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18.5 Electrochemistry: An Introduction
599
Salt bridge
Porous disk
a
Figure 18.4
A salt bridge or a porous-disk
connection allows ions to flow,
completing the electric circuit.
The name galvanic cell honors
Luigi Galvani (1737–1798),
an Italian scientist generally
credited with the discovery
of electricity. These cells are
sometimes called voltaic
cells after Alessandro Volta
(1745–1827), another Italian,
who first constructed cells of
this type around 1800.
Anode: The electrode where
oxidation occurs. Cathode: The
electrode where reduction occurs.
b
The salt bridge contains a strong
electrolyte either as a gel or as
a solution; both ends are covered
with a membrane that allows
only ions to pass.
The porous disk allows ion flow
but does not permit overall mixing
of the solutions in the two
compartments.
idizing agent, and ions in the two aqueous solutions flow from one compartment to the other to keep the net charge zero.
Thus an electrochemical battery, also called a galvanic cell, is a
device powered by an oxidation–reduction reaction where the oxidizing
agent is separated from the reducing agent so that the electrons must travel
through a wire from the reducing agent to the oxidizing agent (Figure 18.5).
Notice that in a battery, the reducing agent loses electrons (which flow
through the wire toward the oxidizing agent) and so is oxidized. The electrode where oxidation occurs is called the anode. At the other electrode, the
oxidizing agent gains electrons and is thus reduced. The electrode where reduction occurs is called the cathode.
We have seen that an oxidation–reduction reaction can be used to generate an electric current. In fact, this type of reaction is used to produce elec-
e–
e–
e–
Cathode
(reduction)
Anode
(oxidation)
–
Ions
Corbis-Bettmann
Oxidizing
agent
Figure 18.5
Alessandro Volta.
Schematic of a battery (galvanic cell).
+
Reducing
agent
600 Chapter 18 Oxidation–Reduction Reactions and Electrochemistry
tric currents in many space vehicles. An oxidation–reduction reaction that
can be used for this purpose is hydrogen and oxygen reacting to form water.
2H2(g) ϩ O2(g) n 2H2O(l)
Oxidation states:
c
0
c
0
a
Q
ϩ1
Ϫ2
(each H)
Notice from the changes in oxidation states that in this reaction, hydrogen
is oxidized and oxygen reduced. The opposite process can also occur. We can
force a current through water to produce hydrogen and oxygen gas.
Electrical
2H2O(l) 87777777n
2H2(g) ϩ O2(g)
energy
This process, where electrical energy is used to produce a chemical change, is
called electrolysis.
In the remainder of this chapter, we will discuss both types of electrochemical processes. In the next section we will concern ourselves with the
practical galvanic cells we know as batteries.
18.6 Batteries
OBJECTIVE:
To learn about the composition and operation of commonly used batteries.
In the previous section we saw that a galvanic cell is a device that uses an
oxidation–reduction reaction to generate an electric current by separating
the oxidizing agent from the reducing agent. In this section we will consider
several specific galvanic cells and their applications.
▲
Remember: The oxidizing agent
accepts electrons and the
reducing agent furnishes
electrons.
Lead Storage Battery
Since about 1915, when self-starters were first used in automobiles, the lead
storage battery has been a major factor in making the automobile a practical means of transportation. This type of battery can function for several
years under temperature extremes from Ϫ30 ЊF to 100 ЊF and under incessant
punishment from rough roads. The fact that this same type of battery has
been in use for so many years in the face of all of the changes in science and
technology over that span of time attests to how well it does its job.
In the lead storage battery, the reducing agent is lead metal, Pb, and the
oxidizing agent is lead(IV) oxide, PbO2. We have already considered a simplified version of this reaction in Example 18.6. In an actual lead storage battery, sulfuric acid, H2SO4, furnishes the Hϩ needed in the reaction; it also furnishes SO42Ϫ ions that react with the Pb2ϩ ions to form solid PbSO4. A
schematic of one cell of the lead storage battery is shown in Figure 18.6.
In this cell the anode is constructed of lead metal, which is oxidized. In
the cell reaction, lead atoms lose two electrons each to form Pb2ϩ ions, which
combine with SO42Ϫ ions present in the solution to give solid PbSO4.
The cathode of this battery has lead(IV) oxide coated onto lead grids.
Lead atoms in the ϩ4 oxidation state in PbO2 accept two electrons each (are
reduced) to give Pb2ϩ ions that also form solid PbSO4.
In the cell the anode and cathode are separated (so that the electrons
must flow through an external wire) and bathed in sulfuric acid. The halfreactions that occur at the two electrodes and the overall cell reaction are
shown on the following page.
18.6 Batteries
601
Figure 18.6
In a lead storage battery each
cell consists of several lead grids
that are connected by a metal
bar. These lead grids furnish
electrons (the lead atoms lose
electrons to form Pb2ϩ ions,
which combine with SO42Ϫ ions
to give solid PbSO4). Because the
lead is oxidized, it functions as
the anode of the cell. The
substance that gains electrons is
PbO2; it is coated onto lead grids,
several of which are hooked
together by a metal bar. The
PbO2 formally contains Pb4ϩ,
which is reduced to Pb2ϩ, which
in turn combines with SO42Ϫ to
form solid PbSO4. The PbO2
accepts electrons, so it functions
as the cathode.
e– flow
H2SO4
electrolyte
solution
Pb metal grid
(anode)
PbO2 coated
onto a lead grid
(cathode)
Anode reaction:
Pb ϩ H2SO4 S PbSO4 ϩ 2Hϩ ϩ 2eϪ oxidation
Cathode reaction:
PbO2 ϩ H2SO4 ϩ 2eϪ ϩ 2Hϩ S PbSO4 ϩ 2H2O reduction
Overall reaction:
Pb(s) ϩ PbO2(s) ϩ 2H2SO4(aq) S 2PbSO4(s) ϩ 2H2O(l)
The tendency for electrons to flow from the anode to the cathode in a
battery depends on the ability of the reducing agent to release electrons and
on the ability of the oxidizing agent to capture electrons. If a battery consists
of a reducing agent that releases electrons readily and an oxidizing agent
with a high affinity for electrons, the electrons are driven through the connecting wire with great force and can provide much electrical energy. It is
useful to think of the analogy of water flowing through a pipe. The greater
the pressure on the water, the more vigorously the water flows. The “pressure” on electrons to flow from one electrode to the other in a battery is
called the potential of the battery and is measured in volts. For example,
each cell in a lead storage battery produces about 2 volts of potential. In an
actual automobile battery, six of these cells are connected to produce about
12 volts of potential.
▲
Dry Cell Batteries
The calculators, electronic watches, CD players, and MP3 players that are so
familiar to us are all powered by small, efficient dry cell batteries. They
are called dry cells because they do not contain a liquid electrolyte. The common dry cell battery was invented more than 100 years ago by George
Leclanché (1839–1882), a French chemist. In its acid version, the dry cell battery contains a zinc inner case that acts as the anode and a carbon (graphite)
rod in contact with a moist paste of solid MnO2, solid NH4Cl, and carbon
602 Chapter 18 Oxidation–Reduction Reactions and Electrochemistry
Cathode
(graphite rod)
Cathode (steel)
Insulation
Anode (zinc container)
Paste of HgO (oxidizing
agent) in a basic medium
of KOH and Zn(OH)2
Figure 18.8
A mercury battery of the type used in small
calculators.
Anode (zinc
inner case)
Paste of MnO2 ,
NH4Cl, and carbon
that acts as the cathode (Figure 18.7). The half-cell reactions are complex but
can be approximated as follows:
Anode reaction:
Zn S Zn2ϩ ϩ 2eϪ oxidation
Cathode reaction:
Figure 18.7
A common dry cell battery.
Ϫ
reduction
2NHϩ
4 ϩ 2MnO2 ϩ 2e S Mn2O3 ϩ 2NH3 ϩ H2O
This cell produces a potential of about 1.5 volts.
In the alkaline version of the dry cell battery, the NH4Cl is replaced with
KOH or NaOH. In this case the half-reactions can be approximated as follows:
Anode reaction:
Zn ϩ 2OHϪ S ZnO(s) ϩ H2O ϩ 2eϪ oxidation
Cathode reaction: 2MnO2 ϩ H2O ϩ 2eϪ S Mn2O3 ϩ 2OHϪ reduction
The alkaline dry cell lasts longer, mainly because the zinc anode corrodes less
rapidly under basic conditions than under acidic conditions.
Other types of dry cell batteries include the silver cell, which has a Zn anode and a cathode that employs Ag2O as the oxidizing agent in a basic environment. Mercury cells, often used in calculators, have a Zn anode and a cathode involving HgO as the oxidizing agent in a basic medium (see Figure 18.8).
An especially important type of dry cell is the nickel–cadmium battery, in
which the electrode reactions are
Anode reaction:
Cd ϩ 2OHϪ S Cd(OH)2 ϩ 2eϪ oxidation
Cathode reaction: NiO2 ϩ 2H2O ϩ 2eϪ S Ni(OH)2 ϩ 2OHϪ reduction
In this cell, as in the lead storage battery, the products adhere to the electrodes. Therefore, a nickel–cadmium battery can be recharged an indefinite
number of times, because the products can be turned back into reactants by
the use of an external source of current.
18.7 Corrosion
OBJECTIVE:
To understand the electrochemical nature of corrosion and to learn some
ways of preventing it.
Most metals are found in nature in compounds with nonmetals such as oxygen and sulfur. For example, iron exists as iron ore (which contains Fe2O3
and other oxides of iron).
C H E M I S T R Y I N F OCUS
Stainless Steel: It’s the Pits
O
Built in New York in
1929, the Chrysler
Building’s stainless
steel pinnacle has
been cleaned only a
few times. Despite
the urban setting, the
material shows few
signs of corrosion.
Image not available due to copyright restrictions
Dreamtime
ne of New York’s giants, the Chrysler Building, boasts a much admired art-deco stainless
steel pinnacle that has successfully resisted corrosion since it was built in 1929. Stainless steel is
the nobility among steels. Consisting of iron,
chromium (at least 13%), and nickel (with molybdenum and titanium added to more expensive
types), stainless steel is highly resistant to the
rusting that consumes regular steel. However,
the cheaper grades of stainless steel have an
Achilles heel—pit corrosion. In certain environments, pit corrosion can penetrate several millimeters in a matter of weeks.
Metallurgy, the science of producing useful
metallic materials, almost always requires some
kind of compromise. In the case of stainless steel,
inclusions of MnS make the steel easier to machine
into useful parts, but such inclusions are also the
source of pit corrosion. Recently a group of British
researchers analyzed stainless steel using a highenergy beam of ions that blasted atoms loose
from the steel surface. Studies of the resultant
atom vapor revealed the source
of the problem. It
turns out that
when the stainless
steel is cooling,
the MnS inclusions
“suck’’ chromium
atoms from the
surrounding area,
leaving behind a
chromium-deficient region. The corrosion occurs
in this region, as illustrated in the accompanying
diagram. The essential problem is that to resist corrosion steel must contain at least 13% Cr atoms.
The low-chromium region around the inclusion is
not stainless steel—so it corrodes just like regular
steel. This corrosion leads to a pit that causes deterioration of the steel surface.
Now that the reason for the pit corrosion is
understood, metallurgists should be able to develop methods of stainless steel formulation that
avoid this problem. One British scientist, Mary
P. Ryan, suggests that heat treatment of the
stainless steel may solve the problem by allowing
Cr atoms to diffuse from the inclusion back into
the surrounding area. Because corrosion of regular steel is such an important issue, finding ways
to make cheaper stainless steel will have a significant economic impact. We need stainless without the pits.
603
604 Chapter 18 Oxidation–Reduction Reactions and Electrochemistry
Corrosion can be viewed as the process of returning metals to their
natural state—the ores from which they were originally obtained. Corrosion
involves oxidation of the metal. Because corroded metal often loses its
strength and attractiveness, this process causes great economic loss. For example, approximately one-fifth of the iron and steel produced annually is
used to replace rusted metal.
Because most metals react with O2, we might expect them to corrode so
fast in air that they wouldn’t be useful. It is surprising, therefore, that the
problem of corrosion does not virtually prevent the use of metals in air. Part
of the explanation is that most metals develop a thin oxide coating, which
tends to protect their internal atoms against further oxidation. The best example of this is aluminum. Aluminum readily loses electrons, so it should be
very easily oxidized by O2. Given this fact, why is aluminum so useful for
building airplanes, bicycle frames, and so on? Aluminum is such a valuable
structural material because it forms a thin adherent layer of aluminum oxide, Al2O3, which greatly inhibits further corrosion. Thus aluminum protects
itself with this tough oxide coat. Many other metals, such as chromium,
nickel, and tin, do the same thing.
Iron can also form a protective oxide coating. However, this oxide is
not a very effective shield against corrosion, because it scales off easily, exposing new metal surfaces to oxidation. Under normal atmospheric conditions, copper forms an external layer of greenish copper sulfate or carbonate
called patina. Silver tarnish is silver sulfide, Ag2S, which in thin layers gives
the silver surface a richer appearance. Gold shows no appreciable corrosion
in air.
Preventing corrosion is an important way of conserving our natural
supplies of metals and energy. The primary means of protection is the application of a coating—most often paint or metal plating—to protect the metal
from oxygen and moisture. Chromium and tin are often used to plate steel
because they oxidize to form a durable, effective oxide coating.
Alloying is also used to prevent corrosion. Stainless steel contains
chromium and nickel, both of which form oxide coatings that protect the
steel.
Cathodic protection is the method most often employed to protect
steel in buried fuel tanks and pipelines. A metal that furnishes electrons
more easily than iron, such as magnesium, is connected by a wire to the
pipeline or tank that is to be protected (Figure 18.9). Because the magnesium
is a better reducing agent than iron, electrons flow through the wire from the
magnesium to the iron pipe. Thus the electrons are furnished by the magnesium rather than by the iron, keeping the iron from being oxidized. As oxidation of the magnesium occurs, the magnesium dissolves, so it must be replaced periodically.
Image not
available due
to copyright
restrictions
18.8 Electrolysis
OBJECTIVE:
To understand the process of electrolysis and learn about the commercial
preparation of aluminum.
Unless it is recharged, a battery “runs down” because the substances in it that
furnish and accept electrons (to produce the electron flow) are consumed.
For example, in the lead storage battery (see Section 18.6), PbO2 and Pb are
consumed to form PbSO4 as the battery runs.
PbO2(s) ϩ Pb(s) ϩ 2H2SO4(aq) S 2PbSO4(s) ϩ 2H2O(l)
C H E M I S T R Y I N F OCUS
Water-Powered Fireplace
H
An electrolytic cell uses
electrical energy to produce a
chemical change that would not
otherwise occur.
fuel. The Aqueon fireplace
uses electrolysis to decompose the water to H2(g) and
O2(g); the hydrogen is then
burned to furnish heat for
the home. The 31,000-Btu
fireplace features copper and
stainless steel and has a contemporary design (see accompanying photo). To operate,
the fireplace is simply hooked
up to the water and electrical
supplies for the home.
Courtesy, Hearth & Home Technologies
ydrogen gas is being touted
as an environmentally friendly
fuel because, unlike fossil fuels,
it does not produce the greenhouse gas carbon dioxide. The
only product of combustion of
H2 is water. As a result, hydrogen is being investigated as a
possible fuel for cars, trucks,
and buses. Now comes a manufacturer, Heat & Glo, that is
showcasing an in-home fireplace that uses water as the
However, one of the most useful characteristics of the lead storage battery is
that it can be recharged. Forcing current through the battery in the direction
opposite to the normal direction reverses the oxidation–reduction reaction.
That is, PbSO4 is consumed and PbO2 and Pb are formed in the charging
process. This recharging is done continuously by the automobile’s alternator, which is powered by the engine.
The process of electrolysis involves forcing a current through a cell to
produce a chemical change that would not otherwise occur.
One important example of this type of process is the electrolysis of water. Water is a very stable substance that can be broken down into its elements by using an electric current (Figure 18.10).
Forced
electric
Charles D. Winters/Photo Researchers, Inc.
2H2O(l) 87777777n
2H2(g) ϩ O2(g)
current
Figure 18.10
The electrolysis of water
produces hydrogen gas at the
cathode (on the left) and oxygen
gas at the anode (on the right).
A nonreacting strong electrolyte
such as Na2SO4 is needed to
furnish ions to allow the flow
of current.
The electrolysis of water to produce hydrogen and oxygen occurs
whenever a current is forced through an aqueous solution. Thus, when the
lead storage battery is charged, or “jumped,” potentially explosive mixtures
of H2 and O2 are produced by the current flow through the solution in the
battery. This is why it is very important not to produce a spark near the battery during these operations.
Another important use of electrolysis is in the production of metals
from their ores. The metal produced in the greatest quantities by electrolysis
is aluminum.
Aluminum is one of the most abundant elements on earth, ranking
third behind oxygen and silicon. Because aluminum is a very reactive metal,
it is found in nature as its oxide in an ore called bauxite (named after Les
Baux, France, where it was discovered in 1821). Production of aluminum
metal from its ore proved to be more difficult than the production of most
other metals. In 1782, Lavoisier, the pioneering French chemist, recognized
aluminum as a metal “whose affinity for oxygen is so strong that it cannot
be overcome by any known reducing agent.” As a result, pure aluminum
605
606 Chapter 18 Oxidation–Reduction Reactions and Electrochemistry
Charles Martin Hall (1863–1914) was a
student at Oberlin College in Ohio when
he first became interested in aluminum.
One of his professors commented that
anyone who could find a way to
manufacture aluminum cheaply would
make a fortune, and Hall decided to give
it a try. The 21-year-old worked in a
wooden shed near his house with an iron
frying pan as a container, a blacksmith’s
forge as a heat source, and galvanic cells
constructed from fruit jars. Using these
crude galvanic cells, Hall found that he
could produce aluminum by passing a
current through a molten mixture of Al2O3
and Na3AIF6. By a strange coincidence,
Paul Heroult, a French chemist who was
born and died in the same years as Hall,
made the same discovery at about the
same time.
The Granger Collection
Figure 18.11
Table 18.2 The Price of
Aluminum,
1855–1990
Date
Price of Aluminum
($/lb)*
1855
$100,000
1885
100
1890
2
1895
0.50
1970
0.30
1980
0.80
1990
0.74
*Note the precipitous drop in price after
the discovery of the Hall–Heroult process
in 1886.
C H A P T E R
metal remained unknown. Finally, in 1854, a process was found for producing metallic aluminum by using sodium, but aluminum remained a very expensive rarity. In fact, it is said that Napoleon III served his most honored
guests with aluminum forks and spoons, while the others had to settle for
gold and silver utensils!
The breakthrough came in 1886 when two men, Charles M. Hall in the
United States (Figure 18.11) and Paul Heroult in France, almost simultaneously discovered a practical electrolytic process for producing aluminum,
which greatly increased the availability of aluminum for many purposes.
Table 18.2 shows how dramatically the price of aluminum dropped after this
discovery. The effect of the electrolysis process is to reduce Al3ϩ ions to neutral Al atoms that form aluminum metal. The aluminum produced in this
electrolytic process is 99.5% pure. To be useful as a structural material, aluminum is alloyed with metals such as zinc (for trailer and aircraft construction) and manganese (for cooking utensils, storage tanks, and highway
signs). The production of aluminum consumes about 4.5% of all electricity
used in the United States.
18
REVIEW
Key Terms
oxidation–reduction
(redox) reactions (18.1)
oxidation (18.1, 18.3)
reduction (18.1, 18.3)
oxidation states (18.2)
oxidizing agent (electron
acceptor) (18.3)
reducing agent (electron
donor) (18.3)
half-reactions (18.4)
electrochemistry (18.5)
electrochemical battery
(galvanic cell) (18.5)
anode (18.5)
cathode (18.5)
electrolysis (18.5, 18.8)
lead storage battery (18.6)
potential (18.6)
dry cell battery (18.6)
corrosion (18.7)
cathodic protection (18.7)
Chapter Review
F
VP
directs you to the Chemistry in Focus feature in the chapter
indicates visual problems
interactive versions of these problems are assignable in OWL
607
1. Sketch a galvanic cell, and explain how it works. Look
at Figures 18.1 and 18.5. Explain what is occurring in
each container and why the cell in Figure 18.5
“works,” but the one in Figure 18.1 does not.
2. Make a list of nitrogen compounds with as many different oxidation states for nitrogen as you can.
Summary
1. Oxidation–reduction reactions involve a transfer of
electrons. Oxidation states provide a way to keep
track of electrons in these reactions. A set of rules is
used to assign oxidation states.
2. Oxidation is an increase in oxidation state (a loss of
electrons); reduction is a decrease in oxidation state
(a gain of electrons). An oxidizing agent accepts electrons, and a reducing agent donates electrons. Oxidation and reduction always occur together.
3. Oxidation–reduction equations can be balanced by
inspection or by the half-reaction method. This
method involves splitting a reaction into two parts
(the oxidation half-reaction and the reduction halfreaction).
4. Electrochemistry is the study of the interchange of
chemical and electrical energy that occurs through
oxidation–reduction reactions.
5. When an oxidation–reduction reaction occurs with
the reactants in the same solution, the electrons are
transferred directly, and no useful work can be obtained. However, when the oxidizing agent is separated from the reducing agent, so that the electrons
must flow through a wire from one to the other, chemical energy is transformed into electrical energy. The
opposite process, in which electrical energy is used to
produce chemical change, is called electrolysis.
6. A galvanic (electrochemical) cell is a device in which
chemical energy is transformed into useful electrical
VP
energy. Oxidation occurs at the anode of a cell; reduction occurs at the cathode.
7. A battery is a galvanic cell, or group of cells, that
serves as a source of electric current. The lead storage
battery has a lead anode and a cathode of lead coated
with PbO2, both immersed in a solution of sulfuric
acid. Dry cell batteries do not have liquid electrolytes
but contain a moist paste instead.
8. Corrosion involves the oxidation of metals to form
mainly oxides and sulfides. Some metals, such as aluminum, form a thin protective oxide coating that inhibits their further corrosion. Corrosion of iron can
be prevented by a coating (such as paint), by alloying,
and by cathodic protection.
3. Which of the following are oxidation–reduction reactions? Explain.
a.
b.
c.
d.
e.
PCl3 ϩ Cl2 S PCl5
Cu ϩ 2AgNO3 S Cu(NO3)2 ϩ 2Ag
CO2 ϩ 2LiOH S Li2CO3 ϩ H2O
FeCl2 ϩ 2NaOH S Fe(OH)2 ϩ 2NaCl
MnO2 ϩ 4HCl S Cl2 ϩ 2H2O ϩ MnCl2
4. Which of the following statements is (are) true? Explain. (There may be more than one answer.)
a. Oxidation and reduction cannot occur independently of each other.
b. Oxidation and reduction accompany all chemical
changes.
c. Oxidation and reduction describe the loss and
gain of electron(s), respectively.
5. Why do we say that when something gains electrons
it is reduced? What is being reduced?
6. The equation Agϩ ϩ Cu S Cu2ϩ ϩ Ag has equal numbers of each type of element on each side of the equation. This equation, however, is not balanced. Why is
this equation not balanced? Balance the equation.
7. In balancing oxidation–reduction equations, why is
it permissible to add water to either side of the equation?
8. What does it mean for a substance to be oxidized? The
term “oxidation” originally came from substances reacting with oxygen gas. Explain why a substance that
reacts with oxygen gas will always be oxidized.
9. Label the following parts of the galvanic cell.
anode
cathode
reducing agent
oxidizing agent
e–
e–
e–
A
C
Active Learning Questions
These questions are designed to be considered by groups of
students in class. Often these questions work well for introducing a particular topic in class.
–
B
Ions
+
D
608 Chapter 18 Oxidation–Reduction Reactions and Electrochemistry
Questions and Problems
18.1 Oxidation–Reduction Reactions
QUESTIONS
10. Why is fluorine always assigned an oxidation state of
Ϫ1? What oxidation number is usually assigned to the
other halogen elements when they occur in compounds? In an interhalogen compound involving fluorine
(such as ClF), which atom has a negative oxidation state?
1. Give some examples of how we make good use of
oxidation–reduction reactions in everyday life.
11. The sum of all the oxidation states of all the atoms in
H3PO4 is
.
2. How do chemists define the processes of oxidation
and reduction? Write a simple equation illustrating
each of your definitions.
12. The sum of all the oxidation states of all the atoms in
PO43Ϫ is
.
3. For each of the following oxidation–reduction reactions, identify which element is being oxidized and
which is being reduced.
a.
b.
c.
d.
Cl2(g) ϩ I2(g) S 2ICl(g)
Cl2(g) ϩ 2Li(s) S 2LiCl(s)
2Na(s) ϩ 2H2O(l) S 2NaOH(aq) ϩ H2(g)
Cl2(g) ϩ 2NaBr(aq) S 2NaCl(aq) ϩ Br2(l)
4. For each of the following oxidation–reduction reactions, identify which element is being oxidized and
which is being reduced.
a.
b.
c.
d.
4B(s) ϩ 3O2(g) S 2B2O3(s)
N2(g) ϩ 2O2(g) S 2NO2(g)
CaC2(s) ϩ H2(g) S CaH2(g) ϩ 2C(s)
CuSO4(aq) ϩ Mg(s) S MgSO4(aq) ϩ Cu(s)
5. For each of the following oxidation–reduction reactions, identify which element is being oxidized and
which is being reduced.
a.
b.
c.
d.
Ca(s) ϩ 2H2O(l) S Ca(OH)2(s, aq) ϩ H2(g)
H2(g) ϩ F2(g) S 2HF(g)
4Fe(s) ϩ 3O2(g) S 2Fe2O3(s)
2Fe(s) ϩ 3Cl2(g) S 2FeCl3(s)
6. For each of the following oxidation–reduction reactions, identify which element is being oxidized and
which is being reduced.
a.
b.
c.
d.
2Cr2S3(s) ϩ 3O2(g) S 2Cr2O3(s) ϩ 6S(s)
P4(s) ϩ 5O2(g) S P4O10(s)
CO2(g) ϩ H2(g) S CO(g) ϩ H2O(g)
2B(s) ϩ 3H2(g) S B2H6(g)
18.2 Oxidation States
QUESTIONS
7. What is an oxidation state? Why do we define such a
concept?
8. Why is the oxidation state of an element in the uncombined state equal to zero?
9. Explain why, although it is not an ionic compound,
we still assign oxygen an oxidation state of Ϫ2 in water, H2O. Give an example of a compound in which
oxygen is not in the Ϫ2 oxidation state.
PROBLEMS
13. Assign oxidation states to all of the atoms in each of
the following.
a. CBr4
b. HClO4
c. K3PO4
d. N2O
14. Assign oxidation states to all of the atoms in each of
the following.
a. CrCl3
b. Ni(OH)2
c. H2S
d. CS2
15. What is the oxidation state of sulfur in each of the following substances?
a. S8
b. H2SO4
c. NaHSO4
d. Na2S
16. What is the oxidation state of nitrogen in each of the
following substances?
a. N2
b. NH3
c. NO2
d. NaNO3
17. What is the oxidation state of chlorine in each of the
following substances?
a. ClF
b. Cl2
c. HCl
d. HClO
18. What is the oxidation state of manganese in each of
the following substances?
a. MnCl2
b. KMnO4
c. MnO2
d. Mn(C2H3O2)3
19. Assign oxidation states to all of the atoms in each of
the following.
a. CuCl2
b. KClO3
c. KClO4
d. Na2CO3
20. Assign oxidation states to all of the atoms in each of
the following.
a. CaO
b. Al2O3
c. PF3
d. P2O5
21. Assign oxidation states to all of the atoms in each of
the following ions:
a. CO32Ϫ
b. NO3Ϫ
c. PO43Ϫ
d. SO42Ϫ
All even-numbered Questions and Problems have answers in the back of this book and solutions in the Solutions Guide.
Chapter Review
22. Assign oxidation states to all of the atoms in each of
the following ions:
a. HSO4Ϫ
b. MnO4Ϫ
c. ClO3Ϫ
d. BrO4Ϫ
18.3 Oxidation–Reduction Reactions
Between Nonmetals
QUESTIONS
23. Oxidation can be defined as a loss of electrons or as
an increase in oxidation state. Explain why the two
definitions mean the same thing, and give an example to support your explanation.
24. Reduction can be defined as a gain of electrons or as
a decrease in oxidation state. Explain why the two definitions mean the same thing, and give an example
to support your explanation.
25. What is an oxidizing agent? What is a reducing agent?
26. Give an example of a simple oxidation–reduction
equation. Identify the species being oxidized and the
species being reduced. Identify the oxidizing agent
and the reducing agent in your example.
27. Does an oxidizing agent donate or accept electrons?
Does a reducing agent donate or accept electrons?
F 28. The “Chemistry in Focus” segment Do We Age by Oxi-
dation? discusses antioxidants. What does it mean for
a chemical to be an antioxidant? How would it work
chemically?
PROBLEMS
29. In each of the following reactions, identify which element is being oxidized and which is being reduced
by assigning oxidation numbers.
a. Fe(s) ϩ CuSO4(aq) S FeSO4(aq) ϩ Cu(s)
b. Cl2(g) ϩ 2NaBr(aq) S 2NaCl(aq) ϩ Br2(l)
c. 3CuS(s) ϩ 8HNO3(aq) S
3CuSO4(aq) ϩ 8NO(g) ϩ 4H2O(l)
d. 2Zn(s) ϩ O2(g) S 2ZnO(s)
30. In each of the following reactions, identify which element is being oxidized and which is being reduced
by assigning oxidation numbers.
a.
b.
c.
d.
2Al(s) ϩ 3S(s) S Al2S3(s)
CH4(g) ϩ 2O2(g) S CO2(g) ϩ 2H2O(g)
2Fe2O3(s) ϩ 3C(s) S 3CO2(g) ϩ 4Fe(s, l)
K2Cr2O7(aq) ϩ 14HCl(aq) S
2KCl(aq) ϩ 2CrCl3(s) ϩ 7H2O(l) ϩ 3Cl2(g)
31. In each of the following reactions, identify which element is being oxidized and which is being reduced
by assigning oxidation states.
a.
b.
c.
d.
2Cu(s) ϩ S(s) S Cu2S(s)
2Cu2O(s) ϩ O2(g) S 4CuO(s)
4B(s) ϩ 3O2(g) S 2B2O3(s)
6Na(s) ϩ N2(g) S 2Na3N(s)
609
32. In each of the following reactions, identify which element is being oxidized and which is being reduced
by assigning oxidation numbers.
a. 4KClO3(s) ϩ C6H12O6(s) S
4KCl(s) ϩ 6H2O(l) ϩ 6CO2(g)
b. 2C8H18(l) ϩ 25O2(g) S 16CO2(g) ϩ 18H2O(l)
c. PCl3(g) ϩ Cl2(g) S PCl5(g)
d. Ca(s) ϩ H2(g) S CaH2(g)
33. Pennies in the United States consist of a zinc core
that is electroplated with a thin coating of copper.
Zinc dissolves in hydrochloric acid, but copper does
not. If a small scratch is made on the surface of a
penny, it is possible to dissolve away the zinc core,
leaving only the thin shell of copper. Identify which
element is oxidized and which is reduced in the reaction for the dissolving of the zinc by the acid.
Zn(s) ϩ 2HCl(aq) S ZnCl2(aq) ϩ H2(g)
34. Iron ores, usually oxides of iron, are converted to the
pure metal by reaction in a blast furnace with carbon
(coke). The carbon is first reacted with air to form carbon monoxide, which in turn reacts with the iron oxides as follows:
F2O3(s) ϩ 3CO(g) S 2Fe(l) ϩ 3CO2(g)
Identify the atoms that are oxidized and reduced, and
specify the oxidizing and reducing agents.
35. Although magnesium metal does not react with water at room temperature, it does react vigorously with
steam at higher temperatures, releasing elemental hydrogen gas from the water.
Mg(s) ϩ 2H2O(g) S Mg(OH)2(s) ϩ H2(g)
Identify which element is being oxidized and which
is being reduced.
36. Potassium iodide in solution reacts readily with many
reagents. In the following reactions, identify the
atoms that are being oxidized and reduced, and specify the oxidizing and reducing agents.
a. Cl2(g) ϩ KI(aq) S KCl(aq) ϩ I2(s)
b. 2FeCl3(aq) ϩ 2KI(aq) S
2FeCl2(aq) ϩ 2KCl(aq) ϩ I2(s)
c. 2CuCl2(aq) ϩ 4KI(aq) S
2CuI(s) ϩ 4KCl(aq) ϩ I2(s)
18.4 Balancing Oxidation–Reduction Reactions
by the Half-Reaction Method
QUESTIONS
37. In what two respects must oxidation–reduction reactions be balanced?
38. Why is a systematic method for balancing oxidation–
reduction reactions necessary? Why can’t these equations be balanced readily by inspection?
All even-numbered Questions and Problems have answers in the back of this book and solutions in the Solutions Guide.