CHAPTER 21 – PRINCIPLES OF REACTIVITY: ELECTRON TRANSFER REACTIONS
Chapter 21 Problem Set
Pages 996-1001
18, 24, 28, 32, 40, 46, 60
How do batteries and fuel cells work? What do they have in common with the reaction of zinc and
HCl?
21.1 OXIDATION-REDUCTION REACTIONS
Redox reactions occur by electron transfer.
Some important examples of redox reactions are the oxygen you breathe, photosynthesis.
A battery is an electrochemical cell that produces a current, or flow of electrons, at a constant voltage.
Electrolysis is the direct application of electricity to produce a chemical change.
Balancing Equations for Oxidation-Reduction Reactions
Consider the reaction of zinc + copper (II) sulfate Zn(s) + CuSO4 (aq)  ZnSO4 (aq) + Cu (s)
The equation for the net chemical reaction is the sum of the two half reactions
Zn(s)  Zn2+(aq) + 2eCu2+(aq) + 2e-  Cu(s)__________
Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu (s)
Cu2+ gains electrons
Is reduced
Is the oxidizing agent
Zn loses electrons
Is oxidized
Is the reducing agent
Notice each half reaction is balanced for mass.
Each half reaction has a balanced charge, known as a charge balance
Therefore the net equation is balanced for mass and charge.
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EXERCISE 21.1 Page 950 Balancing an Equation for an Oxidation-Reduction Reaction
Balance the equation Cr2+(aq) + I2(aq)  Cr3+(aq) + I-(aq)
1. Write balanced ½ reactions
2. Write balanced net ionic equation
3. Identify the oxidizing and reducing agents and the substance oxidized and the substance reduced.
EXERCISE 21.2 Page 954 Balancing Equations for Oxidation-Reduction Reactions in Acid
Solution
Cobalt metal reacts with nitric acid to give a cobalt (III) salt and NO2 gas. The unbalanced net ionic
equation is:
Co(s) + NO3-(aq)  Co3+ (aq) + NO2(g)
1. Balance the equation for the reaction in acid solution.
2. Identify the oxidizing and reducing agents and the substance oxidized and the substance reduced.
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EXERCISE 21.3 Page 957 Balancing Equations for Oxidation-Reduction Reactions in Basic
Solution
Batteries based on the reduction of sulfur appear very promising. One being studied currently involves
the reaction of sulfur with aluminum in base.
Al(s) + S(s)  Al(OH)3(s) + HS-(aq)
1. Balance the equation, showing each balanced half-reaction (in basic solution).
2. Identify the oxidizing and reducing agents, the substance oxidized, and the substance reduced.
Problem-Solving Tips And Ideas Page 957
21.1 Balancing Equations of Oxidation-Reduction Reactions
21.2 CHEMICAL CHANGE LEADING TO AN ELECTRIC CURRENT
Consider the Zn and CuSO4 reaction.
In order to use the reaction as the basis of a battery, zinc metal and Cu2+ ions must be placed in
separate containers.
Electrons can be made to travel an external wire as they go from zinc to copper.
On the way they can do work.
This arrangement only works with a salt bridge, whose function is to allow ions to pass freely from the
compartment where cations are being lost to the compartment where cations are being generated.
A NaSO4 soln could be used as a salt bridge to transfer the ions.
An oxidizing agent and a reducing agent arranged so they can react only if electrons flow through an
outside conductor is called an electrochemical cell, a voltaic cell, or a battery.
In all electrochemical cells oxidation occurs at the anode, and reduction occurs at the cathode.
See Figure 21.5 Page 958 and Figure 21.6 Page 959 for definitions.
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EXERCISE 21.4 Page 961 Electrochemical Cells
A voltaic cell has been assembled with the net reaction
Ni(s) + 2 Ag+(aq)  Ni2+(aq) + 2 Ag(s)
Give the ½ reactions for this electron transfer process, indicating whether each is an oxidation or
reduction and deciding which happens at the anode and which is at the cathode.
What is the direction of electron flow in an external wire connecting the two electrodes?
If a salt bridge connecting the cell compartment contains KNO3, what is the direction of flow of the
nitrate ions?
21.3 ELECTROCHEMICAL CELLS AND POTENTIALS
An electromotive force (emf) is the force that makes electrons move.
The quantity of electric work done is proportional to the number of electrons that go from higher to
lower potential energy and to the size of the potential energy difference.
Electrical work = charge x potential energy difference, charge is measured in coulombs (C)
C = quantity of electrical charge that passes a point in an electrical circuit when a current of 1 ampere
flows for one second.
volt 
1 joule
1 coulomb
or
Work ( joule )  1 volt x 1 coulomb
Cell potential depends upon identity of subs used and the conc of the solutes or pressures of gases at
electrodes. Because the potential varies if the conditions are changed, standard conditions have been
defined: standard soln are 1.0 M and standard gas press are 1 bar.
A standard electrode potential has symbol E°.
Cell potentials are positive values for product-favored reactions
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E° and G°
The standard cell potential is a measure of reaction spontaneity. The relation between the standard cell
potential E° and the standard free energy change G°rxn is:
EQUATION 21.1
G°rxn = -nFE°
Where n = # moles of electrons and F = Faraday constant, 9.6485309 x 104 J/V. mol
EXERCISE 21.5 Page 964 The Relation Between E° and G°rxn
The following reaction has an E° value of –0.76 V. Calculate G°rxn and decide whether the reaction
is product-favored or reactant-favored.
H2(g) + 2 H2O(l) + Zn2+(aq)  Zn s) + 2 H3O+(aq)
Calculating the Potential E° of an Electrochemical Cell
Redox reactions are the sum of 2 half reactions. If a half reaction is assigned a voltage then summing
the half-reaction voltages would then give the cell voltage. But the voltage of the cell measures the
potential energy difference for electrons in 2 different chemical environments. Therefore, we use a
standard half-cell reaction against which all others are measured. The standard half-cell and its
assigned potential is the standard hydrogen electrode,
2 H3O+(aq) + 2e-  H2 (g, 1 bar) + 2 H2O(l)
E° = 0.00 V
A potential of 0.00 V has been assigned to this half reaction. (This value has no physical meaning in
itself, just as the half-reaction alone has no meaning)
The measured voltage of an electrochemical cell is always positive.
*Consider 3 standard electrochemical cells and measure the potential for each cell. We will break each
cell into 2 standard half cells.
READ PGS 964-967 FOR EXPLANATION OF THIS SECTION
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EXERCISE 21.6 Page 968 Determining a Half-Reaction Potential
Given that the reduction of aqueous copper (II) with iron metal has an E°net value of +0.78 V, what is
E° for the half-cell Fe (s)  Fe2+ (aq, 1 M) + 2e-?
Fe(s) + Cu2+(aq, 1 M)  Fe2+(aq, 1 M) + Cu (s)
21.4 USING STANDARD CELL POTENTIALS
If 2 half-cells are combined and the sum of the potentials is negative, what does this tell us about the
electrochemical cell?
We have considered 4 elements that could act as reducing agents: Zn, Cu, Ni, and H2. What are their
reducing abilities?
What are the relative abilities of Zn2+, Ni2+, Cu2+, and H+ to act as oxidizing agents?
See bottom Page 968 and top Page 969
We have created our own Table of Standard Reduction Potentials (see Table 21.1).








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E° values written in form “oxidized form + electrons = reduced form. Oxidizing agents are on the
left, reducing agents on right. All potentials are for reduction reactions.
Signs are reversed if reactions are reversed
All half-reactions are reversible
The more positive the value of E°, the better the oxidizing ability of the ion or compound. In this
table, the strongest oxidizing agent is F2
In this table, the strongest reducing agent is Li (s)
Under standard conditions, any substance on the left of this table will spontaneously oxidize any
substance lower than it on the right.
The algebraic sign of the half-reaction potential is the sign of the electrode when it is attached to
the H2/H3O+ standard cell.
Electrochemical potentials depend on the nature of the reactants and products and their
concentrations, not on the quantities of material used. Changing coefficients does not change value
of E°.
EXERCISE 21.7 Page 971 Predicting E° and the Direction of a Redox Reaction
Is the following reaction product-favored under standard conditions? What is the value of E°net?
H2O2(aq) + 2H3O+ (aq) + Br-(aq)  Br2(l) + 4 H2O(l)
EXERCISE 21.8 Page 973 Constructing an Electrochemical Cell
Draw a diagram of an electrochemical cell using the half-cells Zn (s)/ Zn2+ (aq) and Al (s)/Al3+ (aq).
Decide first on the net reaction, and predict its E° value. Show the direction of electron flow in the
external wire and the directions of ion flow in the salt bridge. Tell which compartment is the anode
and which is the cathode.
21.5 ELECTROCHEMICAL CELLS AT NONSTANDARD CONDITIONS
The Nernst Equation
The standard potential, E° can be corrected by a factor that includes the temperature of the reaction, the
number of moles of electrons transferred between in reactants, and the concentrations of reactants and
products. This relationship is the Nernst Equation.
E = E° - (RT/nF) ln Q When T = 298 K we can rewrite a modified form that we find more useful.
EQUATION 21.2
E  E 
0.0257 V
ln Q
n
at 25 C
*Sometimes the Nernst Equation is written in this format:
E  E 
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0.0592 V
log Q
n
at 25 C
EXERCISE 21.9 Page 975 Using the Nernst Equation
Calculate Enet for the following reaction:
2 Ag+(aq, 0.80M) + Hg(l)  2 Ag(s) + Hg2+(0.0010 M, aq)
Is this reaction product-favored or reactant-favored under these conditions? How does this compare
with the reaction under standard conditions?
E° and the Equilibrium Constant
The cell voltage and even the reaction direction can change when the concentrations of reactants and
products change.
What does it mean when E° had a potential of zero?
E  0  E 
0.0257 V
ln K
n
The Nernst equation can be rewritten as:
n E
EQUATION 21.3
ln K  
0.0257 V
at 25 C
This is a very useful equation because it allows us to obtain the equilibrium constant for a reaction
from a calculation or measurement of E°net
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EXERCISE 21.10 Page 976 E° and Equilibrium Constants
In Exercise 21.9 you calculated Enet for
2 Ag+(aq, 0.80M) + Hg(l)  2 Ag(s) + Hg2+(0.0010 M, aq)
What is the equilibrium constant for this reaction?
21.6 BATTERIES AND FUEL CELLS
Primary Batteries
For a dry cell, what is the anode reaction?
What is the cathode reaction?
The alkaline battery
The mercury battery
Secondary Batteries
Storage batteries can be recharged.
The lead storage battery
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What is the anode reaction?
What is the cathode reaction?
Another recharge battery is the nickel-cadmium battery.
Fuel Cells
A fuel cell is an electrochemical cell in which the reactants are supplied to the cell from an external
source. For example, consider the hydrogen-oxygen cell
21.7 CORROSION: REDOX REACTIONS IN THE ENVIRONMENT
Corrosion is the deterioration of metals in an ox-re rxn. Ex: rust
For corrosion to occur:
1. must be anodic areas where the metal is ox
2. cathodic areas where some subsis reduces (often water)
3. an electrical connection between the two areas
4. electrolyte soln connecting the two areas
How can corrosion of iron be eliminated?
1. Inhibiting the anodic process by painting the metal surface or by allowing a thin oxide to form.
2. Inhibiting the cathodic process by forcing the metal to become the cathode by coating it with a
more readily oxidized metal as in the galvanizing of iron.
3. Best example is galvanized iron. In this case zinc coating forms a sacrificial anode.
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21.8 ELECTROLYSIS: CHEMICAL CHANGE FROM ELECTRIC ENERGY
Electrolysis is the use of electric current to bring about a chemical change.
In electrolysis as well as in a battery, oxidation occurs at the anode and reduction occurs at the cathode.
The electrolysis of molten NaCl is represented by the following half reactions:
Anode, ox
Cathode, re
Net Rxn
PROBLEM-SOLVING TIPS AND IDEAS Page 986
21.2 A Summary of Electrochemical Terminology
Whether you are describing a voltaic cell or an electrolysis cell, the terms “anode” and “cathode”
always refer to the electrodes at which oxidation and reduction occur, respectively. The polarity of the
electrodes is reversed, however, in a voltaic or electrolysis cell.
Type of Cell Electrode
Function
Voltaic
Oxidation
Reduction
Oxidation
Reduction
Electrolysis
Anode
Cathode
Anode
Cathode
Polarity
+
+
-
In a voltaic cell, the negative electrode is the one at which electrons are produced. In an electrolysis
cell, the negative electrode is the one onto which the external source is “pumping” the electrons.
The electrode reactions which will most likely occur will be those requiring the least potential to
overcome their nonspontaneity and the one which occurs most rapidly. The electrode reactions which
will most likely occur will be those requiring the least potential to overcome their nonspontaneity and
the one which occurs most rapidly.
Consider the electrolysis of an aqueous solution of sodium iodide. What are the possible reduction
half –reactions. SEE PG 987-988 for details and rxn
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EXERCISE 21.11 Page 989 Electrolysis of Salts
Predict the results of passing an electric current through each of the following solutions:
1. Molten NaBr
2. Aqueous NaBr
3. Aqueous CuCl2
21.9
COUNTING ELECTRONS
What is the relationship between amps (current), time, and coulombs?
To reduce 1 mol of Ag ions, 1 mol of e- is required.
1 mol of e- = 1 Faraday = 96485 coulombs (charge passed when 1 amp of current flows for 1 sec)
quantity of mass easily determined by mass difference. Current and time easily measured.
Relationship between these quantities is the molar mass of the metal and stoich of ox or reduction
(how many e- are gained or lost ie charge on the ion)
Time (sec) x current (amps)  charge (coulmbs)  moles (electrons)
Moles electrons  moles of metal  mass of metal
EQUATION 21.4 Current ,I ( amperes ,A ) 
electric ch arg e ( coulombs ,C )
time (sec onds ,s )
Exercise 21.12 Page 991 Using the Faraday Constant
In the commercial production of sodium by electrolysis, the cell operates at 7.0 V and a current of 25 x
103A. How many grams of sodium can be produced in one hour?
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21.10 THE COMMERCIAL PRODUCTION OF CHEMICALS BY ELECTROCHEMICAL
METHODS
Aluminum Read Pages 991-992
Chlorine and Sodium Hydroxide. Read Pages 992-994
What are some of the big uses of chlorine?
EXERCISE 21.13 Page 994 Producing Sodium
Sodium metal is produced by electrolysis from molten sodium chloride. The cell operates at 7.0 V
with a current of 25 x 103 A. How many kilowatt-hours of electricity are used to produce 1.00 kg of
sodium metal?
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