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)

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.

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

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

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

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

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?

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 unreactive metal conductor is used as the electrode platinum foil

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.

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.

Galvanic Cells Anode higher potential energy Cathode lower potential energy

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

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

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

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?

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-

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

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

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)

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

Galvanic Cells

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

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)