Please Sit with Your Group

1. Please Sit with Your Group • Please be sure each member of your team has a copy of • Electrochemical Cells Lecture Notes • Electrochemical Cells Problem Set • Today’s reporter is the person who’s box number has the most odd digits. • Next reading assignment: • Zumdahl Chapter 11.5

2. Electrochemical Cells Edward A. Mottel Integrated, First-Year Curriculum in Science, Engineering and Mathematics

3. Electrochemical Cells • Voltaic and electrolytic cells • Electrochemical cell structure • Types of electrochemical cells

4. Electrochemical Cell • A physical arrangement involving • an oxidation reaction • a reduction reaction • The potential of the cell can be • spontaneous • Voltaic cell (also called a Galvanic cell) • non-spontaneous • Electrolytic cell

5. Voltaic Cell • An electrochemical cell which • spontaneously generates a positive electrical potential • can be used for useful work • has Ecell > 0 as constructed • Example • A discharging battery • rechargeable or non-rechargeable

6. Electrolytic Cell • An electrochemical cell which • requires an external energy source to force the cell in a non-spontaneous direction. • has Ecell < 0 as constructed. • Examples • A battery being recharged. • A piece of metal being electroplated.

7. Electrochemical Cell Structure Cathode Electrode Anode Electrode Salt Bridge Cathode Solution Anode Solution

8. Electrochemical Cell Structure Cathode Electrode Anode Electrode Salt Bridge Cathode Solution Anode Solution

9. Electrochemical Cell Structure Cathode Electrode Anode Electrode Salt Bridge Cathode Solution Anode Solution

10. 0.000 V Electrochemical Cell Structure 2.002 V Cathode Electrode Anode Electrode Salt Bridge Cathode Solution Anode Solution

11. Electrochemical Cell Structure • Half-cell reactions • Electrodes • Electron flow • Ion flow • Shorthand notation

12. Al(s) Al3+(aq) + 3 e– Cu2+(aq) + 2 e– Cu(s) Half-Cell Reactions • Each electrochemical cell involves both an oxidation reaction and a reduction reaction. • The oxidation cell and the reduction cell are referred to as half-cells.

13. Al(s) Al3+(aq) + 3 e– Anode Reaction 2.002 V Cathode Electrode Anode Electrode Salt Bridge Cathode Solution Anode Solution

14. Al(s) Al3+(aq) + 3 e– AnodeThe electrode at which oxidation occurs 2.002 V Cathode Electrode Anode Electrode Salt Bridge Cathode Solution Anode Solution

15. Al(s) Al3+(aq) + 3 e– Anode of a Voltaic Cell is Negative 2.002 V Cathode Electrode Anode Electrode e– Salt Bridge - Al Cathode Solution Anode Solution because electrons are released

16. Al(s) Al3+(aq) + 3 e– e– Al 2.002 V Cathode Electrode Anode Electrode Salt Bridge - Cathode Solution Anode Solution

17. Cu2+(aq) + 2 e– Cu(s) Al(s) Al3+(aq) + 3 e– e– Al Cathode Reaction 2.002 V Cathode Electrode Anode Electrode Salt Bridge - Cathode Solution Anode Solution

18. Al(s) Al3+(aq) + 3 e– Cu2+(aq)+ 2 e– Cu(s) e– Al CathodeThe electrode at which reduction occurs 2.002 V Cathode Electrode Anode Electrode Salt Bridge - Cathode Solution Anode Solution

19. Al(s) Al3+(aq) + 3 e– Cu2+(aq)+ 2 e– Cu(s) e– e– Al Cu2+ Cathode of a Voltaic Cell is Positive 2.002 V Cathode Electrode Anode Electrode Salt Bridge - + Cathode Solution Anode Solution because electrons are attracted and consumed

20. Al(s) Al3+(aq) + 3 e– Cu2+(aq)+ 2 e– Cu(s) e– e– Al Cu2+ 2.002 V Cathode Electrode Anode Electrode Salt Bridge - + Cathode Solution Anode Solution

21. Al(s) Al3+(aq) + 3 e– Cu2+(aq)+ 2 e– Cu(s) e– e– Al Cu2+ Electrons are transferredthrough a wire from anode to cathode 2.002 V Cathode Electrode Anode Electrode Salt Bridge - + Cathode Solution Anode Solution

22. Al(s) Al3+(aq) + 3 e– Cu2+(aq)+ 2 e– Cu(s) e– e– Al Cu2+ Electron Current Flowmay be used to perform useful work Cathode Electrode Anode Electrode Salt Bridge - + Cathode Solution Anode Solution Electrical connection is made at the electrodes, the site at which oxidation and reduction occurs.

23. Keeping It Straight Anode Cathode Oxidation Reduction Electrons are attracted and consumed Electrons are released In a voltaic cell it is the negative electrode In a voltaic cell it is the positive electrode

24. Electrons are transferred through a wire from the anode to the cathode. Electron Flow Ion Flow • Anions are attracted to the anode and cations migrate away from anode. Salt Bridge • The salt bridge contains an ionic compound such as KNO3 or NaCl dissolved in a gel such as agar-agar.

25. Draw a Diagram indicate what is happening to all the charged species in the anode cell. List charged species Show their location and their motion

26. e– NO3– K+ + + Al Al NO3– + Al3+ + NO3– + Anode Cell Show the motion of all the charged species

27. Ion Flow • Cations are attracted to the cathode and anions migrate away from cathode. Draw a diagram indicating what is happening to all the charged species in the cathode cell.

28. e– NO3– K+ – – – NO3– – K+ Cu – Cu2+ – Cathode Cell Identify the main species Show the motion of all the charged species

29. Salt Bridge • A salt bridge may be used to physically separate ions in one half-cell from ions in the other half-cell. Draw a diagram indicating what is happening to all the charged species in the salt bridge.

30. K+ NO3– NO3– K+ K+ K+ NO3– Salt Bridge NO3– Al3+

31. Shorthand Line Notation Al(s) | Al3+ (1.00 M) | | Cu2+ (1.00 M) | Cu(s) anode | anode solution | | cathode solution | cathode H2(g, 1 atm), Pt(s) | H+ (1 M) | | Cl– (1 M) | Cl2(g, 1 atm), C(gr) Why is a graphite or a platinum electrode needed?

32. Types of Electrochemical Cells • Concentration Cell • Standard Redox Cell • Non-standard (Combination) Redox Cell

33. Concentration Cell • The oxidation and reduction reactions are identically reverse of each other. • The observed cell potential is due solely to differences in concentrations of the solutions involved. • Low potentials generated (mV)

34. Zn(s) Zn2+(0.23 M) + 2 e– Zn2+ (1.00 M) + 2 e– Zn(s) Concentration Cell • Example: Zn(s) | Zn2+ (0.23 M) | | Zn2+ (1.00 M) | Zn(s) Write the oxidation and reduction half-cell reactions taking place in this cell.

35. 2 + 2 + [ Zn ] [ Zn ( 0 . 23 M )] anode = = 2 + 2 + [ Zn ] [ Zn ( 1 . 00 M )] cathode Concentration Cell • Example: Zn(s) | Zn2+ (0.23 M) | | Zn2+ (1.00 M) | Zn(s) Write the Q term for this cell.

36. Zn(s) Zn2+(0.23 M) + 2 e– Zn2+ (1.00 M) + 2 e– Zn(s) Concentration Cell • Example: Zn(s) | Zn2+ (0.23 M) | | Zn2+ (1.00 M) | Zn(s) Determine the standard cell potential for this cell.

37. Concentration Cell • E°cell = 0.00 V • Low potentials generated (mV)

38. Standard Redox Cell • The oxidation and reduction reactions are different. • Concentrations of solutions are 1 M and reactant gas pressures are 1 atm. • The observed cell potential is due to the differences in the activity of the reactants.

39. Ni(s) Ni2+(1.00 M) + 2 e– Ag+ (1.00 M) +e– Ag(s) Standard Redox Cell • Example: Ni(s) | Ni2+ (1.00 M) | | Ag+ (1.00 M) | Ag(s) Write the oxidation and reduction half-cell reactions taking place in this cell. Write the Q term for this cell.

40. 2 + [ Ni ( 1 . 00 M )] Q = + 2 [ Ag ( 1 . 00 M )] Standard Redox Cell • Example: • Ni(s) | Ni2+ (1.00 M) | | Ag+ (1.00 M) | Ag(s) = 1 Why is this called a standard redox cell?

41. Ni(s) Ni2+(1.00 M) + 2 e– Ag+ (1.00 M) +e– Ag(s) Standard Redox Cell • Example: Ni(s) | Ni2+ (1.00 M) | | Ag+ (1.00 M) | Ag(s) Determine the standard cell potential for this cell.

42. Standard Redox Cell • E°cell¹ 0.00 V • Potentials (voltage) generated can be quite high

43. Non-standard (Combination) Redox Cell • The oxidation and reduction reactions are different. • The solution concentrations are not 1 M. • Gas pressures are not 1 atm.

44. Mn(s) Mn2+(1.00 M) + 2 e– Pb2+ (0.23 M) + 2 e– Pb(s) Non-standard (Combination) Redox Cell • Example: Mn(s) | Mn2+ (1.00 M) | | Pb2+ (0.23 M) | Pb(s) Write the oxidation and reduction half-cell reactions taking place in this cell. Write the Q term for this cell.

45. 2 + [ Mn ( 1 . 00 M )] Q = 2 + [ Pb ( 0 . 23 M )] Non-standard (Combination) Redox Cell • Example: Mn(s) | Mn2+ (1.00 M) | | Pb2+ (0.23 M) | Pb(s) = 4.3

46. Mn(s) Mn2+(1.00 M) + 2 e– Pb2+ (0.23 M) + 2 e– Pb(s) Non-standard (Combination) Redox Cell • Example: Mn(s) | Mn2+ (1.00 M) | | Pb2+ (0.23 M) | Pb(s) Determine the standard cell potential for this cell. Why is this called a non-standard redox cell?

47. Non-standard (Combination) Redox Cell • The majority of the observed cell potential is due to the differences in the activity of the reactants, modified slightly by non-standard conditions. • E°cell¹ 0.00 V • Potentials generated can be quite high (V)

48. Electrode Materials • Inert electrodes can or must be used in some instances. • The reactant or product is a gas or liquid. • The reactant and product of a half-cell are soluble. • The product is being plated out onto an inert electrode.

49. Inert Electrodes • Examples: Pt(s) | Cr2+ (1.00 M), Cr3+ (1.00 M) | | Cu2+ (1.00 M) | Au(s) Co(s) | Co2+ (0.789 M) | | Hg2+ (0.50 M) | Hg(l),Pt(s) H2(g, 30 atm), C (gr) | KOH(0.789 M) | | KOH(0.789 M) | O2(g, 20 atm), C (gr)

50. Pt(s) | Cr2+(1.0 M), Cr3+(1.0 M) | | Cu2+ (1.0 M) | Au(s) • Draw a beaker diagram for this cell. • Identify what is being oxidized and what is being reduced. • Indicate the flow of all cations, anions and electrons in your diagram. • What is the standard cell potential? • What is the Q term?