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Galvanic Cells

Galvanic Cells. anode oxidation. cathode reduction. spontaneous redox reaction. 2 e - + 2H + (1 M ) H 2 (1 atm ). Zn ( s ) Zn 2+ (1 M ) + 2 e -. Zn ( s ) + 2H + (1 M ) Zn 2+ (1 M ) + H 2 (1 atm ). Standard Reduction Potentials.

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Galvanic Cells

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  1. Galvanic Cells anode oxidation cathode reduction spontaneous redox reaction

  2. 2e- + 2H+ (1 M) H2 (1 atm) Zn (s) Zn2+ (1 M) + 2e- Zn (s) + 2H+ (1 M) Zn2+ (1 M) + H2 (1 atm) Standard Reduction Potentials Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s) Anode (oxidation): Cathode (reduction):

  3. Zn Cu

  4. Brief Activity Series

  5. E0 is for the reaction as written • The more positive E0 the greater the tendency for the substance to be reduced • The half-cell reactions are reversible • The sign of E0 changes when the reaction is reversed • Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E0

  6. How many moles of Cu can be plated out of a solution containing Cu2+ ions if a 1.50 amp current is passed through for 300 s? RXN: Cu2+ + 2e-→ Cu(s) Notes: 1 amp = C/s F = 9.65 x 104 C/ mole e- Find: C → moles e-s → moles Cu 300 s mole e- 9.65 x 104 C 1 mole Cu 2 mole e-s 1.50 C s = moles Cu 2.33 x 10-3

  7. Effect of Concentration on EMF A Battery going dead. ∆G = ∆Go + RT lnQ -nFE = -nFEo + RT ln Q Dividing by –nF gives Rise to the Nerst Equation.

  8. The Nerst Equation: Q: aA + bB ↔ cC + dD n = # of e- s transferred

  9. The Nerst Equation allows us to find voltage under nonstandard conditions Example: Determine E (voltage) for: Fe(s) + Cd2+(aq)→ Fe2+(aq) + Cd(s) When [Fe2+] = 0.10 M and [Cd2+] = 1.0 M @ 298K Half RXNs: Fe(s) → Fe2+ + 2e- Eo = +0.44 V Cd2+ + 2e-→ Cd(s) Eo = -0.40 V Eo = +0.04 V + Value means spontaneous

  10. Now try: [Fe2+] = 1.0 M and [Cd2+] = 0.01 M @ 298K - Value means non spontaneous and will go in opposite direction

  11. What does an E = 0 value mean? For: Fe(s) + Cd2+(aq) → Fe2+(aq) + Cd(s) Eo = +0.04V K = 22

  12. Concentration Cells Galvanic cell from two half-cells composed of the same material but differing in ion concentrations.

  13. Batteries Pb (s) + SO2- (aq) PbSO4 (s) + 2e- 4 PbO2(s) + 4H+(aq) + SO2-(aq) + 2e- PbSO4(s) + 2H2O (l) 4 Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO2- (aq) 2PbSO4 (s) + 2H2O (l) 4 Lead storage battery Anode: Cathode:

  14. Batteries Zn (s) Zn2+ (aq) + 2e- + 2NH4(aq) + 2MnO2(s) + 2e- Mn2O3(s) + 2NH3(aq) + H2O (l) Zn (s) + 2NH4+ (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3(s) Dry cell Leclanché cell Anode: Cathode:

  15. Batteries Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e- HgO (s) + H2O (l) + 2e- Hg (l) + 2OH-(aq) Zn(Hg) + HgO (s) ZnO (s) + Hg (l) Mercury Battery Anode: Cathode:

  16. Batteries Solid State Lithium Battery

  17. Batteries 2H2 (g) + 4OH- (aq) 4H2O (l) + 4e- O2(g) + 2H2O (l) + 4e- 4OH-(aq) 2H2 (g) + O2 (g) 2H2O (l) A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning Anode: Cathode:

  18. Cathodic Protection of an Iron Storage Tank

  19. Electrolysis is the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur. Electrolysis of molten NaCl

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