1. Please Pick Up Electrochemical Cells Problem Set

2. Electrochemical Cells Edward A. Mottel Department of Chemistry Rose-Hulman Institute of Technology

3. 6/12/2015 Electrochemical Cells  Reading assignment: Chang: Chapter 19.1-19.2  A physical arrangement designed for electron flow involving an oxidation reaction a reduction reaction

4. 6/12/2015 Voltaic Cell also called a Galvanic cell  An electrochemical cell which spontaneously generates a positive electrical potential can be used for useful work has E cell > 0 as constructed  Example A discharging battery ·rechargeable or non-rechargeable Corrosion of a piece of iron

5. 6/12/2015 Electrolytic Cell  An electrochemical cell which requires an external energy source to force the cell in a non-spontaneous direction. has E cell < 0 as constructed.  Examples A battery being recharged. A piece of metal being electroplated.

6. 6/12/2015 Electrochemical Cell Structure Cathode Solution Anode Solution Anode Electrode Cathode Electrode Salt Bridge

7. 6/12/2015 Electrochemical Cell Structure Cathode Solution Anode Solution Anode Electrode Cathode Electrode Salt Bridge

8. 6/12/2015 Electrochemical Cell Structure Cathode Solution Anode Solution Anode Electrode Cathode Electrode Salt Bridge

9. 6/12/2015 Electrochemical Cell Structure Cathode Solution Anode Solution Anode Electrode Cathode Electrode Salt Bridge 0.000 V 2.002 V

10. 6/12/2015 Electrochemical Cell Structure  Half-cell reactions  Electrodes  Electron flow  Ion flow  Shorthand notation

11. 6/12/2015 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. Al(s)Al 3+ (aq) + 3 e – Cu 2+ (aq) + 2 e – Cu(s)

12. 6/12/2015 Anode Reaction Cathode Solution Anode Solution Anode Electrode Cathode Electrode Salt Bridge 2.002 V Al(s)Al 3+ (aq) + 3 e –

13. 6/12/2015 Anode The electrode at which oxidation occurs Cathode Solution Anode Solution Anode Electrode Cathode Electrode Salt Bridge 2.002 V Al(s)Al 3+ (aq) + 3 e –

14. 6/12/2015 Anode of a Voltaic Cell is Negative Cathode Solution Anode Solution Anode Electrode Cathode Electrode Salt Bridge 2.002 V Al(s)Al 3+ (aq) + 3 e – Al e–e– - because electrons are released

15. 6/12/2015 Cathode Solution Anode Solution Anode Electrode Cathode Electrode Salt Bridge 2.002 V Al(s)Al 3+ (aq) + 3 e – Al e–e– -

16. 6/12/2015 Cathode Reaction Cathode Solution Anode Solution Anode Electrode Cathode Electrode Salt Bridge 2.002 V Al(s)Al 3+ (aq) + 3 e – Cu 2+ (aq) + 2 e – Cu(s) Al e–e– -

17. 6/12/2015 Cathode The electrode at which reduction occurs Cathode Solution Anode Solution Anode Electrode Cathode Electrode Salt Bridge 2.002 V Al(s)Al 3+ (aq) + 3 e – Cu 2+ (aq)+ 2 e – Cu(s) Al e–e– -

18. 6/12/2015 Cathode of a Voltaic Cell is Positive Cathode Solution Anode Solution Anode Electrode Cathode Electrode Salt Bridge 2.002 V Al(s)Al 3+ (aq) + 3 e – Al e–e– Cu 2+ e–e– because electrons are attracted and consumed + - Cu 2+ (aq)+ 2 e – Cu(s)

19. 6/12/2015 Cathode Solution Anode Solution Anode Electrode Cathode Electrode Salt Bridge 2.002 V Al(s)Al 3+ (aq) + 3 e – Al e–e– Cu 2+ e–e– - + Cu 2+ (aq)+ 2 e – Cu(s)

20. 6/12/2015 Electrons are transferred through a wire from anode to cathode Cathode Solution Anode Solution Anode Electrode Cathode Electrode Salt Bridge 2.002 V Al(s)Al 3+ (aq) + 3 e – Al e–e– Cu 2+ e–e– - + Cu 2+ (aq)+ 2 e – Cu(s)

21. 6/12/2015 Electron Current Flow may be used to perform useful work Cathode Solution Anode Solution Anode Electrode Cathode Electrode Salt Bridge Al(s)Al 3+ (aq) + 3 e – Al e–e– Cu 2+ e–e– - + Cu 2+ (aq)+ 2 e – Cu(s) Electrical connection is made at the electrodes, the site at which oxidation and reduction occurs.

22. 6/12/2015 Keeping It Straight Electrons are released In a voltaic cell it is the negative electrode Electrons are attracted and consumed In a voltaic cell it is the positive electrode A node O xidation C athode R eduction

23. 6/12/2015  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 KNO 3 or NaCl dissolved in a gel such as agar-agar.

24. 6/12/2015 indicate what is happening to all the charged species in the anode cell. List charged species Show their location and their motion Draw a Diagram

25. 6/12/2015 Anode Cell Al Al 3+ e–e– NO 3 – K+K+ Al Show the motion of all the charged species + + + + + NO 3 –

26. 6/12/2015 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.

27. 6/12/2015 Cathode Cell Identify the main species

28. 6/12/2015 Cathode Cell Show the motion of all the charged species e–e– – – – – – – Cu 2+ Cu NO 3 – K+K+ K+K+ Identify the main species

29. 6/12/2015 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. 6/12/2015 Salt Bridge

31. 6/12/2015 Salt Bridge NO 3 – K+K+ K+K+ K+K+ Al 3+ K+K+ NO 3 –

32. 6/12/2015 Shorthand Line Notation Al(s) | Al 3+ (1.00 M) | | Cu 2+ (1.00 M) | Cu(s) Why is a graphite or a platinum electrode needed? anode | anode solution | | cathode solution | cathode H 2 (g, 1 atm), Pt(s) | H + (1 M) | | Cl – (1 M) | Cl 2 (g, 1 atm), C(gr)

33. 6/12/2015 Types of Electrochemical Cells  Concentration Cell  Standard Redox Cell  Non-standard (Combination) Redox Cell

34. 6/12/2015 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)

35. 6/12/2015 Concentration Cell  Example: Zn(s) | Zn 2+ (0.23 M) | | Zn 2+ (1.00 M) | Zn(s) Zn(s)Zn 2+ (0.23 M) + 2 e – Zn 2+ (1.00 M) + 2 e – Zn(s) Write the oxidation and reduction half-cell reactions taking place in this cell.

36. 6/12/2015 Concentration Cell Zn anode cathode    [] [] 2 2 ZnM M    [(.)] [(. 2 2 023 100  Example: Zn(s) | Zn 2+ (0.23 M) | | Zn 2+ (1.00 M) | Zn(s) Write the Q term for this cell.

37. 6/12/2015 Concentration Cell  Example: Zn(s) | Zn 2+ (0.23 M) | | Zn 2+ (1.00 M) | Zn(s) Zn(s)Zn 2+ (0.23 M) + 2 e – Zn 2+ (1.00 M) + 2 e – Zn(s) Determine the standard cell potential for this cell.

38. 6/12/2015 Concentration Cell  E ° cell = 0.00 V  Low potentials generated (mV)

39. 6/12/2015 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.

40. 6/12/2015 Standard Redox Cell  Example: Ni(s) | Ni 2+ (1.00 M) | | Ag + (1.00 M) | Ag(s) Ni(s)Ni 2+ (1.00 M) + 2 e – Write the oxidation and reduction half-cell reactions taking place in this cell. Ag + (1.00 M) + e – Ag(s) Write the Q term for this cell.

41. 6/12/2015 Standard Redox Cell  Example: Ni(s) | Ni 2+ (1.00 M) | | Ag + (1.00 M) | Ag(s) Q Ni Ag    [(.M)] [(. 2 M 2 100 1  Why is this called a standard redox cell?

42. 6/12/2015 Standard Redox Cell  Example: Ni(s) | Ni 2+ (1.00 M) | | Ag + (1.00 M) | Ag(s) Ni(s)Ni 2+ (1.00 M) + 2 e – Determine the standard cell potential for this cell. Ag + (1.00 M) + e – Ag(s)

43. 6/12/2015 Standard Redox Cell  E ° cell  0.00 V  Potentials (voltage) generated can be quite high

44. 6/12/2015  The oxidation and reduction reactions are different.  The solution concentrations are not 1 M.  Gas pressures are not 1 atm. Non-standard (Combination) Redox Cell

45. 6/12/2015  Example: Mn(s) | Mn 2+ (1.00 M) | | Pb 2+ (0.23 M) | Pb(s) Write the oxidation and reduction half-cell reactions taking place in this cell. Mn(s) Mn 2+ (1.00 M) + 2 e – Pb 2+ (0.23 M) + 2 e – Pb(s) Write the Q term for this cell. Non-standard (Combination) Redox Cell

46. 6/12/2015 Non-standard (Combination) Redox Cell  Example: Mn(s) | Mn 2+ (1.00 M) | | Pb 2+ (0.23 M) | Pb(s) Q Mn Pb    [(.M)] [(.M 2 2 100 023 

47. 6/12/2015  Example: Mn(s) | Mn 2+ (1.00 M) | | Pb 2+ (0.23 M) | Pb(s) Mn(s) Mn 2+ (1.00 M) + 2 e – Pb 2+ (0.23 M) + 2 e – Pb(s) Why is this called a non-standard redox cell? Determine the standard cell potential for this cell. Non-standard (Combination) Redox Cell

48. 6/12/2015  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 ° cell  0.00 V  Potentials generated can be quite high (V) Non-standard (Combination) Redox Cell

49. 6/12/2015 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.

50. 6/12/2015 Inert Electrodes Examples H 2 (g, 30 atm), C(gr) | KOH (0.789 M) | | KOH (0.789 M) | O 2 (g, 20 atm), C(gr) Pt(s) | Cr 2+ (1.00 M), Cr 3+ (1.00 M) | | Cu 2+ (1.00 M) | Au(s) Co(s) | Co 2+ (0.789 M) | | Hg 2+ (0.50 M) | Hg( l ), Pt(s)

51. 6/12/2015 Pt(s) | Cr 2+ (1.0 M), Cr 3+ (1.0 M) | | Cu 2+ (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?

52. 6/12/2015