Galvanic Cells

1. Galvanic Cells • What will happen if a piece of Zn metal is immersed in a CuSO4 solution? • A spontaneous redox reaction occurs: Zn (s) + Cu2+(aq) Zn2+(aq) + Cu (s) • Spontaneous reaction: a reaction that doesn’t need to be driven by an outside source of energy

2. Galvanic Cells • The previous reaction occurred when the Zn metal was in directcontact with the Cu2+ ions. • Redox reactions can also occur when reactants are indirectly in contact with each other in a galvanic (voltaic) cell.

3. Galvanic Cells • Galvanic (voltaic) cell: • A device in which a spontaneous redox reaction occurs as electrons are transferred from the reductant to the oxidant through an external circuit • used to perform electrical work using the energy released during a spontaneous redox reaction.

4. Galvanic Cells • In a galvanic cell, the two half reactions occur in separate compartments called half-cells. • 1 half-cell contains the oxidation half reaction • 1 half-cell contains the reduction half reaction • Each half cell contains: • electrode • electrolyte solution

5. Galvanic Cells • The two half cells are connected by • external circuit (wire) between the electrodes • salt bridge between the electrolyte solutions • ionic solution that will not react with other components in the galvanic cell • NaNO3 • completes the electrical circuit

6. Galvanic Cells Zn (s) + Cu2+(aq) Zn2+(aq) + Cu (s) electrode electrode Oxidation half cell Reductionhalf cell

7. Galvanic Cells • Two types of electrodes: • anode: • the electrode at which oxidation occurs • located in the oxidation half-cell • the “negative” electrode • electrons are released here • cathode: • the electrode at which reduction occurs • located in the reduction half-cell • the “positive” electrode • electrons move toward (are gained at) the cathode

8. Galvanic Cells Consider the following reaction: Zn(s) + Ni2+ (aq) Zn2+ (aq) + Ni (s) • Which metal will be the anode? • Which metal will be the cathode?

9. Galvanic Cells • In some galvanic cells, one (or both) of the half reactions does not involve a metal: Cr2O72- (aq) + 14 H+ (aq) + 6 I- (aq) 2 Cr3+ (aq) + 3 I2 (s) + 7 H2O (l) • In these cases, an unreactivemetal conductor is used as the electrode • platinum foil

10. Galvanic Cells Zn (s) + 2 H+ (aq) Zn2+ (aq) + H2 (g) • Oxidation half-reaction: • Zn (s) Zn2+ (aq) + 2 e- • Reduction half-reaction: • 2 H+ (aq) + 2 e- H2 (g) • In this case a standard hydrogen electrode is used as the cathode.

11. Cell EMF • The redox reactions occurring in a galvanic cell are spontaneous. • Why do electrons flow spontaneously from one electrode to the other? • Electrons flow spontaneously because there is a difference in potential energy between the anode and the cathode.

12. Galvanic Cells • Anode • higher potential energy • Cathode • lower potential energy

13. Galvanic Cells • The difference in electrical potential between the anode and the cathode is called the cell potential or cell voltage (Ecell) • measured in volts • Standard cell potential (Eocell): • the cell potential measured under standard conditions • 25oC • 1M concentrations of reactants and products in solution • or 1 atm pressure for gases

14. Galvanic Cells • Eocell depends on the half-cells or half-reactions present • Standard potentials have been assigned to each individual half-cell • By convention, the standard reduction potential (Eored) for each half cell is used and tabulated

15. Galvanic Cells • Standard reduction potential: • potential of a reduction half-reaction under standard conditions • measured relative to the reduction of H+ to H2 under standard conditions: 2H+ (aq, 1M) + 2 e- H2 (g, 1 atm)Eored = 0 V

16. Galvanic Cells • As Eored becomes increasingly positive, the driving force for reduction increases. • Reduction becomes more spontaneous • Reaction occurs at cathode F2(g) + 2e- 2 F-(aq) Eored = +2.87 V Ag+(aq) + e- Ag (s) Eored = + 0.80 V Which reaction is more spontaneous as written? Which reaction will tend to occur at the cathode if the two reactions were combined in a galvanic cell?

17. Galvanic Cells • As Eored becomes increasingly negative, the driving force for oxidation increases. Li+(aq) + e- Li (s) Eored = -3.05 • The negative reduction potential indicates that the reverse (oxidation) half-reaction is spontaneous. • The reaction that occurs at the anode is: Li (s) Li+(aq) + e-

18. Galvanic Cells Example: Given the following standard reduction potentials, which of the metals will be most easily oxidized? Ag+(aq) + e- Ag (s) Eored = 0.80 V Zn2+(aq) + 2 e- Zn (s) Eored = -0.76 V Na+ (aq) + e- Na (s) Eored = -2.71 V

19. Galvanic Cells • Standard cell potential Eocell = Eored (cathode) - Eored (anode) reduction oxidation

20. Galvanic Cells Example: What is the Eocell for the following reaction? Zn (s) + Cu2+(aq) Zn2+(aq) + Cu (s)

21. Galvanic Cells

22. Galvanic Cells Example: Given the following reduction half-reactions, identify the metal at the anode, the balanced reaction for the galvanic cell, and the Eocell. Al3+ (aq) + 3 e- Al (s) Eored = -1.66 V Fe2+ (aq) + 2 e- Fe (s) Eored = -0.440 V

23. Galvanic Cells

24. Galvanic Cells • Oxidizing Agent (oxidant): • the substance that causes another to be oxidized • the substance that is reduced • the substance that gains electrons • The strongest oxidizing agent is the substance that has the greatest tendency to be reduced. • The most positive Eored

25. Galvanic Cells Example: Use the reduction potentials given in Appendix E to determine which of the following is the stronger oxidizing agent: Br2 (l) or I2 (s)