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Section 10

Section 10. Electrochemical Cells and Electrode Potentials. Electrochemistry Oxidation/Reduction Reactions. “Redox” reactions involve electron transfer from one species to another Ox 1 + Red 2  Red 1 + Ox 2 Ox 1 + ne -  Red 1 ( Reduction ½ reaction)

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Section 10

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  1. Section 10 Electrochemical Cells and Electrode Potentials

  2. ElectrochemistryOxidation/Reduction Reactions • “Redox” reactions involve electron transfer from one species to another • Ox1 + Red2 Red1 + Ox2 • Ox1 + ne-  Red1 (Reduction ½ reaction) • Red2  Ox2 + ne-(Oxidation ½ reaction) • “Reducing agent” donates electrons (is oxidezed) • “Oxidizing agent” accepts electrons (is reduced)

  3. ElectrochemistryOxidation/Reduction Reactions • Typical oxidizing agents:Standard Potentials,V • O2 + 4H+ + 4e- 2H2O +1.229 • Ce4+ + e-  Ce3+ +1.6 (acid) • MnO4- + 8H+ + 5e-  Mn2+ + 4H2O +1.51 • Typical reducing agents: • Zn2+ + 2e-  Zno -0.763 • Cr3+ + e-  Cr2+ -0.408 • Na+ + e-  Nao -2.714

  4. The salt bridge allows charge transfer through the solution and prevents mixing. The spontaneous cell reaction (Fe2+ + Ce4+ = Fe3+ + Ce4+) generates the cell potential. The cell potential depends on the half-reaction potentials at each electrode. The Nernst equation describes the concentration dependence. A battery is a voltaic cell. It goes dead when the reaction is complete (Ecell = 0). Fig. 12.1. Voltaic cell. ©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

  5. ElectrochemistryStandard Reduction Potentials • Half-Reaction Potentials: • They are measured relative to each other • Reference reduction half-reaction: • standard hydrogen electrode (SHE) • normal hydrogen electrode (NHE) • 2H+(a=1.0) + 2e- H2(g 1atm) 0.0000 volts

  6. The more positive the Eo, the better oxidizing agent is the oxidized form (e.g., MnO4-). The more negative the Eo, the better reducing agent is the reduced form (e.g., Zn). ©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

  7. ElectrochemistryReduction Potentials • General Conclusions: • 1. The more positive the electrode potential, the stronger an oxidizing agent the oxidized form is and the weaker a reducing agent the reduced form is • 2. The more negative the reduction potential, the weaker the oxidizing agent is the oxidized formis and the stronger the reducing agent the reduced form is.

  8. ElectrochemistryOxidation/Reduction Reactions • Typical oxidizing agents:Standard Potentials,V • O2 + 4H+ + 4e- 2H2O +1.229 • Ce4+ + e-  Ce3+ +1.6 (acid) • MnO4- + 8H+ + 5e-  Mn2+ + 4H2O +1.51 • Typical reducing agents: • Zn2+ + 2e-  Zno -0.763 • Cr3+ + e-  Cr2+ -0.408 • Na+ + e-  Nao -2.714

  9. ElectrochemistryOxidation/Reduction Reactions • Net Redox Reactions:Standard Potentials,V • MnO4- Mn2+ • MnO4- + 8H+ + 5e-  Mn2+ + 4H2O +1.51 • Sn4+ + 2e-  Sn2+ +0.154 • Balanced Net Ionic Reaction: • 2MnO4- + 16H+ + 5Sn2+  2Mn2+ + 5Sn4+ + 8H2O

  10. ElectrochemistryVoltaic Cell • The spontaneous (Voltaic) cell reaction is the one that gives a positive cell voltage when subtracting one half-reaction from the other. • Eocell = Eoright– Eoleft = Eocathode – Eoanode =Eo+ - Eo- • Which is the Anode? The Cathode? • Convention: • The anode is the electrode where oxidation occurs  the more negative half-reaction potential • The cathode is the electrode where reduction occurs  the more positive half-reaction potential • anode  solution  cathode

  11. ElectrochemistryOxidation/Reduction Reactions • Net Redox Reactions:Standard Potentials,V • MnO4- Mn2+ • MnO4- + 8H+ + 5e-  Mn2+ + 4H2O +1.51 • Sn4+ + 2e-  Sn2+ +0.154 • Balanced Net Ionic Reaction: • 2MnO4- + 16H+ + 5Sn2+  2Mn2+ + 5Sn4+ + 8H2O • Eocell = Eocat – Eoan = (+1.51 – (+0.154)) = +1.36 V

  12. ElectrochemistryNernst Equation • Effects of Concentrations on Potentials: • aOx + ne- bRed • E = Eo – (2.3026RT/nF) log([Red]b/[Ox]a • Where E is the reduction at specific conc., • Eo is standard reduction potential, n is number of electrons involved in the half reaction, • R is the gas constant (8.3143 V coul deg-1mol-1), • T is absolute temperature, • and F is the Faraday constant (96487 coul eq-1). • At 25oC(298.16K) the value of 2.3026RT/F is 0.05916 • Note: Concentrations should be activities

  13. Electrochemistry • Calculations: • MnO4- + 8H+ + 5e-  Mn2+ + 4H2O Eo = +1.51 V • For [H+] = 1.0M, [MnO4-] = 0.10M, [Mn2+] = 0.010M • E = Eo – 0.05916/5 (log ([Mn2+]/[MnO4-][H+]8) • E = +1.51 – 0.1183(-1) = +1.63 V vs NHE • Note: This is more positive than Eo • Greater tendency to be reduced compared to standard conditions.

  14. Electrochemistry • Calculations: • Silver electrode/silver chloride deposit/0.010M NaCl • AgCl + 1e- Ago + Cl- E = ? • Ag+ + 1e-  Ago Eo = +0.799 V • AgCl  Ag+ + Cl- Ksp= 1.8 x 10-10 • AgCl + e-  Ago + Cl- • E = Eo - (0.05916/1) Log (1/[Ag+]) • [Ag+] = Ksp/[Cl-] = 1.8 x 10-10/(0.010) = 1.8 x 10-8 • E = +0.799 – (0.05916)(7.74) = +0.341 V

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