Chapter 18: Electrochemistry
2 types of cells:
1. Galvanic Cell: a spontaneous chemical reaction is used to produce electrical energy
2. Electrolytic Cell: a nonspontaneous reaction is made to occur by forcing an electric current
through the cell (there will be an energy source in the circuit i.e. a battery)
Galvanic Cells
Electrical enegy is measures in volts (V) by a voltmeter
Components:
2 electrodes immersed in solutions containing metal cations
Anode: oxidation (loss of e- occurs)
Cathode: reduction (gain of e- occurs)
2 compartments connected by a salt bridge or porous disk
(usually glass tube or filter paper) see figure at right
 Oxidizing agent gets reduced at the cathode
 Reducing agent gets oxidized at the anode
Line notation
reactant/product ‖ reactant/product (salt bridge indicated by ‖)
anode reaction cathode reaction
Example 1: Give the correct line notation for the galvanic cell pictured below.
Standard Reduction Potentials
 Cell potential = Ecell = potential of cell to do work on surroundings by driving an electric current
through a wire
 Potential = 1 volt is produced when 1 Joule of energy moves 1 coulomb of electric charge across
a potential.
 Magnitude of cell potential is a measure of the driving force behind an electrochemical reaction.
 Also referred to as electromotive force or emf
 Tables of reduction potentials give standard voltages for reduction half reactions measured at
standard conditions of 1 atm, 1 M solutions and 25 ⁰C (will be provided on AP exam, see last
page of notes)
Calculating Cell Potential of Galvanic Cell
 Break down reaction into half reactions (chapter 4)
 Use Table of Standard Reduction Potentials to calculate the cell potential of the overall reaction
Example 2: Consider a galvanic cell based on the reaction: Al + Ni2+ → Al3+ + Ni
Give the balanced cell reaction and calculate the cell potential, E⁰cell, for the reaction.
(E⁰cell = E⁰ox + E⁰red)
Spontaneous Reactions
Gibbs Free Energy, ΔG⁰, as mentioned in chapter 17, can be calulcated from the cell potential E⁰cell.
ΔG⁰ = -nFE⁰cell
n = moles of electrons transferred in redox reaction
F = Faraday’s constant = 96 485 C/mol eA spontaneous reaction has a negative value for ΔG⁰ or a positive value for E⁰cell.
Example 3: Will 1 M HCl dissolve silver metal and form Ag+ solution?
Example 4: Bromine, Br2, can oxidize iodide, I-, to iodine, I2. However, Br2, cannot oxidize chlorine, Cl2,
to chlorine, Cl2. Explain why the first reaction occurs and the second reaction does not.
Cell Dependence on Concentration
Eample 5: From Example 2, the galvanic cell based on the reaction: 2 Al + 3 Ni2+ → 2 Al3+ + 3 Ni
has a E⁰cell = 1.43 V under standard condition (all solutions = 1 molar)
According to LeChâtelier’s Principle, increasing the concentration of Ni2+ will shift the reaction to the
________.
 Increases the ____________ ______ on the e Increases the _____ _______________
Nernst Equation: relationship between cell potential and concentrations at 25 ⁰C
Ecell = E⁰cell - (0.0591/n) log Q
Ecell = non-standard conditions
n = moles of electrons transferred
Q = mass action quotient at specified conditions
Example 6: Calculate the cell potential for the above reaction in which [Al3+] = 2.00 M and
[Ni2+] = 0.750 M.
Determining the Strength of Oxidizing and Reducing Agents
** oxidation and reduction always occur together
 The larger (more positive) E⁰red indicates the stronger oxidizing agent
 The larger (more positive) E⁰ox indicates the stronger reducing agent
Example 7: Classify each of the following as an oxidizing agent, reducing agent or both.
Within each list, arrange in order of increasing strength as oxidizing and reducing agents.
Br2, Mg, Fe2+, I2, Cl-, Cu2+
Oxidizing Agents:
Reducing Agents:
Electrolytic cells
The reverse reaction of the previous example is nonspontaneous and can be made to occur by the
addition of an external power source.
2 Al3+ + 3 Ni → 2 Al + 3 Ni2+
Set up similar to a galvanic cell with the replacement of an external power source (battery) for the
voltmeter. (see figure below)
For example, if an object is to be plated with copper, make it the cathode and immerse it into a copper
(II) sulfate solution.
Write the reaction that will occur at the cathode:
The anode can also be made of copper where _________________ of copper occurs.
Reactions that Occur in an Electrolytic Cell
In an aqueous solution, the oxidation or reduction of water must be considered
Example 8: A solution of copper (II) sulfate is electrolyzed. Calculate the cell potential of the reaction,
E⁰cell.
Stoichiometry of Electrolytic Processes
How much chemical change occurs with the flow of a given current for a specified time.
 How much metal was plated (formed)
 How long an electroplating process will take
 How much current is required to produce a specified amount of metal over a period of time
Example 9: A current of 10.0 A is passed through a solution containing M2+ for 30.0 min. It produces
5.94 g of metal, M. Determine the identity of the metal, M.
Comparison of Galvanic and Electrolytic Cells
Both
 Reduction always occurs and cathode
 Oxidation always occurs at anode
 e- travel in the wire
 positive ions move toward the cathode (excess of negative ions at the cathode from the
reduction of the positive ions in solution
Galvanic
 e- travel from anode to cathode
 oxidation at anode produces positive ions, so negative ions in the salt bridge must move to the
anode to maintain electrical neutrality
Electrolytic
 e- travel from battery to the anode
 doesn’t make sense to say e- travel from anode to cathode