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Chem. 133 – 3/6 Lecture

Chem. 133 – 3/6 Lecture. Announcements. Grading: I’m still working on the Electronics Lab and Exam 1 Seminars: Friday (regular): David Heppner (Stanford) – Bioinorganic Next Week (Analytical Candidate Tu and Th 3:30 to 4:30) Today’s Lecture (Electrochemistry): Galvanic Cells

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Chem. 133 – 3/6 Lecture

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  1. Chem. 133 – 3/6 Lecture

  2. Announcements • Grading: I’m still working on the Electronics Lab and Exam 1 • Seminars: • Friday (regular): David Heppner (Stanford) – Bioinorganic • Next Week (Analytical Candidate Tu and Th 3:30 to 4:30) • Today’s Lecture (Electrochemistry): • Galvanic Cells • Electrolytic Cells • Nernst Equation and Applications

  3. ElectrochemistryGalvanic Cells What are galvanic cells? Cells that use chemical reactions to generate electrical energy Batteries are examples of useful galvanic cells Example reaction If reactants are placed in a beaker, only products + heat are produced When half reactions are isolated on electrodes, electrical work can be produced GALVANIC CELL voltmeter Zn(s) Ag(s) Zn(s) + 2Ag+→ Zn2+ + 2Ag(s) AgNO3(aq) ZnSO4(aq) Salt Bridge

  4. ElectrochemistryGalvanic Cells Description of how example cell works Reaction on anode = oxidation Anode = Zn electrode (as the Eº for Zn2+ is less than for that for Ag+) So, reaction on cathode must be reduction and involve Ag Oxidation produces e-, so anode has (–) charge (galvanic cells); current runs from cathode to anode Salt bridge allows replenishment of ions as cations migrate to cathode and anions toward anodes GALVANIC CELL voltmeter Ag+ + e- → Ag(s) Zn(s) Ag(s) + – AgNO3(aq) ZnSO4(aq) Zn(s) → Zn2+ + 2e- Salt Bridge

  5. ElectrochemistryGalvanic Cells Cell notation Example Cell: Zn(s)|ZnSO4(aq)||AgNO3(aq)|Ag(s) GALVANIC CELL voltmeter Zn(s) Ag(s) “|” means phase boundary left side for anode (right side for cathode) “||” means salt bridge AgNO3(aq) ZnSO4(aq) Salt Bridge

  6. ElectrochemistryGalvanic Cells Given the following cell, write the cell notation: GALVANIC CELL voltmeter – reads +0.43 V Pt(s) Ag(s) – + AgCl(s) NaCl(aq) FeSO4 (aq), Fe2(SO4)3(aq) Salt Bridge

  7. ElectrochemistryStandard Reduction Potential A half cell or electrode, is half of a galvanic cell A standard electrode is one under standard conditions (e.g. 1 M AgNO3(aq)) Standard reduction potential (Eº) is cell potential when reducing electrode is coupled to standard hydrogen electrode (oxidation electrode) Large + Eº means easily reduced compounds on electrode Large – Eº means easily oxidized compounds on anode Pt(s) Ag(s) H2(g) AgNO3(aq) H+(aq)

  8. ElectrochemistryElectrolytic Cells Used in more advanced electrochemical analysis (not covered in detail) Uses voltage to drive (unfavorable) chemical reactions Example: use of voltage to oxidize phenol in an HPLC electrochemical detector (E° of 0 to 0.5 V needed) anode (note: oxidation driven by voltage, but now + charge) cathode (reduction, - charge)

  9. ElectrochemistryThe Nernst Equation The Nernst Equation relates thermodynamic quantities to electrical quantities for a cell reaction Thermodynamics: ΔG = ΔGº + RTlnQ ΔG = free energy, Q = reaction quotient so, -nFE = -nFEº + RTlnQ, or E = Eº– (RT/nF)lnQ more often seen as: E = Eº– (0.05916/n)logQ (although only valid at T = 298K) Note: in calculations, E is for reductions (even if oxidation actually occurs at that electrode) Equation for electrodes or full cells, although text uses Ecell = E+– E- where + and – refer to voltmeter leads

  10. ElectrochemistryThe Nernst Equation Example: Determine the voltage for aAg/AgCl electrode when [Cl-] = 0.010 M if Eº = 0.222 V (at T = 25°C)?

  11. ElectrochemistryApplications of The Nernst Equation Examples: The following electrode, Cd(s)|CdC2O4(s)|C2O42- is used to determine [C2O42-]. It is paired with a reference electrode that has an E value of 0.197 V (vs. the S.H.E.) with the reference electrode connected to the + end of the voltmeter. If Eº for the above reduction reaction is -0.522 V, and the measured voltage is 0.647 V, what is [C2O42-]?

  12. ElectrochemistryApplications of The Nernst Equation Application of Nernst Equation is most common in potentiometry In potentiometry measured voltage is related to log[x] (where x is the analyte) this provides a method to analyze analytes over a broad concentration range (e.g. pH electrodes function well from about pH 1 to pH 11 or over 10 orders of magnitude)

  13. ElectrochemistryApplications of The Nernst Equation Relating the Nernst Equation to Equilibrium Equations Example problem: It is desired to use the reaction Zn(CN)2(s) + 2e-↔ Zn(s) + 2CN- to measure [CN-] in suspected poisoned drinks. However, the Eº value is not available. Given that Eº = -0.762 V for Zn2+ + 2e-↔ Zn(s), and Ksp = 3.0 x 10-16 for Zn(CN)2(s) ↔ Zn2+ + 2CN-, calculate Eº for the first reaction.

  14. ElectrochemistryPotentiometry Overview (Chapter 14) Potentiometry is the use of measured voltages to provide chemical information Equipment Reference Electrode Indicator Electrode or ion-selective electrode Voltmeter Most Common Applications Measurement of specific ions (usually with ion-selective electrodes) Redox titrations (to keep track of the extent of a reaction)

  15. ElectrochemistryPotentiometry – Reference Electrodes Role of Reference Electrodes Provide other half-cell to complete circuit Designed so that the voltage is near constant (even when conditions change or when current occurs) Common Reference Electrodes silver/silver chloride: AgCl(s) + e-↔ Ag(s) + Cl- calomel (Hg2Cl2): Hg2Cl2(s) + 2e-↔ Hg(l) + 2Cl- Purpose of saturated Cl- conditions: less variability in [Cl-] as current forces reaction

  16. ElectrochemistryPotentiometry – Indicator Electrodes Metal (Reactive) Electrodes simple electrodes to measure dissolved metal use can be extended to anions (e.g. Cl- in Ag/AgCl electrode) fairly limited use Inert Electrodes e.g. Pt or graphite electrodes serve as an electron conduit to solution without electrode material participating in reaction used commonly in redox titrations described in Ch. 15 and in the types of electrolysis methods described in Ch. 16 Ion Selective Electrodes membrane based electrode to be described later Ag+ Fe3+ Fe2+ e- Ag(s) Ag(s) e- Pt(s)

  17. ElectrochemistryApplications of The Nernst Equation Application of Nernst Equation is most common in potentiometry In potentiometry measured voltage is related to log[x] (where x is the analyte) this provides a method to analyze analytes over a broad concentration range (e.g. pH electrodes function well from about pH 1 to pH 11 or over 10 orders of magnitude)

  18. ElectrochemistryApplications of The Nernst Equation Relating the Nernst Equation to Equilibrium Equations Example problem: It is desired to use the reaction Zn(CN)2(s) + 2e-↔ Zn(s) + 2CN- to measure [CN-] in suspected poisoned drinks. However, the Eº value is not available. Given that Eº = -0.762 V for Zn2+ + 2e-↔ Zn(s), and Ksp = 3.0 x 10-16 for Zn(CN)2(s) ↔ Zn2+ + 2CN-, calculate Eº for the first reaction.

  19. ElectrochemistryPotentiometry Overview (Chapter 14) Potentiometry is the use of measured voltages to provide chemical information Equipment Reference Electrode Indicator Electrode or ion-selective electrode Voltmeter Most Common Applications Measurement of specific ions (usually with ion-selective electrodes) Redox titrations (to keep track of the extent of a reaction)

  20. ElectrochemistryPotentiometry – Reference Electrodes Role of Reference Electrodes Provide other half-cell to complete circuit Designed so that the voltage is near constant (even when conditions change or when current occurs) Common Reference Electrodes silver/silver chloride: AgCl(s) + e-↔ Ag(s) + Cl- calomel (Hg2Cl2): Hg2Cl2(s) + 2e-↔ Hg(l) + 2Cl- Purpose of saturated Cl- conditions: less variability in [Cl-] as current forces reaction

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