1. ELECTROCHEMISTRY REDOX REVISITED! 24-Nov-97Electrochemistry (Ch. 21) & Phosphorus and Sulfur (ch 22)1

2. 24-Nov-97Electrochemistry (Ch. 21) & Phosphorus and Sulfur (ch 22) 2

3. 3 ELECTROCHEMISTRY redox reactions electrochemical cells electrode processes construction notation cell potential and  G o standard reduction potentials (E o ) non-equilibrium conditions (Q) batteries corrosion Electric automobile

4. 4 CHEMICAL CHANGE  ELECTRIC CURRENT With time, Cu plates out onto Zn metal strip, and Zn strip “disappears.” Zn is oxidized and is the reducing agent Zn(s)  Zn 2+ (aq) + 2e- Cu 2+ is reduced and is the oxidizing agent Cu 2+ (aq) + 2e-  Cu(s) http://www.youtube.com/wat ch?v=0oSqPDD2rMA

5. 5 Electrons travel thru external wire. Salt bridge allows anions and cations to move between electrode compartments. This maintains electrical neutrality. ANODE OXIDATION CATHODE REDUCTION

6. 6 This is the STANDARD CELL POTENTIAL, E o E o is a quantitative measure of the tendency of reactants to proceed to products when all are in their standard states at 25  C. CELL POTENTIAL, E o For Zn/Cu, voltage is 1.10 V at 25  C and when [Zn 2+ ] and [Cu 2+ ] = 1.0 M.

7. 7 E o and  G o E o is related to  G o, the free energy change for the reaction.  G o = - n F E o F = Faraday constant = 9.6485 x 10 4 J/Vmol n = the number of moles of electrons transferred. Michael Faraday 1791-1867 Discoverer of electrolysis magnetic props. of matter electromagnetic induction benzene and other organic chemicals n for Zn/Cu cell ? n = 2 Zn / Zn 2+ // Cu 2+ / Cu

8. 8 For a reactant-favored reaction - electrolysis cell: Electric current  chemistry Reactants  Products  G o > 0 and so E o < 0 (E o is negative) For a product-favored reaction – battery or voltaic cell: Chemistry  electric current Reactants  Products  G o 0 (E o is positive) E o and  G o (2)  G o = - n F E o

9. 9 STANDARD CELL POTENTIALS, E o Can’t measure half- reaction E o directly. Therefore, measure it relative to a standard HALF CELL: the S tandard H ydrogen E lectrode ( SHE ). 2 H + (aq, 1 M) + 2e- H 2 (g, 1 atm) E o = 0.0 V

10. 10 BEST Reducing agent ? ? STANDARD REDUCTION POTENTIALS Half-Reaction E o (Volts) Cu 2+ + 2e-  Cu+ 0.34 Oxidizing ability of ion Reducing ability of element 2 H + + 2e-  H 2 0.00 Zn 2+ + 2e-  Zn -0.76 BEST Oxidizing agent ? ? Cu 2+ Zn

11. 11 Using Standard Potentials, E o H 2 O 2 /H 2 O +1.77 Cl 2 /Cl - +1.36 O 2 /H 2 O +1.23 In which direction does the following reaction go? Cu(s) + 2 Ag + (aq)  Cu 2+ (aq) + 2 Ag(s) Hg 2+ /Hg +0.86 Sn 2+ /Sn -0.14 Al 3+ /Al -1.66 Ag + /Ag +0.80 Cu 2+ / Cu +0.34 See Table 21.1, App. J for E o (red.) Which is the best oxidizing agent: O 2, H 2 O 2, or Cl 2 ? Which is the best reducing agent: Sn, Hg, or Al ? As written: E o = (-0.34) + 0.80 = +0.43 V reverse rxn: E o = +0.34 + (-0.80) = -0.43 V

12. 12 Cells at Non-standard Conditions For ANY REDOX reaction, Standard Reduction Potentials allow prediction of direction of spontaneous reaction If E o > 0 reaction proceeds to RIGHT (products) If E o < 0 reaction proceeds to LEFT (reactants) E o only applies to [ ] = 1 M for all aqueous species at other concentrations, the cell potential differs E cell can be predicted by Nernst equation

13. 13 Cells at Non-standard Conditions (2) E o only applies to [ ] = 1 M for all aqueous species at other concentrations, the cell potential differs E cell can be predicted by Nernst equation E = E o - ln (Q) RT nF n = # e- transferred F = Faraday’s constant = 9.6485 x 10 4 J/Vmol Q is the REACTION QUOTIENT (recall ch. 16, 20) At equilibrium  G = 0 E = 0 Q = K  G o, E o refer to ALL REACTANTS relative to ALL PRODUCTS

14. 14 Example of Nernst Equation Q. Determine the potential of a Daniels cell with [Zn 2+ ] = 0.5 M and [Cu 2+ ] = 2.0 M; E o = 1.10 V A. Zn / Zn 2+ (0.5 M) // Cu 2+ (2.0 M) / Cu E = 1.10 - (0.0257) ln ( [Zn 2+ ]/[Cu 2+ ] ) 2 E = 1.10 - (-0.018) = 1.118 V Zn(s) + Cu 2+ (aq)  Zn 2+ (aq) + Cu(s) Q = ? [Zn 2+ ] [Cu 2+ ] E = E o - ln (Q) RT nF

15. 15 Nernst Equation (2) Q. What is the cell potential and the [Zn 2+ ], [Cu 2+ ] when the cell is completely discharged? A. When cell is fully discharged: chemical reaction is at equilibrium E = 0  G = 0 Q = Kand thus 0 = E o - (RT/nF) ln (K) or E o = (RT/nF) ln (K) or ln (K) = nFE o /RT = (n/0.0257) E o at T = 298 K So... K = e = 1.5 x 10 37 (2)(1.10)/(.0257) E = E o - ln (Q) RT nF Determine K c from E o by K c = e (nFE o /RT)

16. 16 Primary (storage) Batteries Anode (-) Zn  Zn 2+ + 2e- Cathode (+) 2 NH 4 + + 2e-  2 NH 3 + H 2 Common dry cell (LeClanché Cell) Mercury Battery (calculators etc) Anode (-) Zn (s) + 2 OH - (aq)  ZnO (s) + 2H 2 O + 2e- Cathode (+) HgO (s) + H 2 O + 2e-  Hg (l) + 2 OH - (aq)

17. 17 Secondary (rechargeable) Batteries Nickel-Cadmium 11_NiCd.mov 21m08an5.mov Anode (-) Cd + 2 OH - Cd(OH) 2 + 2e-  Cathode (+) NiO(OH) + H 2 O + e- Ni(OH) 2 + OH -  DISCHARGE RE-CHARGE

18. 18 Secondary (rechargeable) Batteries (2) Lead Storage Battery 11_Pbacid.mov 21mo8an4.mov Con-proportionation reaction - same species produced at anode and cathode RECHARGEABLE Cathode (+) E o = +1.68 V PbO 2 (s) + HSO 4 - + 3 H + + 2e- PbSO 4 (s) + 2 H 2 O  Overall battery voltage = 6 x (0.36 + 1.68) = 12.24 V Anode (-) E o = +0.36 V Pb (s) + HSO 4 -  PbSO 4 (s) + H + + 2e-

19. 19 Corrosion - an electrochemical reaction Electrochemical or redox reactions are tremendously damaging to modern society e.g. - rusting of cars, etc: anode: Fe -  Fe 2+ + 2 e- net: 2 Fe(s) + O 2 (g) + 2 H 2 O (l)  2 Fe(OH) 2 (s) Mechanisms for minimizing corrosion sacrificial anodes (cathodic protection) (e.g. Mg) coatings - e.g. galvanized steel - Zn layer forms (Zn(OH) 2.xZnCO 3 ) this is INERT (like Al 2 O 3 ); if breaks, Zn is sacrificial cathode: O 2 + 2 H 2 O + 4 e -  4 OH - E OX = +0.44 E RED = +0.40 E cell = +0.84

20. 20 Electrolysis of Aqueous NaOH Anode : E o = -0.40 V 4 OH -  O 2 (g) + 2 H 2 O + 2e- Cathode : E o = -0.83 V 4 H 2 O + 4e-  2 H 2 + 4 OH - E o for cell = -1.23 V since E o 0 - not spontaneous ! - ONLY occurs if E external > 1.23 V is applied Electric Energy  Chemical Change 11_electrolysis.mov 21m10vd1.mov

21. Go to Molecular Workbench and find the “other activites”. Select from the top list “How a battery works” and do all sections. You will not regret it! electrochemical cell animation : http://www.youtube.com/watch?v=A0VUsoeT 9aM http://www.youtube.com/watch?v=A0VUsoeT 9aM 24-Nov-97Electrochemistry (Ch. 21) & Phosphorus and Sulfur (ch 22) 21