notes_Lecture_14_110510

notes_Lecture_14_110510 - Lecture 14: The Nernst Equation...

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Lecture 14: The Nernst Equation Reading: Outline – How does E cell depend on concentration? (11.4) Nernst equation Concentration Cells – Electrolytic cells: an example of nonspontaneous electron flow (11.7) Problems for Extra Practice – 47, 49, 53, 55, 59 (Chapter 11)
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cell and Δ G (cont.) Developing the “big picture” Also recall: Δ G = w max The Gibbs energy changes as a system proceeds towards equilibrium. ..this change can be used to do useful (non-PV) work.
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Which of the following balanced redox reactions has the largest equilibrium constant? Zn + 2 + Cu Cu + 2 + Zn E cell 0 = - 1.10 V 2 Hg + 2 + 2 I - Hg 2 + 2 + I 2 E cell 0 = + 0.38 V 2 AgCl + Fe 2 Ag + 2 Cl - + Fe + 2 E cell 0 = + 0.66 V 2 H + + S - 2 S + H 2 E cell 0 = + 0.48 V
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Concentration and E cell • Consider the following redox reaction: Δ G ° = -nFE ° cell < 0 (spontaneous) • What if [Cu 2+ ] = 3 M? Expect driving force for product formation to increase. Therefore Δ G decreases, and E cell increases How does E cell depend on concentration? Fe(s) + Cu 2+ (aq) Fe 2+ (aq) + Cu(s) E ° cell = +0.78 V
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Concentration and E cell (cont.) • Recall, in general: Δ G = Δ G ° + RTln(Q) • We also know that: Δ G = -nFE cell -nFE cell = -nFE ° cell + RTln(Q) E cell = E ° cell - (RT/nF)ln(Q) E cell = E ° cell - (0.0591 V/n)log(Q) The Nernst Equation
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• Back to our example: Δ G ° = -nFE ° cell < 0 (spontaneous) [Cu 2+ ] = 3 M, everything else at standard state. Fe(s) + Cu
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notes_Lecture_14_110510 - Lecture 14: The Nernst Equation...

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