1. 19.2 Galvanic Cells 19.3 Standard Reduction Potentials 19.4 Spontaneity of Redox Reactions 19.5 The Effect of Concentration on Emf 19.8 Electrolysis Chapter 19 Electrochemistry Semester 2/2013 Ref: http://www.mhhe.com/chemistry/chang

2. 19.2 Galvanic Cells Spontaneous(natural) redox reaction anode oxidation cathode reduction

3. Cell = half-cell + half – cell OxidationReduction AnodeCathode In Galvanic cell … Zn (s) + Cu 2+ (aq) Cu (s) + Zn 2+ (aq) Zn is oxidized to Zn 2+ ion  Zn electrode is Anode (Reducing Agent) Cu 2+ is reduced to Cu  Cu electrode is Cathode (Oxidizing Agent)

4. Galvanic Cells The difference in electrical potential between the anode and cathode is called: cell voltage electromotive force (emf) cell potential Cell Diagram Cell Equation Zn (s) + Cu 2+ (aq) Cu (s) + Zn 2+ (aq) [Cu 2+ ] = 1 M & [Zn 2+ ] = 1 M Cell Notation Zn (s) | Zn 2+ (1 M) || Cu 2+ (1 M) | Cu (s) anodecathode

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

6. 19.3 Standard Reduction Potentials Standard reduction potential (E 0 ) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm. E 0 = 0 V Standard hydrogen electrode (SHE) 2e - + 2H + (1 M) H 2 (1 atm) Reduction Reaction

7. E 0 = 0.76 V cell Standard emf (E 0 ) cell 0.76 V = 0 - E Zn /Zn 0 2+ E Zn /Zn = -0.76 V 0 2+ Zn 2+ (1 M) + 2e - Zn E 0 = -0.76 V E 0 = E H /H - E Zn /Zn cell 00 + 2+ 2 Standard Electrode Potentials E 0 = E cathode - E anode cell 00 Zn (s) | Zn 2+ (1 M) || H + (1 M) | H 2 (1 atm) | Pt (s)

8. Standard Electrode Potentials Pt (s) | H 2 (1 atm) | H + (1 M) || Cu 2+ (1 M) | Cu (s) 2e - + Cu 2+ (1 M) Cu (s) H 2 (1 atm) 2H + (1 M) + 2e - Anode (oxidation): Cathode (reduction): H 2 (1 atm) + Cu 2+ (1 M) Cu (s) + 2H + (1 M) E 0 = E cathode - E anode cell 00 E 0 = 0.34 V cell E cell = E Cu /Cu – E H /H 2++ 2 000 0.34 = E Cu /Cu - 0 0 2+ E Cu /Cu = 0.34 V 2+ 0

9. Note: The more positive E 0 the greater the tendency for the substance to be reduced The half-cell reactions are reversible The sign of E 0 changes when the reaction is reversed Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E 0

10. 19.4 Spontaneity of Redox Reactions  G = -nFE cell  G 0 = -nFE cell 0 n = number of moles of electrons in reaction F = 96,500 J V mol = 96,500 C/mol  G 0 = -RT ln K = -nFE cell 0 E cell 0 = RT nF ln K (8.314 J/K mol)(298 K) n (96,500 J/V mol) ln K = = 0.0257 V n ln K E cell 0 = 0.0592 V n log K E cell 0 E 0 cell > 0 spontaneous reaction

11. Spontaneity of Redox Reactions

12. 19.5 The Effect of Concentration on Cell Emf  G =  G 0 + RT ln Q  G = -nFE  G 0 = -nFE 0 -nFE = -nFE 0 + RT ln Q E = E 0 - ln Q RT nF Nernst equation At 298 K ln = 2.303log - 0.0257 V n ln Q E 0 E = - 0.0592 V n log Q E 0 E =

13. 19.8 Electrolysis is the process in which electrical energy is used to cause a non spontaneous chemical reaction to occur.

14. Electrolysis of Water 19.8

15. Electrolysis and Mass Changes Quantitative Aspects Case (i) Na + + 1eNa 1 mol. of electron produces 1 mol of Na Atom(22g) 1 F (96500 C) Case (ii) Mg 2+ + 2eMg 2 mol. of electron produces 1 mol of Mg Atom(24g) 2 F (2x 96500C) Case (iii) Al 3 + + 3eAl 3 mol. of electron produces 1 mol of Al Atom(26g) 3 F (3 x 96500 C)

16. charge ( C ) = current (A) x time (s) 1 mole of electron = 96500 coulomb 1 mol. of Na atom = 22 g 1 mol. of Mg atom = 24 g 1 mol. of Al atom = 26 g