Galvanic Cells
 When Zn metal is immersed in an aqueous
solution of CuSO4, a spontaneous redox reaction
occurs:
Zn (s) + Cu2+ (aq)  Zn2+ (aq) + Cu (s)
 This same redox reaction can also occur when
reactants are indirectly in contact with each
other in a galvanic (voltaic) cell.
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
Reduction
half 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
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
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
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,
driving force for reduction increases.
 Reduction becomes more spontaneous
 Reaction occurs at cathode
the
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?
Galvanic Cells
 As
Eored becomes increasingly negative,
driving force for oxidation increases.
Li+ (aq) + e-
Li (s)
the
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
Galvanic Cells
 Standard cell potential
Eocell = Eored (cathode) - Eored (anode)
reduction
oxidation
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 halfreactions, 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)