1. Section 10 Electrochemical Cells and Electrode Potentials

2. 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) Red 2  Ox 2 + ne - (Oxidation ½ reaction) “Reducing agent” donates electrons (is oxidezed) “Oxidizing agent” accepts electrons (is reduced)

3. Electrochemistry Oxidation/Reduction Reactions Typical oxidizing agents: Standard Potentials,V –O 2 + 4H + + 4e -  2H 2 O+1.229 –Ce 4+ + e -  Ce 3+ +1.6(acid) –MnO 4 - + 8H + + 5e -  Mn 2+ + 4H 2 O+1.51 Typical reducing agents: –Zn 2+ + 2e -  Zn o -0.763 –Cr 3+ + e -  Cr 2+ -0.408 –Na + + e -  Na o -2.714

4. Fig. 12.1. Voltaic cell. The salt bridge allows charge transfer through the solution and prevents mixing. The spontaneous cell reaction (Fe 2+ + Ce 4+ = Fe 3+ + Ce 4+ ) 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 (E cell = 0). The salt bridge allows charge transfer through the solution and prevents mixing. The spontaneous cell reaction (Fe 2+ + Ce 4+ = Fe 3+ + Ce 4+ ) 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 (E cell = 0). ©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

5. Electrochemistry Standard 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 + (  1.0) + 2e -  H 2(g 1atm) 0.0000 volts

6. The more positive the E o, the better oxidizing agent is the oxidized form (e.g., MnO 4 - ). The more negative the E o, the better reducing agent is the reduced form (e.g., Zn). The more positive the E o, the better oxidizing agent is the oxidized form (e.g., MnO 4 - ). The more negative the E o, the better reducing agent is the reduced form (e.g., Zn). ©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)

7. Electrochemistry Reduction 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. Electrochemistry Oxidation/Reduction Reactions Typical oxidizing agents: Standard Potentials,V –O 2 + 4H + + 4e -  2H 2 O+1.229 –Ce 4+ + e -  Ce 3+ +1.6(acid) –MnO 4 - + 8H + + 5e -  Mn 2+ + 4H 2 O+1.51 Typical reducing agents: –Zn 2+ + 2e -  Zn o -0.763 –Cr 3+ + e -  Cr 2+ -0.408 –Na + + e -  Na o -2.714

9. Electrochemistry Oxidation/Reduction Reactions Net Redox Reactions: Standard Potentials,V MnO 4 -  Mn 2+ MnO 4 - + 8H + + 5e -  Mn 2+ + 4H 2 O+1.51 Sn 4+ + 2e -  Sn 2+ +0.154 Balanced Net Ionic Reaction: 2MnO 4 - + 16H + + 5Sn 2+  2Mn 2+ + 5Sn 4+ + 8H2O

10. Electrochemistry Voltaic Cell The spontaneous (Voltaic) cell reaction is the one that gives a positive cell voltage when subtracting one half- reaction from the other. E o cell = E o right – E o left = E o cathode – E o anode =E o + - E o - 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. Electrochemistry Oxidation/Reduction Reactions Net Redox Reactions: Standard Potentials,V MnO 4 -  Mn 2+ MnO 4 - + 8H + + 5e -  Mn 2+ + 4H 2 O+1.51 Sn 4+ + 2e -  Sn 2+ +0.154 Balanced Net Ionic Reaction: 2MnO 4 - + 16H + + 5Sn 2+  2Mn 2+ + 5Sn 4+ + 8H2O E o cell = E o cat – E o an = (+1.51 – (+0.154)) = +1.36 V

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

13. Electrochemistry Calculations: MnO 4 - + 8H + + 5e -  Mn 2+ + 4H 2 O E o = +1.51 V For [H + ] = 1.0M, [MnO 4 - ] = 0.10M, [Mn 2+ ] = 0.010M E = E o – 0.05916/5 (log ([Mn 2+ ]/[MnO 4 - ][H + ] 8 ) E = +1.51 – 0.1183(-1) = +1.63 V vs NHE Note: This is more positive than E o Greater tendency to be reduced compared to standard conditions.

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