Chapter 12 Electrochemistry

Chapter 12 Electrochemistry - L31: Chapter 12:...

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L31: Chapter 12: Electrochemistry The interchange of chemical and electrical energy Review: RedOx rxns (Ch. 6.4) Thermodynamics (Ch. 8) Electrochemical Cells Gibbs Free Energy and Cell Voltage Free Energy, Concentration and K: Nernst Equation Batteries and Fuel Cells Special topics: Methanol and bacterial fuel cells Nanoelectrodes 1 micrometer Shewanella
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Electric current: generated by electrons’ flow Redox rxns: involve transfer of electrons between reagents CuSO 4 + Zn = ZnSO 4 + Cu Half rxns: Cu 2+ + 2e -> Cu 0 reduction rxn Zn 0 –2e -> Zn 2+ oxidation rxn Cu 2+ is reduced <-> Cu 2+ is oxidizing agent Zn is oxidized <-> Zn is reducing agent Balancing redox rxns: see Chapter 6.4
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Electrochemical cell – sustained electron flow spatially separate reducing and oxidizing agents +2 Zn|Zn 2+ ||Cu 2+ |Cu
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Electrochemical cell Anode (oxidation) (attracts anions) Cathode (reduction) (attracts cations) - + http://www2.wwnorton.com/college/chemistry/gilbert/tutorials/ch17.htm; 17.1 section 4
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Cell potential E cell : Potential difference Δ E cell driving force for electrons’ flow (electromotive force, emf) Units: Volt (V) 1V = 1 J/C 1 Volt = 1 Joule of work per 1 Coulomb of charge transferred Galvanic (voltaic) cell: operates spontaneously, Δ E cell >0 Can apply external potential, Δ E ext , then the cell operates as defined by Δ E cell + Δ E ext - can reverse electrons’ flow: electrolytic cell
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Standard reduction potentials Each cell is characterized completely by the cell potential, an intensive property (value does not change with amount) Problem: cannot tabulate all combinations Idea: assign a potential to each half rxn and tabulate them. Then construct a potential for the specific cell (i.e., full redox rxn) from tabulated potentials of half rxns Problem: cannot measure a potential for half a cell Solution: use reference cell and measure potential of each half-rxns with respect to the reference!
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Standard reduction potential: One electrode is a standard hydrogen electrode: 2H + + 2e -> H 2 , [H + ]=1M, p H2 =1 atm, E 0 0 V
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Standard hydrogen reduction reaction: 2H + + 2e -> H 2 , [H + ]=1M, pH 2 =1 atm, E 0 0 V
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Potentials are measured w.r.t. the hydrogen electrode; Reaction is spontaneous when E 0 >0 Example: Cu 2+ + 2e -> Cu 0 E 0 =0.34 V [Cu 2+ ] = 1 M, T = 25º C Overall rxn: Cu 2+ + H 2 -> Cu 0 + 2H + E cell = 0.34 – 0.00 = 0.34 V rxn goes from left to right Zn 2+ + 2e -> Zn 0 E 0 = -0.76 V Overall rxn: Zn 2+ + H 2 -> Zn 0 + 2H + E cell = -0.76 – 0.00 = -0.76 V rxn goes from right to left 2H + + 2e -> H 2 , E 0 =0 V
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What about Cu/CuSO 4 || Zn/ZnSO 4 cell? (1) Cu 2+ + 2e -> Cu 0 E 0 = 0.34 V (2) Zn 2+ + 2e -> Zn 0 E 0 = -0.76 V Overall rxn: (1) + (2) Cu 2+ + Zn = Zn 2+ + Cu E cell = 0.34 –( -0.76 ) = 1.1 V E cell > 0 rxn goes from left to right spontaneously .
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Describe completely the galvanic cell based on the following half rxns under standard conditions: (1) Ag + + e -> Ag E 0 = 0.80 V (2) Fe 3+ + e -> Fe 2+ E 0 = 0.77 V To have E cell > 0 - reverse rxn (2): Ag + + e -> Ag E 0 = 0.80 V Fe 2+ -> Fe 3+ + e E 0 = -0.77 V Overall: Ag + + Fe 2+ -> Ag + Fe 3+ E 0 cell = 0.80 + ( -0.77 )=0.03 V Anode: Pt/Fe 2+ /Fe 3+ , [Fe 2+ ]=[Fe 3+ ]=1M (Pt is inert) Cathode: Ag/Ag + , [Ag + ]=1M Which is stronger oxidizing agent, Ag + or Fe 3+ ?
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Chapter 12 Electrochemistry - L31: Chapter 12:...

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