1. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Chemistry FIFTH EDITION by Steven S. Zumdahl University of Illinois

2. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2 Chemistry FIFTH EDITION Chapter 17 Electrochemistry

3. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 3 Electrochemistry The study of the interchange of chemical and electrical energy. Processes involve Oxidation-Reduction Reactions.

4. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 4 Primarily Concerned with Two Processes 1) Galvanic Cells: Generation of an electric current from a spontaneous oxidation-reduction chemical reaction. 2) Electrolysis (Opposite Process): The use of a current to produce a chemical change.

5. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 5 Review of Terms (Chapter 4) oxidation-reduction (redox) reaction: involves a transfer of electrons from the reducing agent to the oxidizing agent. oxidation: loss of electrons reduction: gain of electrons

6. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 6 Figure 4.19 A Summary of an Oxidation- Reduction Process

7. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 7 Half-Reactions The overall reaction is split into two half-reactions  one involving oxidation and one reduction. 8H + + MnO 4  + 5Fe 2+  Mn 2+ + 5Fe 3+ + 4H 2 O Reduction: 8H + + MnO 4  + 5e   Mn 2+ + 4H 2 O Oxidation: 5Fe 2+  5Fe 3+ + 5e 

8. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 8 Rules for Assigning Oxidation States (p164-171) 1. Oxidation state of an atom in an element = 0 2. Oxidation state of monatomic ion = charge 3. Oxygen =  2 in covalent compounds (except in peroxides where it =  1) 4. H = +1 in covalent compounds 5. Fluorine =  1 in compounds 6. Sum of oxidation states = 0 in compounds Sum of oxidation states = charge of the ion

9. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 9 Balancing by Half-Reaction Method (in Acidic Solution) 1.Write separate reduction, oxidation reactions. 2.For each half-reaction:  Balance elements (except H, O)  Balance O using H 2 O  Balance H using H +  Balance charge using electrons

10. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 10 Balancing by Half-Reaction Method (continued) 3.If necessary, multiply by integer to equalize electron count. 4.Add half-reactions. 5.Check that elements and charges are balanced.

11. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 11 Half-Reaction Method - Balancing in Base 1.Balance as in acid. 2.Add OH  that equals H + ions (both sides!) 3.Form water by combining H +, OH . 4.Check elements and charges for balance.

12. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 12 Half-Reactions The overall reaction is split into two half- reactions, one involving oxidation and one reduction. 8H + + MnO 4  + 5Fe 2+  Mn 2+ + 5Fe 3+ + 4H 2 O Reduction: 8H + + MnO 4  + 5e   Mn 2+ + 4H 2 O Oxidation: 5Fe 2+  5Fe 3+ + 5e 

13. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 13 That is, Fe 2+ will lose e - (transfer electrons) to MnO 4 - which will gain the e -. If Fe 2+ and MnO 4 - are in the same solution, electrons will transfer directly when the reactants collide.  No Useful Work!!

14. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 14 Therefore, one needs to separate the oxidizing agent and the reducing agent. Require the electron transfer to occur through a wire. The current produced in the wire by the electron flow can be directed through a device to provide useful work.

15. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 15 Figure 17.1 Schematic of a Method to Separate the Oxidizing and Reducing Agents of a Redox Reaction Electrons should flow through the wire from Fe 2+ to MnO 4 1-.

16. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 16 Electrons should flow through the wire from Fe 2+ to MnO 4 -. BUT------- Current flows for an instant & then CEASES! This in initial flow of electrons leads to a charge buildup, charge separation.

17. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 17 Left compartment: receives e - ’s becomes negatively charged. Right compartment: loses e - ’s becomes positively charged. Sustained electron flow will not occur under these conditions.

18. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 18 Solve this problem by connecting compartments: Use Salt Bridge or Porous disk. Figure 17.2 Galvanic Cells

19. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 19 Salt Bridge: U-tube filled with a strong electrolyte. Porous Disk: Disk in a tube connecting the two solutions. It has tiny passages that allow hindered flow of ions. Both devices allow ion flow without extensive mixing of the solutions.

20. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 20 Electrons flow through the wire from the reducing agent to the oxidizing agent. Ions flow from one compartment to the other to keep the net charge zero.

21. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 21 Galvanic Cell A device in which chemical energy is changed to electrical energy.

22. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 22 The Reaction takes place at the interface between the electrode and the solution where the electron transfer occurs.

23. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 23 Electrodes Anode and Cathode OXIDATION occurs at the ANODE. REDUCTION occurs at the CATHODE.

24. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 24 Animated Galvanic Cell http://www.chembio.uoguelph.ca/educmat/ch m19105/galvanic/galvanic1.htmhttp://www.chembio.uoguelph.ca/educmat/ch m19105/galvanic/galvanic1.htm http://www.chem.iastate.edu/group/Greenbow e/sections/projectfolder/flashfiles/electroChem /voltaicCellEMF.htmlhttp://www.chem.iastate.edu/group/Greenbow e/sections/projectfolder/flashfiles/electroChem /voltaicCellEMF.html

25. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 25 Cell Potential Cell Potential or Electromotive Force (emf): The “pull” or driving force on the electrons.

26. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 26 Figure 17.3 An Electrochemical Process Involves Electron Transfer at the Interface Between the Electrode and the Solution Pulls e - through wire From reducing agent.

27. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 27 Figure 17.3 An Electrochemical Process Involves Electron Transfer at the Interface Between the Electrode and the Solution Supplies e - which travel through the wire to cathode

28. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 28 Volt (V): Unit of electrical potential 1 V = 1 J/C 1V = 1 joule of work per coulomb of charge transferred

29. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 29 Measuring Cell Potential 1) Traditional Voltmeter Some useful work of the cell is lost through fictional heating as the current flows through the wire. 2) Potentiometer: This device applies voltage in opposition to the cell potential. It is adjusted until current flow in the cell circuit stops. Value on potentiometer is equal in magnitude but opposite in sign. 3) Digital Voltmeter: See next slide.

30. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 30 Digital Voltmeters: Today these measure voltage while drawing only a negligible amt. of current.

31. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 31 Next  Section 17.2 Standard Reduction Potential

32. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 32 GALVANIC CELL 1) Always an Oxidation-Reduction Reaction 2) Consists of Two Half-Reactions 3) Total Cell Potential = Sum of Cell Potentials for each Half-reaction

33. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 33 Consider a Galvanic Cell for the following Reaction: 2 H + (aq) + Zn(s)  Zn 2+ (aq) + H 2 (g) Half Reactions: (1) Oxidation (Anode): Zn  Zn 2+ + 2e - (2) Reduction (Cathode): 2H + + 2e -  H 2 (g)

34. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 34 Observed Potential = 0.76 V A Zn/H + Galvanic Cell All cell compartments are in their standard states.

35. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 35 Cell Potentials can only be measured for the total cell. There is no way to measure the potentials of the individual electrode processes.

36. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 36 Scientific community has accepted the assignment for the following reaction 2H + + 2e -  H 2 as having a Cell Potential exactly equal to ZERO. Other Half-Cell Potentials are calculated based on this assignment.

37. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 37 Standard Hydrogen Electrode (SHE) Platinum electrode (chemically inert conductor) in contact with 1 M H + ions and bathed in H 2 gas at 1 atm. 2H + + 2e -  H 2 This process is arbitrarily assigned a value of exactly Zero Volts.

38. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 38 Standard Hydrogen Electrode 

39. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 39 E  cell = Sum of the 2 half potentials E  cell = E  (H +  H 2 ) + E  (Zn  Zn 2+ ) E  cell = 0.76 V (measured) Since E  (H +  H 2 ) = 0, Then E  (Zn  Zn 2+ ) = 0.76 V

40. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 40 Consider Zn (s) + Cu 2+ (aq)  Zn 2+ (aq) + Cu (s) Anode: Zn (s)  Zn 2+ + 2e - Cathode: Cu 2+ + 2e -  Cu(s) Electrode compartments in standard state.

41. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 41 Figure 17.6 A Zn/Cu Galvanic Cell Measured Cell Potential = 1.10 V

42. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 42 E  cell = E  (Zn  Zn 2+ ) + E  (Cu 2+  Cu) 1.10 V = 0.76 V + E  (Cu 2+  Cu) Therefore, E  (Cu 2+  Cu) = 0.34 V

43. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 43 Standard Reduction Potential at 25°C (298 K) See Table 17.1 Accepted Convention to give the potential of half reactions as Reduction Processes.

44. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 44 Standard Reduction Potential at 25°C (298 K) Recall Zn (s)  Zn 2+ + 2e - E ° = 0.76V When a half-reaction is reversed, the sign of the potential must be reversed. Zn 2+ + 2 e -  Zn (s)E ° = - 0.76V (See Table 17.1)

45. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 45 Standard Reduction Potentials The E  values corresponding to reduction half- reactions with all solutes at 1M and all gases at 1 atm. Cu 2+ + 2e   Cu E  = 0.34 V vs. SHE SO 4 2  + 4H + + 2e   H 2 SO 3 + H 2 O – E  = 0.20 V vs. SHE

46. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 46 Calculating the E ° for a Galvanic Cell Obtain the Standard Reduction Potentials from the Table or Appendix. Reverse the sign of the potential for the substance being oxidized. To achieve balanced equation, the # of e - lost must equal the # of e - gained; however, the value of E ° is not changed when a half-rxn is multiplied by an integer. E  is an intensive property.

47. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 47 Galvanic Cell A Galvanic Cell will always run spontaneously in the direction that produces a positive cell potential. That is, E ° cell > 0

48. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 48 Read Line-Notation. Anode compartment written on the left. Electrode/anode listed on the far left. Vertical lines separate phases. Double vertical lines imply salt bridge or porous disk. Cathode components after double lines with substance constituting the electrode/cathode written at the far right.

49. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 49 Mg (s) | Mg 2+ (aq) || Al 3+ (aq) | Al (s) Anode Cathode

50. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 50 Complete Description of a Galvanic Cell The Balanced Cell Rxn. & the E ° cell. The direction of e - flow Designation of the anode & cathode. The nature of each electrode & the ions present in each compartment.

51. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 51 If all components in the reaction are dissolved ions or bubbled in gases, Then a non-reacting (chemically inert) conductor is used as an electrode, usually platinum.

52. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 52 Problem Describe completely a galvanic cell based on the following half-reactions. Fe 2+ + 2e -  FeE  = - 0.44 V MnO 4 - + 5e - + 8 H +  Mn 2+ + 4 H 2 O E  = 1.51 V Which way will be Spontaneous?

53. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 53 Which way will be Spontaneous? Fe  Fe 2+ + 2e - E  = + 0.44 V MnO 4 - + 5e - + 8 H +  Mn 2+ + 4 H 2 O E  = 1.51 V E  cell = 0.44 V + 1.51 V = 1.95 V

54. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 54 Balance Equation 5 x [Fe  Fe 2+ + 2e - ] 2 x [MnO 4 - + 5e - + 8 H +  Mn 2+ + 4 H 2 O] 5 Fe (s) + 2MnO 4 - (aq) + 16 H + (aq)  5 Fe 2+ (aq) + 2Mn 2+ (aq) + 8 H 2 O (l)

55. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 55 Line Notation Fe | Fe 2+ (aq) || MnO 4 - (aq), H + (aq), Mn 2+ (aq) | Pt (s)

56. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 56 Figure 17.7 A Schematic of a Galvanic Cell

57. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 57 Read Sample Exercise 17.2.

58. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 58 Figure 17.8 A Schematic of a Galvanic Cell