1. Zn (s) + Cu2+ (aq)  Zn2+ (aq) + Cu (s)
Galvanic Cells
When Zn metal is immersed in an aqueous solution of CuSO4, a spontaneous redox reaction occurs:
This same redox reaction can also occur when reactants are indirectly in contact with each other in a galvanic (voltaic) cell.
Zn (s) + Cu2+ (aq)  Zn2+ (aq) + Cu (s)

2. 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.

3. 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

4. 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

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

6. 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

7. Zn (s) + Ni2+ (aq) Zn2+ (aq) + Ni (s)
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?

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

9. Zn (s) + 2 H+ (aq)  Zn2+ (aq) + H2 (g)
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.

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

11. Galvanic Cells Anode higher potential energy Cathode
lower potential energy

12. 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

13. 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

14. 2H+ (aq, 1M) + 2 e- H2 (g, 1 atm) Eored = 0 V
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

15. Galvanic Cells
As Eored becomes increasingly positive, the driving force for reduction increases.
Reduction becomes more spontaneous
Reaction occurs at cathode
F2 (g) + 2e F- (aq) Eored = V
Ag+ (aq) + e Ag (s) Eored = 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?

16. Li+ (aq) + e- Li (s) Eored = -3.05
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-

17. 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 = V
Na+ (aq) + e Na (s) Eored = V

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

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

20. 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 = V
Fe2+ (aq) + 2 e-  Fe (s) Eored = V

21. Galvanic Cells

22. 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

23. 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)