2. Electric energy Chemical energy Electrolysis Galvanic cell Chapter 8 Electrochemistry

3. Electrochemistry Study of chemical reactions that can produce electricity or use electricity to produce desired product. Study of interchange of chemical and electrical energy Electrochemical reaction always involves oxidation-reduction reactions – Electron transfer reactions – Electrons transferred from one substance to another Also called redox reactions 2

4. Galvanic Cells

5. The Daniell Cell Flow of Zn 2+ Flow of SO 4 2- Half-cell Half-cell reaction

6. 5 Needed to complete circuit Tube filled with solution of an electrolyte – Salt composed of ions not involved in cell reaction – KNO 3 and KCl often used Porous plugs at each end of tube – Prevent solution from pouring out – Enable ions from salt bridge to migrate between half- cells to neutralize charges in cell compartments Anions always migrate toward anode Cations always migrate toward cathode Salt Bridge Salt bridge

7. Anode compartment Half-cell Half-reaction: Oxidation Cathode compartment Half-cell Half-reaction: Reduction Zn (s) → Zn 2+ (aq) +2e - Cu 2+ (aq) + 2e - → Cu (s) Cu 2+ (aq) + Zn (s) → Cu (s) + Zn 2+ (aq) total cell reaction 2e -

8. 7 Cell Notation Cu (s) |Cu 2+ (aq) ||Ag + (aq) |Ag (s) Single slash = boundary between phases (solid electrode and aqueous solution of ions) Double slash represents salt bridge – Separates cell reactions In each half (half-cell) – Electrodes appear at outsides – Reaction electrolytes in inner section – Species in same state separated with ; – Concentrations shown in ( ) anodecathode anode electrode anode electrolyte cathode electrolyte cathode electrode Salt Bridge

9. 8 Learning Check Write the standard cell notation for the following electrochemical cells: Fe (s) + Cd 2+ (aq)  Cd (s) + Fe 2+ (aq) Anode = ox = Fe (s) Cathode = red = Cd 2+ (aq) Fe (s) |Fe 2+ (aq) ||Cd 2+ (aq) |Cd (s) Al (s) + Au 3+ (aq)  Al 3+ (aq) + Au (s) Anode = ox = Al (s) Cathode = red = Au 3+ (aq) Al (s) |Al 3+ (aq) ||Au 3+ (aq) |Au (s)

10. Your Turn! Write the standard cell notation (Pt electrodes) for the following reaction: 2Mn 3+ (aq) + 2I - (aq) → Mn 2+ (aq) + I 2 (s) A. Pt(s)|Mn 3+ (aq); Mn 2+ (aq)||I - (aq)|I 2 (s)|Pt(s) B. Pt(s)|I - (aq)|I 2 (s)||Mn 3+ (aq); Mn 2+ (aq)|Pt(s) C. Mn 3+ (aq)|Pt(s); Mn 2+ (aq)||I - (aq)|I 2 (s)|Pt(s) D. Pt(s)|Mn 3+ (aq); I - (aq)||Mn 2+ (aq)|I 2 (s)|Pt(s) Oxidation reaction is on the right and reduction reaction is on the left of the salt bridge (||). 9

11. Reaction can be performed without harnessing electricity!  G of reaction: maximum work over and above volume work (electricity) that can be harnessed from the chemical reaction.

12. Cu wire is dipped into Zn 2+ solution: nothing happens. Cu wire is dipped into Ag + solution: Ag + has higher tendency to be reduced than Cu 2+. Cu 2+ has higher tendency to be reduced than Zn 2+. 2Ag + (aq) + Cu (s) → 2Ag (s) + Cu 2+ (aq)

13. Work harnessed!!

14. Ag + has higher tendency to be reduced than Cu 2+. Cu 2+ has higher tendency to be reduced than Zn 2+. How can we know? Electrode Potential Reflects tendency towards reduction Problem: Only potential difference can be measured between two half-cells.

15. Hydrogen Standard Electrode This electrode used as standard. EMF of all other electrodes measured with reference to this electrode. H 2(g) → 2 H + (aq) +2e -

16. Standard Reduction Potential SHE as anode (Oxidation). The other electrode cathode (Reduction). H 2(g) → 2 H + (aq) +2e - Cu 2+ (aq) + 2e - → Cu (s)

17. Voltage: Potential difference Measured voltage = Potential of reduction electrode - Potential of anode electrode

18. H 2(g) → 2 H + (aq) +2e - Cu 2+ (aq) + 2e - → Cu (s) Cu 2+ (aq) + H 2(g) → Cu (s) + 2H + (aq) Spontaneous at Standard conditions Reduction Potential Cu 2+ (aq) + 2e - → Cu (s) E o red =0.34 V

19. H 2(g) → 2 H + (aq) +2e - Zn 2+ (aq) + 2e - → Zn (s) Zn 2+ (aq) + H 2(g) → Zn (s) + 2H + (aq) Zn (s) → Zn 2+ (aq) +2e - 2 H + (aq) +2e - → H 2(g) Zn (s) + 2H + (aq) → Zn 2+ (aq) + H 2(g)

20. Reduction Potential Zn 2+ (aq) + H 2(g) → Zn (s) + 2H + (aq) nonspontaneous Electricity must be applied to force this process to take place! Zn 2+ (aq) + 2e - → Zn (s) E o red = -0.76 V

22. Cu 2+ (aq) + 2e - → Cu (s) Zn (s) → Zn 2+ (aq) +2e - Cu 2+ (aq) + Zn (s) → Cu (s) + Zn 2+ (aq)

24. Electrochemical thermodynamics  G of reaction: maximum work over and above volume work (available work) (electricity) that can be harnessed from the chemical reaction.

25. Electrical heater Heating elements resistive Work done

35. Concentration Cells Consider the cell presented on the left. The 1/2 cell reactions are the same, it is just the concentrations that differ. Will there be electron flow?

36. Concentration Cells (cont.) AgAg + + e - -E 1/2 Anode: Ag + + e - Ag E 1/2 Cathode: E cell = E° cell - (0.0591/n)log(Q) 0 V E cell = - (0.0591)log(0.1) = 0.0591 V 1

37. Concentration Cells (cont.) Another Example: What is E cell ?

38. Concentration Cells (cont.) E cell = E° cell - (0.0591/n)log(Q) 0 Fe 2+ + 2e - Fe 2 e - transferred…n = 2 2 E cell = -(0.0296)log(.1) = 0.0296 V anodecathode e-e-