecell and DG: • DG = -nFEcell • F = Faraday constant = 96500 C/mol e-’s, n = # e-’s transferred redox rxn • 1 J = CV • G < 0, Ecell > 0 = spontaneous • Under standard state conditions: • DG° = -nFE°cell and • DG° = -RTlnK so • -nFE°cell = -RTlnK

  46. Calorimetric Data DH° DS° Electrochemical Data Composition Data DG° E°cell Equilibrium constants K

  47. Example 5 Calculate E°cell, DG°, and K for the voltaic cell that uses the reaction between Ag and Cl2 under standard state conditions at 25°C.

  48. The Nernst Equation • DG depends on concentrations • DG = DG° + RTlnQ and • DG = -nFEcell and DG° = -nFE°cell thus • -nFEcell = -nFE°cell + RTlnQ or • Ecell = E°cell - (RT/nF)lnQ (Nernst eqn.)

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