16
CHAPTER
Oxidation-Reduction
Reactions
Chapter Preview
Sections
16.1 The Nature of OxidationReduction Reactions
MiniLab 16.1 Corrosion of Iron
ChemLab
Copper Atoms and Ions:
Oxidation and Reduction
16.2 Applications of OxidationReduction Reactions
MiniLab 16.2 Testing for Alcohol
by Redox
552
Why Do Things Rust?
W
hen iron corrodes, iron metal reacts with
oxygen from the air and water to form
iron(III) oxide—rust. Rust is the result of
an oxidation-reduction reaction in which iron
metal loses electrons to oxygen. Given time, and
oxygen from the air and water, all of the drums
in this photo will rust away completely.
Start-up Activities
What I Already Know
Observing an Oxidation–
Reduction Reaction
Rust is the result of a reaction of iron and oxygen. Iron
nails can also react with substances other than oxygen,
as you will find out in this experiment.
Safety Precautions
Always wear safety goggles and an apron in the
laboratory.
Materials
• test tube
• iron nail
• steel wool or sandpaper
• 1M copper(II) sulfate (CuSO4)
Procedure
1. Use a piece of steel wool to polish the end of an
iron nail.
2. Add about 3 mL 1.0M CuSO4 to a test tube. Place the
polished end of the nail into the CuSO4 solution. Let
stand and observe for about 10 minutes. Record your
observations.
Review the following concepts
before studying this chapter.
Chapter 3: patterns of valence
electrons
Chapter 5: predicting oxidation
number from the periodic table
Chapter 6: types of reactions
Reading Chemistry
Look through the Section Previews
for this chapter, jotting down some
key ideas. As you read through the
chapter, make an outline using the
key ideas you wrote down. For each
topic, review any new vocabulary
words.
Preview this chapter’s content and
activities at chemistryca.com
Analysis
What is the substance found clinging to the nail? What
happened to the color of the copper(II) sulfate solution?
Write the balanced chemical equation for the reaction
you observed.
553
SECTION
16.1
The Nature of OxidationReduction Reactions
O
✓ Identify the substances that are oxidized and those that
are reduced in a redox
reaction.
xygen undergoes many reactions when it encounters other substances.
One of these reactions is responsible for the browning of fruits. Another forms the rust that eats away at the metal parts of bikes and cars. In
both of these cases, a type of reaction called oxidation is taking place. You
can probably guess how this reaction got its name; oxygen is a reactant. But
you will learn that not
all oxidation reactions
involve oxygen. And
oxidation reactions
are never lonely
because they always
have partners—
reduction reactions.
You will see what the
characteristics of
these reactions are
and why they always
take place together.
✓ Distinguish oxidizing and reducing
agents in redox
reactions.
What is oxidation-reduction?
SECTION PREVIEW
Objectives
✓ Analyze the characteristics of an
oxidation-reduction
reaction.
✓ Distinguish
between oxidation
reactions and reduction reactions by
definition.
Review Vocabulary
Buffer: solution that
resists changes in pH
when moderate
amounts of acids or
bases are added.
New Vocabulary
oxidation-reduction
reaction
oxidation
reduction
oxidizing agent
reducing agent
554
Chapter 16
Oxygen is the most abundant element in Earth’s crust. It is very reactive
and can combine with almost every other element. An element that bonds
to oxygen to form a new compound, called an oxide, usually loses electrons
because oxygen is more electronegative. You will recall that an electronegative element has a strong attraction for electrons. Because of this
strong attraction, oxygen is able to pull electrons away from other atoms.
The reactions in which elements combine with oxygen to form oxides were
among the first to be studied by early chemists, who grouped them together and called them oxidation reactions. Later, chemists realized that some
other nonmetal elements can combine with substances in the same way as
oxygen and that these reactions are similar to oxidation reactions. Modern
chemists use the term oxidation to refer to any chemical reaction in which
an element or compound loses electrons to another substance.
A common oxidation reaction occurs when iron metal loses electrons
to oxygen. Each year in the United States, corrosion of metals—especially
the iron in steel—costs billions of dollars as automobiles, ships, and
bridges and other structures are slowly eaten away. Figure 16.1 shows
some of this damage and how it can be prevented.
Oxidation-Reduction Reactions
Figure 16.1
Corrosion of Iron
When iron corrodes, iron metal reacts with oxygen to
form iron(III) oxide—rust. Corrosion of iron can be
prevented by covering the surface of exposed steel with
paint or other coatings such as plastic. If the protective
coating is damaged or cracked, rust forms quickly. 䊳
䊴 Steel can be protected from oxidation if it is coated with a
more active metal such as zinc. Zinc loses electrons to
oxygen more readily than iron does, so the zinc is oxidized preferentially, forming a tough protective layer of
zinc oxide. The coating of zinc and zinc oxide prevents
the formation of rust by keeping oxygen from reaching
the iron. Steel that has been coated with zinc is called galvanized steel. The bucket on the left has been galvanized.
Redox
What happens to the zinc in galvanized steel? It reacts with oxygen to
form zinc oxide in the following reaction.
2Zn(s) O2(g) ˇ 2Zn2(s) 2O2(s)
Does this type of reaction look familiar? You learned in Chapter 6 that this
is classified as a synthesis reaction. You also know that early chemists called
it an oxidation reaction because oxygen is a reactant. The formation of zinc
oxide falls into another, broader class of reactions characterized by the
transfer of electrons from one atom or ion to another. This type of reaction
is called an oxidation-reduction reaction, commonly known as a redox
reaction. Many important chemical reactions are redox reactions. Formation of rust is one example; combustion of fuels is another. In each redox
reaction, one element loses electrons, and another element takes them.
How do atoms or ions lose electrons in a redox reaction? If you examine
the equation for the reaction between zinc and oxygen more closely, you can
see which atoms are gaining electrons and which are losing them. You also
can determine where the electrons go during a redox reaction by comparing
the oxidation number of each type of atom or ion before and after the reaction takes place. Recall from Chapter 5 that the oxidation number of an ion
is equal to its charge. All elements, when in their free form, have a charge of
zero and are assigned an oxidation number of zero. In the formation of zinc
oxide, the zinc atom and the diatomic oxygen molecule that react each has
an oxidation number of zero. In the ionic compound formed, each oxide
ion has a 2 charge and an oxidation number of 2. Because the compound must be neutral, the total positive charge must be 4; thus, each zinc
ion must have a charge and an oxidation number of 2.
16.1
The Nature of Oxidation-Reduction Reactions
555
Oxidation
You have learned that a reaction in which an element loses electrons is
called an oxidation reaction. The element that loses the electrons becomes
more positively charged; that is, its oxidation number increases. That element is said to be oxidized during the reaction. Zinc is oxidized during the
formation of zinc oxide because metallic zinc atoms each lose two electrons.
The oxidation reaction can be written by itself to show how zinc changes
during the redox reaction. Here’s what happens to each atom of zinc.
2+
Zn ˇ [Zn] 2e –
reduction:
re (L) back
ducere (L) to lead
In a reduction
reaction, the addition of electrons
results in a
decrease in oxidation number of an
atom or ion.
(loss of electrons)
Reduction
What happens to the electrons that are lost by the zinc atom? Electrons
do not wander around by themselves; they must be transferred to another
atom or ion. This is why oxidation reactions never occur alone. They are
always paired with reduction reactions. A reduction reaction is one in
which an element gains one or more electrons. The element that picks up
the electrons and becomes more negatively charged during the reaction is
said to be reduced. Its oxidation number decreases, or is reduced. Because
oxidation and reduction reactions occur together, each is referred to as a
half-reaction.
In every redox reaction, at least one element undergoes reduction while
another undergoes oxidation. Just as a successful pass in football requires a
quarterback to throw the ball and a receiver to catch it, a redox reaction
must have one element that gives up electrons and one that accepts them.
The electronic structure of both reactants changes during a redox reaction.
Figure 16.2 shows the movement of electrons in the formation of zinc
oxide. Oxygen accepts the electrons that zinc loses. Oxygen is reduced
during the reaction between zinc and oxygen because each oxygen atom
gains two electrons. Like the oxidation reaction, the reduction reaction
can be written by itself. Here’s what happens to each atom of oxygen.
Figure 16.2
Formation of Zinc Oxide
In the formation of zinc
oxide, the zinc atom loses two
electrons during the reaction,
becoming a zinc ion. Its oxidation number increases from
zero to 2. The
oxygen atom gains
the two electrons
Zn
from zinc, becoming an oxide ion.
Its oxidation num2e 8e 18e
ber decreases from
zero to 2.
2e
O 2e – ˇ O
2–
(gain of electrons)
[Zn]2
2e 6e
2e 8e 18e
556
Chapter 16
2
2e 8e
Zinc
atom
Oxygen
atom
Oxidation-Reduction Reactions
Zinc
ion
Oxide
ion
Corrosion of Iron
Corrosion is the term generally used to describe the oxidation of a
metal during its interaction with the environment. In this MiniLab, you
will study the corrosion of a nail and determine the factors that affect
this process.
Procedure
1. Dissolve a package of clear,
unflavored gelatin in about
200 mL of warm water. Stir in
2 mL of phenolphthalein solution and 2 mL of potassium
hexacyanoferrate(III) solution.
Pour the prepared solution into
a widemouth glass jar or petri
dish to a depth of about 1 cm.
2. In the liquid gelatin, place a
plain iron nail, an aluminum
nail, a galvanized iron nail, and
a painted iron nail of the type
used for paneling. Space the
nails far apart.
3. Label the jar or petri dish with
your name, and leave it for several hours or overnight. Handle
the jar or dish carefully until
the gelatin has set.
4. Record your observations
regarding any interactions of
the nails with the substances in
the gelatin.
1
Analysis
1. Which of the nails have
reacted with the substances
in the gelatin? What is the
evidence of corrosion?
2. If any of the nails have not
corroded in the solution, can
you suggest a reason why
they haven’t? What methods
are commonly used to prevent or minimize corrosion?
3. Any blue color in the gelatin
is due to the formation of
iron(II) ions and their interaction with the hexacyanoferrate(III) ion. Any pink or
red color in the gelatin is due
to the gaining of electrons by
oxygen and water molecules,
forming basic hydroxide ions
that turn the phenolphthalein pink. Which of these
reactions is oxidation, and
where does it occur on the
reacting nail?
Combining the Half-Reactions
The equation for the reduction half-reaction shows one atom of oxygen
reacting. However, oxygen is not found in nature as single atoms; two
atoms combine to form a diatomic molecule of O2. The reduction equation must be multiplied by 2 to reflect this. Thus, the balanced equation
for the reduction reaction is written as follows.
O2 4e – ˇ 2 O
2–
Note that four electrons are gained by the oxygen molecule. To produce
those four electrons, two atoms of zinc must take part in the reaction.
Therefore, the balanced equation for the oxidation reaction must be written as follows.
2+
2Zn ˇ 2[Zn] 4e –
16.1
The Nature of Oxidation-Reduction Reactions
557
Figure 16.3
Summarizing the Reaction
Two zinc atoms combine with
one diatomic oxygen molecule to form two formula
units of zinc oxide. Because
zinc loses four electrons in the
oxidation reaction and oxygen gains four electrons in
the reduction reaction, all
electrons are accounted for;
the two half-reactions are
balanced.
Each atom of oxygen accepts 2e from a zinc atom and is reduced.
+4e 0
from Zn
2Zn O2 ˇ 2Zn 2O
0
0
2+
2–
–4e to O2
Each atom of zinc donates 2e to an oxygen atom and is oxidized.
The balanced overall equation for the reaction now can be written as
shown in Figure 16.3. This equation is the same as the equation for the
formation of zinc oxide that you read at the beginning of the discussion
on redox reactions. Now you know that it represents the net oxidationreduction reaction and is the sum of an oxidation half-reaction and a
reduction half-reaction.
If an element is gaining electrons, why is this called a reduction reaction? After all, you don’t gain weight when you reduce. You have learned
that the reason is because there is a reduction in the charge or oxidation
number of an atom of the substance that is reduced. An older, historic
reason for the use of the term reduction is that the name was first applied
to processes in furnaces in which metals are isolated from their ores at
high temperatures, Figure 16.4. During these processes, oxygen is
removed from the ores in which it is combined with the metal, so the ore
is reduced to the free metal. There is a reduction in the amount of solid
material and a considerable decrease in volume.
Figure 16.4
Furnaces: Old and Modern
For thousands of years, metals have been used by many different cultures
for making jewelry, cookware, and weapons. Because metals are normally
found combined with other elements as ores, furnaces operating at high
temperatures are used to separate the free metal from other elements. The
positively charged metal ions in the ore are reduced to the elemental state,
while oxygen and other negatively charged elements in the ore are oxidized. Iron smelting in medieval England is shown here. 䊲
䊱 A modern industrial iron blast
furnace is shown here.
558
Chapter 16
Oxidation-Reduction Reactions
Identifying a Redox Reaction
The oxidation of zinc is a redox reaction in which oxygen is a reactant.
You have learned that elements other than oxygen can accept electrons and
become reduced during redox reactions. You are already familiar with the
explosive reaction in which sodium and chlorine combine to form table salt.
2Na(s) Cl2(g) ˇ 2NaCl(s)
Are electrons transferred during this reaction? Yes, because each sodium
atom loses one electron to become a sodium ion with a charge of 1. The
oxidation number of sodium increases from 0 to 1. Each chlorine atom
gains one electron to form a chloride ion. The oxidation number of chlorine decreases from 0 to 1. Therefore, this is another example of a redox
reaction.
ˇ
–
2Na Cl 2 ˇ 2Na+ 2Cl
0
0
Figure 16.5
The Reaction Between Zinc
and Copper(II) Sulfate
A blue copper(II) sulfate solution gradually becomes colorless if a strip of zinc metal is
placed in it. The zinc gives up
electrons, becoming oxidized
to zinc ions. The colorless
Zn2 ions that form go into
solution. The Cu2 ions pick
up electrons from zinc and
become reduced to copper
metal atoms, which are
deposited on the strip. 䊲
Oxidizing and Reducing Agents
Another redox reaction that doesn’t involve oxygen occurs when a strip
of zinc metal is placed in a solution of copper(II) sulfate. The progress of
this reaction can be followed easily because a readily observable change
takes place. As shown in Figure 16.5, copper metal quickly begins to form
on the zinc strip.
2e
0
Cu reducedˇ
to Cu
ˇ
2+
Cu2 is the oxidizing agent, and
Zn0 is the reducing agent. 䊳
0
Zn
reduced to
Cu
2+
oxidizing
agent
Zn
0
reducing
agent
16.1
ˇ
ˇ
Cu0
ˇ 2e
ˇ Zn2+
oxidized to
2+
Zn
ˇ
2e –
oxidized to
The Nature of Oxidation-Reduction Reactions
559
Copper Atoms and Ions:
Oxidation and Reduction
Copper atoms and ions often take part in reactions by losing or gaining electrons, which are
oxidation and reduction, respectively. If copper
atoms lose electrons to form positive ions, copper
is oxidized. Other atoms or ions must gain the
electrons that copper atoms lose. These atoms or
ions are reduced and are called oxidizing agents.
In this ChemLab, you will observe two reactions
that involve the oxidation or reduction of copper.
Problem
What are some typical reactions that involve
the oxidation or reduction of copper?
Objectives
Observe reactions that involve the oxidation or
reduction of copper.
Classify the reactants as substance oxidized,
reducing agent, substance reduced, and oxidizing agent.
•
•
Materials
copper(II) oxide
powdered charcoal (carbon)
weighing paper
balance
large Pyrex or Kimax test tubes (2)
1-hole rubber stopper fitted with glass tube with
bend as shown
150-mL beakers (2)
graduated cylinder, small
glass stirring rod
Bunsen or Tirrill burner
ring stand
test-tube clamp
560
Chapter 16
Oxidation-Reduction Reactions
thermal glove
limewater (calcium hydroxide solution)
Safety Precautions
Care should be taken in handling hot objects
and when working around open flames. Do not
breathe in the fumes that are produced during
the teacher demonstration in step 1.
1. Teacher Demonstration Your teacher will
perform this reaction as a demonstration either
in the fume hood or outside the building.
CAUTION: Do not perform this procedure by
yourself. A 1-cm square
of copper foil will be
placed in a porcelain
evaporating dish. First,
5 mL of water, then
5 mL of concentrated
HNO3 will be added.
Note the color of the
evolved gas and the
color of the resulting solution. Record your
observations in a table similar to the one under
Data and Observations.
2. On a piece of weighing paper, thoroughly mix
approximately 1 g of copper(II) oxide with
twice its volume of powdered charcoal. Place
the mixture in a clean, dry Pyrex or Kimax test
tube. Add about 10 mL of limewater to a second test tube, and stand it in a 150-mL beaker.
Assemble the apparatus as shown here, with
the copper-oxide test tube sloped slightly
downward and the delivery tube extending
into the limewater.
3. Heat the mixture in the test tube, gently at first
and then strongly. As soon as you notice a
change in the limewater, carefully remove the
stopper and delivery tube from the reaction
test tube. CAUTION: Do not stop heating as
long as the tube is in the limewater. Record your
observations of the limewater in your table.
4. Continue heating the reaction test tube until a
glow spreads throughout the reactant mixture.
Turn off the burner.
5. After the reaction test tube has cooled to nearly room temperature, empty the contents into
a beaker that is about half full of water. In a
sink, slowly stir the mixture while running
water into the beaker until all the unreacted
charcoal has washed away. Observe the product that remains in the beaker, and record
your observations.
gas is NO. In the second reaction, if limewater
becomes cloudy and white, carbon dioxide gas
has reacted with the calcium hydroxide to
form insoluble calcium carbonate. Using this
information, analyze your data and observations. Determine which reactants (Cu and
HNO3 for the first reaction, CuO and C for
the second reaction) were oxidized and which
were reduced in each reaction.
3. Classifying Classify each of the four reactants
as an oxidizing agent or a reducing agent.
1. The mass of copper produced in the second
reaction is less than the mass of the reacting
copper(II) oxide. Why, then, is the gain of
electrons known as reduction?
2. What mass of copper may be produced from
the reduction of 1.000 metric ton of
copper(II) oxide? Hint: Determine the formula mass of copper(II) oxide.
3. If the chlorine gas used at a water-treatment
plant reacts with organic materials in the
water to yield chloride ions, how would you
classify the chlorine gas in terms of oxidation
and reduction?
Observations
1. Interpreting Data What evidence of chemical
change did you observe in each reaction?
2. Interpreting Data In the first reaction, a bluecolored solution indicates the presence of
Cu2 ions, a brown gas is NO2, and a colorless
Step 1: Gas
Solution
Step 3: Limewater
Step 5: Product
16.1
The Nature of Oxidation-Reduction Reactions
561
For more practice with solving
problems, see Supplemental
Practice Problems,
Appendix B.
What role do the copper ions play in the redox reaction? Each copper
ion is reduced to uncharged copper metal when it accepts electrons from
the zinc metal. Because the copper ion is the agent that oxidizes zinc
metal to the zinc ion, Cu2 is called an oxidizing agent. An oxidizing
agent is the substance that gains electrons in a redox reaction. The oxidizing agent is the material that’s reduced. Because oxidation and reduction
go hand in hand, a reducing agent must be present. Zinc metal is the
agent that supplies electrons and reduces the copper ion to copper metal;
therefore, zinc is called the reducing agent. A reducing agent is the substance that loses electrons in a redox reaction. The reducing agent is the
material that’s oxidized. Figure 16.6 summarizes the roles of oxidizing
and reducing agents in redox reactions.
Loses
electrons
e e
Gains
electrons
e e
ˇ
ˇ
Reduced
oxidizing
agent
Reducing
agent
Oxidizing
agent
Oxidized
reducing
agent
Figure 16.6
Oxidizing and Reducing Agents
When electrons are transferred from one element to another, a combination
of an electron-gaining—or reduction—reaction and an electron-losing—or
oxidation—reaction takes place. This combination is called a redox reaction.
The element that is reduced oxidizes another element by attracting electrons
from it, so it is called an oxidizing agent. The element that is oxidized reduces
the first element by transferring electrons to it, so it is called a reducing agent.
SECTION REVIEW
Understanding Concepts
1. Name and define the two half-reactions that
make up a redox reaction.
2. Identify which reactant is reduced and which is
oxidized in each of the following reactions.
a) C5H12(l) 8O2(g) ˇ 5CO2(g) 6H2O(g)
b) 2Al(s) 3Cu2+(aq) ˇ 2Al3(aq) 3Cu(s)
c) 2Cr3(aq) 3Zn(s) ˇ 2Cr(s) 3Zn2(aq)
d) 2Au3(aq) 3Cd(s) ˇ 2Au(s) 3Cd2(aq)
3. What is the oxidizing agent when iron rusts?
What is the reducing agent?
Thinking Critically
base to form a salt and water. Determine the
oxidation number for each element. Is this a
redox reaction? Explain.
2KOH(aq) H2SO4(aq) ˇ
K2SO4(aq) 2H2O(l)
Applying Chemistry
5. Antioxidants Compounds that are easily oxidized can act as antioxidants to prevent other
compounds from being oxidized. Vitamins C and
E protect living cells from oxidative damage by
acting as antioxidants. Why does adding lemon
juice to fruit salad prevent browning of the fruit?
4. Applying Concepts The following equation
represents the reaction between an acid and a
562
Chapter 16
Oxidation-Reduction Reactions
chemistryca.com/self_check_quiz
SECTION
Applications of OxidationReduction Reactions
16.2
N
SECTION PREVIEW
atural redox reactions are going on around you every day, everywhere.
This is partly due to the abundance of oxygen, which acts as the oxidizing agent as it is reduced in some redox reactions. Other oxidizing
agents take part in different redox reactions, especially in environments
where not much oxygen gas is found. Near the vents of volcanoes, where sulfur compounds explode out from deep within Earth, enormous deposits of
solid yellow sulfur are found. The element sulfur acts both as an oxidizing
agent and as a reducing agent in the reaction that forms the sulfur deposits.
Can you tell which sulfur compound serves each function in this reaction?
2H2S(g) SO2(g) ˇ 3S(s) 2H2(g) O2(g)
Note that more than one element in a reaction can be oxidized or reduced.
The sulfur in hydrogen sulfide and the oxygen in sulfur dioxide both are oxidized. Sulfur in sulfur dioxide and hydrogen in hydrogen sulfide both are
reduced. Each reactant acts as both a reducing agent and an oxidizing agent.
Objectives
✓ Analyze common
redox processes to
identify the oxidizing
and reducing agents.
✓ Identify some redox
reactions that take
place in living cells.
Review Vocabulary
Oxidation: reaction
in which an element
loses electrons.
Say Cheese: Redox in Photography
Understanding natural redox reactions such as the one that occurs in
sulfur volcanoes has allowed chemists to develop many processes that
make use of oxidation and reduction reactions. Without them, photographs or steel wouldn’t exist, and stains would be much harder to
remove from clothing.
16.2
Applications of Oxidation-Reduction Reactions
563
Figure 16.7
Early Photos
In a daguerreotype, a redox
reaction between silver and
iodine fumes produced a layer
of light-sensitive silver iodide
on the surface of the polished
photographic plate. Exposure
to light caused decomposition
of the silver iodide into elemental silver, which was then
treated with the fumes of
heated mercury to form
bright amalgam areas. The
image of Paris shown here
was made by Daguerre himself.
photograph:
photos (GK) light
graphein (GK) to
write
Light is used to
record the image
of an object in a
photograph.
564
Chapter 16
Leonardo da Vinci described a primitive “camera” before 1519, in which
someone had to trace images focused on a glass plate inside a box. However, it wasn’t until 1838 that the French inventor L.J.M. Daguerre successfully fixed the images in a camera on highly polished, silver-plated copper
to make the first photographs. These early photographs were called
daguerreotypes in his honor, Figure 16.7.
Modern photographic film is made of a plastic backing covered with a
layer of gelatin, in which millions of grains of silver bromide are embedded. When light strikes a grain, silver and bromide ions are converted into
their elemental forms through a redox reaction. The equation for this
redox reaction is as follows.
2Ag 2Br ˇ 2Ag Br2
The reaction begins when the shutter on a camera is opened. Light
from the scene being photographed passes through the camera’s lens and
shutter and strikes the light-sensitive silver bromide on the film. The light
energy causes electrons to be ejected from a few of the bromide ions, oxidizing them to elemental bromine. The electrons are transferred to silver
ions, reducing them to metallic silver atoms. These grains are now activated. The developing chemicals continue the redox reaction by causing the
activated grains to be converted to metallic silver. In areas where the light
is brightest, more grains are activated, and after developing, they become
the darker areas. No silver atoms form in areas of the film that are not
struck by light, and that part of the film remains transparent. The exposed
film is then developed into a negative, during which time the remaining
AgBr and Br2 is washed away. Figure 16.8 describes the developing and
printing processes.
Oxidation-Reduction Reactions
Figure 16.8
Developing and Printing Pictures
Making photographic negatives by developing exposed film involves several steps. The process describes how
black-and-white pictures are made. For color photos, light-sensitive dyes are combined with the silver bromide in
layers on the film.
䊴 1. The exposed film is transferred to a canister, where it is developed using a solution of a reducing agent, or developer. The
organic compound hydroquinone is usually used for this purpose. The developer reduces all the silver ions to silver atoms in
any grain of silver bromide that was hit by light, but it does not
react with silver ions in grains that were not exposed to light.
Because metallic silver is dark and silver bromide is light, an
image having light and dark areas is produced.
䊴 2. After the film has been developed, a solution of a fixer containing thiosulfate ions is added. Thiosulfate ions react with unreduced silver ions to form a soluble complex, which is washed
away. This prevents unreduced silver ions from becoming
reduced and darkening slowly over time. The reaction follows.
AgBr(s) 2S2O32(aq) ˇ [Ag(S2O3)2]3(aq) Br(aq)
䊴 3. The fixed film is washed to remove any remaining developer or
fixer solution. The photographic negative is the reverse of the
image photographed; that is, light areas in the scene are dark
on the film, and vice versa.
䊴 4. When light is shone
through the negative onto
light-sensitive photographic
paper, a photographic print
is made. The print is positive; light and dark areas
are identical to those in
the scene.
16.2
Applications of Oxidation-Reduction Reactions
565
PHYSICS
CONNECTION
Solid Rocket Booster Engines
If you have ever built and launched a model rocket,
you probably noticed that the rocket engine was made of
a solid, highly combustible material packed into a cardboard tube. After ignition, the expansion and expulsion
of the gases produced enough downward force to launch
the lightweight rocket quickly into the air. Space shuttles
use a similar type of technology, but on a much larger
scale.
Engine systems The space shuttle has two different
engine systems. The three main engines attached directly
to the shuttle operate on liquid hydrogen and liquid
oxygen reservoirs carried in the large, centrally located
disposable fuel tank. The two smaller, reusable, strap-on
booster rockets on either side of the main fuel tank are
loaded with a solid fuel, which undergoes a powerful,
thrust-producing, oxidation-reduction reaction that
helps boost the shuttle into orbit.
Solid rocket fuel The solid rocket fuel is a mixture
containing 12 percent aluminum powder, 74 percent
ammonium perchlorate, and 12 percent polymer binder. Once ignited, the engine cannot
be extinguished. The extremely reactive ammonium perchlorate supplies oxygen to the
easily oxidized aluminum powder, providing a greatly exothermic and fast reaction. The
purpose of the polymer binder is to hold the
mixture together and to help it burn evenly.
The overall redox reaction is shown here.
Connecting to Chemistry
ˇ
24e–
8Al 3NH4ClO4 ˇ 4Al2O3 3NH4Cl
Shuttle forces Each solid rocket booster
weighs 591 000 kg at liftoff, produces 11.5
million N of force, and operates for about
two minutes into the flight. For comparison,
a 1000-kg car accelerating from 0 to 26.8 m/s
(60 mph) in 7 seconds would require a force
of only 3830 N. The tremendous release of
chemical energy and expansion of hot gases
due to the oxidation-reduction reaction
through the engine of the solid rocket booster produces the tremendous thrust needed to
get the 2 million-kg shuttle from 0 to almost
700 m/s (1500 mph) in just 132 seconds.
566
Chapter 16
Oxidation-Reduction Reactions
1. Applying Pow2. Acquiring Informadered aluminum is
tion Investigate the
used in another
lives and research of
greatly exothermic
Robert Goddard
reaction, the therand Werner Von
mite reaction,
Braun, who both
which is used for
experimented with
welding metals. The
rockets in the 1930s
reaction is as
and helped guide
shown.
the United States
into the space age.
2Al Fe2O3 ˇ
Write a short report
Al2O3 2Fe
about these men.
What role does the
powdered aluminum
play in this reaction?
Having a Blast: Redox in a Blast Furnace
Iron is seldom found in the elemental form needed to make steel.
Metallic iron must be separated and purified from iron ore—usually
hematite, Fe2O3. This process takes place in a blast furnace in a series of
redox reactions. The major reaction in which iron ore is reduced to iron
metal uses carbon monoxide gas as a reducing agent.
First, a blast of hot air causes coke, a form of carbon, to burn, producing CO2 and heat. Limestone, CaCO3, which is mixed with the iron ore in
the furnace, decomposes to form lime (CaO) and more carbon dioxide.
The carbon dioxide then oxidizes the coke in a redox reaction to form
carbon monoxide, which is used to reduce the iron ore to iron. The
process is outlined here and illustrated in Figure 16.9.
CaCO3(s) ˇ CaO(s) CO2(g)
The Stone, Bronze, and
Iron Ages are historical
periods named after
the most common
material that was used
for making tools during
each time. The Bronze
Age came before the
Iron Age because copper and tin, the elements that are melted
together to form the
alloy bronze, were both
widely available and
easily accessible metals.
Bronze is stronger than
either copper or tin
alone. The Iron Age
came later because iron
is harder to reduce to
elemental form. It
requires smelting at a
higher temperature
than bronze.
CO2(g) C(s) ˇ 2CO(g)
ˇ
6e –
2Fe (s) 3O (s) 3CO(g) ˇ 2Fe (l) 3CO2(g)
3+
2–
0
Redox in Bleaching Processes
Bleaches can be used to remove stains from clothing. Where do the
stains go? Bleach does not actually remove the chemicals in stains from
the fabric; it reacts with them to form colorless compounds. In chlorine
bleaches, an ionic chlorine compound in the bleach reacts with the compounds responsible for the stain. This ionic compound is sodium
hypochlorite (NaOCl). The hypochlorite ions oxidize the molecules that
cause dark stains.
OCl(aq) stain molecule(s) ˇ Cl(aq) oxidized stain molecule(s)
(colored)
(colorless)
Figure 16.9
䊱 Molten iron is drawn off at the bottom of the furnace. A combination
of by-products known as slag is also
removed at the bottom.
Limestone,
coke, and
iron ore
Blast Furnace
Iron ore (Fe2O3), coke
(C), and limestone
(CaCO3) are added at
the top of the furnace.
Hot air at about 900°C,
blasted into the bottom
of the furnace, burns the
coke in an exothermic
reaction. This reaction
causes temperatures in a
blast furnace to reach
about 2000°C. 䊳
16.2
Exhaust
gases
Compressed
air
Molten iron
Slag
Applications of Oxidation-Reduction Reactions
567
Testing for Alcohol by Redox
Organic alcohols react with orange dichromate ions, producing bluegreen chromium(III) ions. This reaction is used in a Breathalyzer test to
test for the presence of alcohol in a person’s breath. In this MiniLab,
you will use this reaction to test for the presence of alcohol in a number
of household hygiene, cosmetic, and cleaning products.
2
Procedure
CAUTION: Do not allow
dichromate reagent to come into
contact with skin. Wash with
1. Label five small test tubes with
large volumes of water if it does.
the names of the products to
4. Observe and record any color
be tested.
changes that occur within one
minute.
2. Place approximately 1 mL of each
Analysis
product in the
1. Which of the products that you
appropriate tube.
tested contain alcohol? Was the
presence of alcohol noted on
3. Wearing apron
the label of the products?
and goggles, add
three drops of
2. If the orange Cr2O72 ion
dichromate
reacts with alcohol to produce
reagent to each
the blue-green Cr3 ion, what
tube, and stir to
substance is the reducing agent
mix the solutions.
in the reaction?
Lab
See page 870 in
Appendix F for
Testing the Oxidation
Power of Bleach
Bleaches containing hypochlorite should be used carefully because
hypochlorite is a powerful oxidizing agent that can damage delicate fabrics. These bleaches usually have a warning label telling the user to test an
inconspicuous part of the fabric before using the product. In addition to
acting as a bleaching agent, hypochlorite ions are also used as disinfectants, as Figure 16.10 shows.
Figure 16.10
Hypochlorite as a Disinfectant
Hypochlorite is used in disinfectants to kill
bacteria in swimming pools and in drinking
water. In both cases, the hypochlorite ions
act as oxidizing agents. Bacteria are killed
when important compounds in them are
destroyed by oxidation. In this photo, the
amount of chlorine in the water is being
monitored. Chlorine reacts with the water to
form hypochlorite ions.
568
Chapter 16
Oxidation-Reduction Reactions
Breathalyzer Test
The alcohol in beverages, hair spray, and mouthwashes is
ethanol. Ethanol is a volatile liquid that evaporates rapidly at
room temperature. Because of this volatility, drinking an
alcoholic beverage results in a level of gaseous ethanol in the
breath that is proportional to the level of alcohol in the
bloodstream. About 50 percent of all automobile accidents
that result in a fatality are caused by intoxicated drivers. Law
officers can determine quickly whether a person is legally
intoxicated by using an instrument called a breath analyzer,
or Breathalyzer.
2. During a Breathalyzer test,
a person blows into the
mouthpiece of the bag.
1. A simple Breathalyzer device has an
inflatable plastic bag attached to a tube
containing an orange solution of potassium dichromate and sulfuric acid.
3. If alcohol vapors are present
in the person’s breath,
ethanol undergoes a redox
reaction with the dichromate.
As ethanol is oxidized, the
orange Cr6 ions are reduced
to blue-green Cr3 ions.
Thinking Critically
4. The exact color produced depends
on the amount of alcohol in the
breath. The color change that is produced during the test is compared
to standard color mixtures of the
two chromium ions to get an estimate of the blood alcohol level.
1. Suppose a person
used mouthwash
shortly before taking
a Breathalyzer test.
What might be the
result?
16.2
2. How would the color
produced in a
Breathalyzer test
change as the
ethanol content of
the blood increases?
Applications of Oxidation-Reduction Reactions
569
Figure 16.11
The Green Lady
The green color of the
Statue of Liberty in New
York Harbor is due to a
layer of patina, or protective coating, that covers the copper sheets
making up the statue.
The presence of the patina helps keep the statue
from corroding further
because oxygen cannot
get through the patina
to reach the copper layers underneath.
Corrosion of Metals
Did you know that the Statue of Liberty is made of copper sheets
attached to a steel skeleton? Why does it appear green rather than the
reddish-brown color of copper? When copper is exposed to humid air
that contains sulfur compounds, it undergoes a slow oxidation process.
Under these conditions, the copper metal atoms each lose two electrons to
produce Cu2 ions, which form the compounds CuSO4 3Cu(OH)2 and
Cu2(OH)2CO3. These compounds are responsible for the green coat or
patina found on the surface of copper objects that have been exposed to
air for long periods of time, Figure 16.11.
You have learned that iron is oxidized by oxygen in the air to form rust.
Aluminum is a more active metal than iron. As a result of its greater activity, aluminum is oxidized more quickly than iron. If this is true, why does
an aluminum can degrade much more slowly than a tin can, which is
made of iron-containing steel that is coated with a thin layer of tin? The
reason is that, like copper, aluminum is oxidized to form a compound
that coats the metal and protects it from further corrosion, as shown in
Figure 16.12. Aluminum reacts with oxygen to form aluminum oxide in a
redox reaction.
4Al(s) 3O2(g) ˇ 2Al2O3(s)
A coating of aluminum oxide is tough and does not flake off easily, as iron
oxide rust does. When rust flakes fall off a surface, additional metal is
exposed to air and becomes corroded.
Figure 16.12
Corrosion of Iron
and Aluminum
Because iron rust is
porous and flaky, it
does not form a
good protective
coating for itself. 䊳
A tin coating offers some protection
to the iron. However, if a hole or
crack develops in the thin tin coating, the underlying iron corrodes
rapidly. A tin-coated steel can will
degrade completely in about 100
years. The aluminum oxide coating
on an aluminum can is tough and
closely packed. It protects the
underlying aluminum from further
corrosion so that the can will take
about 400 years to degrade. 䊳
570
Chapter 16
Corrosive solution
Corrosive solution
Tin
Corrosive solution
Aluminum oxide
Iron metal
Iron metal
Aluminum metal
Steel can
Tin-coated steel can
Aluminum can
Oxidation-Reduction Reactions
Chemistry
Lightning-Produced Fertilizer
Did you know that one of the main nutrients
plants need is nitrogen? Although the air surrounding Earth is almost 80 percent nitrogen, the
nitrogen is in the form of N2 molecules, a form that
most plants and animals cannot use. Nitrogen from
the air is converted to a form that
plants can use by a process called
nitrogen fixation. Plants can best use
nitrogen when it is in the form of the
ammonium ion, NH4, where the
nitrogen has an oxidation number of
3, but they can also use the nitrate
ion, NO3, with nitrogen having an
oxidation number of 5.
Nitrogen fixation Nitrogen can be
fixed for plants in three ways: by lightning, by nitrogen-fixing bacteria living
in the roots of plants or in the soil,
and by commercial synthesis reactions
such as the Haber ammonia process.
Nitrogen is a fairly inert gas
because the triple bond of N2 is
strong and resists breaking. However,
the exceptionally high energy and temperatures of
lightning can easily break bonds and allow for
recombination of gases in the atmosphere.
Lightning-driven reactions In the process of
lightning-driven nitrogen fixation, nitrogen and
oxygen combine to form nitrogen monoxide.
Nitrogen monoxide then combines with more oxygen to form nitrogen dioxide. This nitrogen dioxide mixes with water in the air to form nitric acid
and more nitrogen monoxide, which is available to
continue the cycle.
N2 O2 ˇ 2NO
2NO O2 ˇ 2NO2
3NO2 H2O ˇ 2HNO3 NO
due to the dissolved nitric acid, HNO3, from nitrogen fixation. As the rain soaks into the soil, bacteria convert the nitrate ions into ammonium ions.
How does nature’s manufacturing of fixed
nitrogen compare with commercial production of
fixed nitrogen? Lightning may seem uncommon,
but it is estimated that there are approximately
10 000 lightning storms
every day over the surface of Earth. Stated
another way, lightning
strikes 100 times a second on the planet as a
whole. Approximately
10 billion kg of nitrogen
are fixed yearly in the
atmosphere. Biological
agents such as bacteria
fix about 100 billion kg
of nitrogen yearly, and
an amount equal to that
is fixed through the
manufacture of fertilizer
and other industrial
processes.
Exploring Further
1. Classifying Nitrogen fixation in the soil is
accomplished by bacteria living in the roots of
certain plants. Name some of these plants.
2. Applying In each of the three equations
shown, what is oxidized and what is reduced?
3. Acquiring Information The process by which
nitrogen is put back into the air is called denitrification. Find out what conditions are necessary for this process and what reaction occurs.
Fertilizer production The pH of rainwater is naturally slightly acidic, and some of this acidity is
To learn more about the nitrogen cycle, visit the
Chemistry Web site at chemistryca.com
16.2
Applications of Oxidation-Reduction Reactions
571
Silver Tarnish: A Redox Reaction
Imagine if, along with your usual chores of taking out the trash, washing dishes, feeding your pets, and taking care of your younger siblings,
you also had to polish the silver—as people did back in your greatgrandparents’ days. How would you find time for any fun? Fortunately,
other materials such as stainless steel have replaced most “silverware.”
Why do silver utensils have to be polished, but those made of stainless
steel or aluminum don’t? Silver becomes tarnished through a redox reaction that is a form of corrosion, as rusting is. Tarnish is formed on the
surface of a silver object when silver reacts with H2S in air. The product,
black silver sulfide, forms the coating of tarnish on the silver.
O2(g) 4Ag(s) 2H2S(g) ˇ 2Ag2S(s) 2H2O(l)
Many commercial silver polishes contain abrasives that help to remove
tarnish. Unfortunately, they also remove some of the silver. A more gentle
way to remove tarnish from the surface of a silver object involves another
redox reaction. In this reaction, aluminum foil scraps act as a reducing agent.
tarnish:
terne (OF) dull,
wan
The shiny surfaces
of many metal
objects lose luster
and become dull
and tarnished as
the metal atoms
undergo oxidation.
ˇ
6e –
2Al (s) 6Ag (s) 3S (s) 6H2O(l) ˇ 6Ag (s) 2Al (aq) 6OH (aq) 3H2S(g)
0
+
2–
0
3+
This reaction is essentially the reverse of the reaction that forms tarnish.
Here, silver ions in the Ag2S tarnish are reduced to silver atoms, while aluminum atoms in the foil are oxidized to aluminum ions. The tarnishremoving solution usually includes baking soda (sodium hydrogen carbonate) to help remove any aluminum oxide coating that forms and to
make the cleaning solution more conductive. Figure 16.13 shows how this
method of silver cleaning is done.
Figure 16.13
Removing Silver Tarnish
Even though corrosion is an unwanted redox reaction, removing the
tarnish makes use of another redox reaction. A nest of crumpled
aluminum foil scraps is made at the bottom of a large pot. The tarnished
silver object is added, making sure the silver is in contact with the foil
scraps. Baking soda is added, and the silver is covered with water. When
the pot is heated on a stove, the silver sulfide tarnish is reduced to silver
atoms, and the silver object becomes shiny and bright.
572
Chapter 16
Oxidation-Reduction Reactions
CHEMISTRY
&TECHNOLOGY
Forensic Blood Detection
The gas station at the corner was robbed, and the
cashier was shot. On television, police announce that
Suspect A has been taken into custody. They have
confiscated a jacket, allegedly worn by the suspect.
After preliminary examination by the police department, the jacket is sent to a forensic laboratory for
scientific investigation. One of the first tests a technician at the laboratory will carry out determines
whether or not there are blood stains on the jacket.
The Luminol Test
The technician may choose from several chemical tests for blood, all based on the fact that the
hemoglobin in blood catalyzes the oxidation of a
number of organic indicators to produce a colored product that emits light, or luminesces.
The technician on this case chooses the luminol test. Luminol has an organic double-ring
structure, shown below. In 1928, German
chemists first observed the blue-green luminescence when the compound was oxidized in alkaline solution. It was soon found that a number of
oxidizing agents, such as hydrogen peroxide,
bring about the luminescence. Later, workers
noted that the luminescence was greatly
enhanced by the presence of blood, which led to
its current use in forensic investigations.
NH2
O
last for a few minutes, the technician photographs the spots and their telltale light.
Ruling Out with Luminol
You may wonder if this relatively simple procedure will serve to convict Suspect A. Certainly
not. However, if the test had been negative, Suspect A might have been cleared from suspicion. A
negative result ensures that a stain is not blood.
But, because this is not the case with the stains
on the jacket, the luminol test is preliminary and
will be used with other tests.
The luminol test is especially useful because it
works well with both fresh and dried blood.
Luminol has one particularly useful feature. The
same stains can be
made luminescent
over and over again if
the spray is allowed
to dry and the stains
are resprayed.
A positive test
should not be taken
as absolute proof of
blood because luminol reacts with copper and cobalt ions, as well as
with the iron in hemoglobin. However, it reacts
much more strongly with hemoglobin. A large
number of forensic authorities believe that the
luminol test has value as a preliminary sorting
technique.
NH
Luminol
DISCUSSING THE TECHNOLOGY
NH
O
The technician carefully mixes an alkaline
solution of luminol with aqueous sodium peroxide and, in a darkened workplace, sprays the solution onto suspected spots on the jacket. Bingo!
An intense, blue-green chemiluminescence is
emitted from several spots. Because the glow will
1. Applying If a luminol test yields a positive reaction, what
is the next logical
step?
2. Hypothesizing
Why can it almost
16.2
never be assumed
that stains are
uncontaminated,
although stain evidence is important
in a criminal investigation?
Applications of Oxidation-Reduction Reactions
573
Figure 16.14
Chemiluminescence
䊴 When lightning is produced by an electrical discharge in the atmosphere, electrons in molecules
of O2 and N2 gases are excited to higher energy
levels. Energy from the electricity breaks the molecules into atoms. When the atoms recombine to
form molecules and the electrons return to lower
energy levels, light energy is released through
chemiluminescence.
When luminol is oxidized
and is observed in the dark,
an eerie blue-green glow is
produced through chemiluminescence. 䊳
Chemiluminescence: It’s Cool
Nitrous oxide is produced by a redox reaction between oxygen
and nitrogen during
lightning storms. It was
discovered and studied
in the late 1700s by
Joseph Priestley, who
found that inhaling it
resulted in unusual side
effects including laughing, singing, and fighting. For this reason, it
was called laughing
gas. Its anesthetic
properties were discovered by accident in
Connecticut in 1844 at
a public demonstration
given for amusement
when a man who
inhaled nitrous oxide
cut his leg badly in a
scuffle but felt no pain
until the gas wore off.
574
Chapter 16
Some redox reactions can release light energy at room temperature.
The production of this kind of cool light by a chemical reaction is called
chemiluminescence. The light from chemiluminescent reactions can be
used in emergency light sticks that work without an external energy
source. You may recall learning in Chapter 6 how these light sticks work.
Now you know that the reaction that takes place when the two solutions
in the light sticks are mixed involves an oxidation and a reduction.
Some chemiluminescent redox reactions occur naturally in the atmosphere as a result of lightning, Figure 16.14. Other chemiluminescent
reactions involve luminol, an organic compound that emits cool light
when it is oxidized. Luminol reactions are utilized by forensic chemists to
analyze evidence in crime investigations. They spray luminol onto a location where the presence of blood is suspected. If blood is present, the
iron(II) ions in the blood oxidize the luminol to form a chemiluminescent compound that glows in the dark. The iron is reduced by the luminol. Figure 16.14 shows the glow from the oxidized form of luminol.
Biochemical Redox Processes
How are bears able to stay warm enough to keep from freezing during
their winter hibernation? How do marathon runners get the energy to finish a race without stopping to eat? In both cases, fats stored in the body
are oxidized. Oxygen molecules from the air are reduced as they gain electrons to form water. In a series of redox reactions called respiration,
Oxidation-Reduction Reactions
Figure 16.15
Keeping Warm
Although it is common to think that only mammals keep
warm, in truth, all plants and animals maintain a temperature at which their enzymes function best. Plants keep from
freezing because heat is produced as a by-product of respiration and photosynthesis. One of the first plants to poke
through the snow in early spring is the heat-producing skunk
cabbage. The heat it releases allows it to get a head start on
other plants and also contributes to the unpleasant odor
that gives it its name.
energy is released. Figure 16.15 shows one effect of this heat in plants.
Respiration will be discussed in Chapter 19. Many other redox reactions
take place in living things. Electrons are transferred between molecules in
redox reactions during photosynthesis and in the reactions that fireflies
use to flash light signals to potential mates. You will study photosynthesis
in Chapter 20.
Some organisms can use the energy released during redox reactions to
convert chemical energy into light energy, a process called bioluminescence. You are probably familiar with the flashing lights given off by fireflies during courtship, but did you know that many different organisms—
including some fish, at least one type of mushroom, and a caterpillar
known as a glowworm—also are bioluminescent? Figure 16.16 shows bioluminescence in fireflies.
Now that you have learned what redox reactions are and have read
about some of the processes of which they are part, you can reexamine
the redox reaction that makes cut fruit turn brown. The color is due to
brown pigments that are formed by the oxidation of colorless compounds
normally present in the cells of the fruit. Oxygen in the air is the oxidizing
agent that reacts with the colorless compounds to produce the brown pigments. The oxygen is reduced when it accepts electrons from the pigments, so the pigments function as reducing agents. This combination of
oxidation and reduction
goes hand in hand in a
redox reaction because
electrons that are lost by
one element must be
gained by another.
bioluminescent:
bios (GK) life
lumen (L) light
escentis (L) beginning to be, have,
or do
A bioluminescent
substance undergoes a chemical
reaction in living
things in which
potential energy in
chemical bonds is
converted into
light energy.
Figure 16.16
Firefly Signals
Fireflies use flashes of light to attract mates. Light
energy is released during an enzyme-catalyzed
redox reaction. Luciferase is the name given to the
enzyme that speeds up the reaction in which the
organic molecule luciferin is oxidized.
16.2
Applications of Oxidation-Reduction Reactions
575
Figure 16.17
Antioxidants
Vitamin C owes its antioxidant properties
to the fact that it reacts so readily with
oxygen. When added to a food product,
oxygen reacts preferentially with vitamin C, thereby sparing the food
product from oxidation. Other
anti-oxidant food additives
include the synthetic compounds BHA and BHT and
the natural antioxidant,
vitamin E.
Myoglobin, found in
muscle tissue, is an
iron-containing protein
that stores oxygen.
Myoglobin in living
muscle tissue is bound
to oxygen and is a red
color. It becomes pale
purple after death
when the oxygen is
lost. Heating meat
results in oxidation of
the iron in myoglobin,
which then has the
brown color that tells
you the meat is cooked.
The skin of a fruit keeps oxygen out, which is why unbroken fruit does
not turn brown. Coating cut fruit with an antioxidant can prevent browning and keep a fruit salad looking fresh longer. The vitamin C in lemons is a
good antioxidant. If lemon juice is squirted onto cut banana or apple slices,
they will not brown as quickly because the vitamin C reacts with oxygen
more readily than do the fruit-browning compounds, Figure 16.17.
Connecting Ideas
For more practice with solving
problems, see Supplemental
Practice Problems,
Appendix B.
Most reactions involve electron transfer and thus are redox reactions. You
have learned to identify which element is reduced and which is oxidized
when you are given the equation for a redox reaction. You might wonder
why one element accepts electrons from another and whether you can
predict which element will be oxidized and which will be reduced. Learning
to make those predictions is the next step in your study of electron-transfer
processes in compounds and will help you understand how redox reactions
in batteries produce electricity.
SECTION REVIEW
Understanding Concepts
Thinking Critically
1. What role does the reducing agent hydroquinone play in the production of a photographic negative?
2. How is most of the iron that is used for making
steel purified from iron ores?
3. Why do aluminum cans degrade more slowly
than cans made of iron?
4. Applying Concepts Oxygen is required for the
production of light by fireflies. What role does
the oxygen play in the reaction?
576
Chapter 16
Oxidation-Reduction Reactions
Applying Chemistry
5. Bleaching Why can’t rust stains be removed
with bleach?
chemistryca.com/self_check_quiz
CHAPTER 16 ASSESSMENT
REVIEWING MAIN IDEAS
16.1 The Nature of Oxidation-Reduction
Reactions
■
■
■
Oxidation occurs when an atom or ion loses
one or more electrons and attains a more positive oxidation number. Reduction takes place
when an atom or ion gains electrons and
attains a more negative oxidation number.
Oxidation and reduction reactions always
occur together in a net process called a redox
reaction.
An oxidizing agent is the substance that gains
electrons and is reduced during a redox reaction. A reducing agent is the substance that
loses electrons and is oxidized during a redox
reaction.
■
■
■
16.2 Applications of Oxidation-Reduction Reactions
■
■
■
In photography, light triggers the reduction of
silver ions to silver metal on photographic film.
Bleach removes stains from clothing by oxidizing colored molecules to form colorless
molecules.
Metals such as copper and aluminum are
resistant to corrosion even though they are
easily oxidized because the products of their
reactions with oxygen form protective coatings on the surface of the metal.
Chemiluminescent reactions in emergency
light sticks, lightning, and the luminol reaction convert the energy of chemical bonds
into light energy.
Some organisms use redox reactions to produce light, which they use in communication.
This light production is called bioluminescence.
Cut fruits turn brown because compounds in
the fruit cells react with oxygen in a redox
reaction to produce brown pigments. Coating
the fruits with antioxidants can prevent this
browning.
Vocabulary
For each of the following terms, write a sentence that shows
your understanding of its meaning.
oxidation
oxidation-reduction
reaction
UNDERSTANDING CONCEPTS
1. What is the difference between an oxidizing
agent and a reducing agent?
2. Which of the changes indicated are oxidations
and which are reductions?
a) Cu becomes Cu2
b) Sn4 becomes Sn2
c) Cr3 becomes Cr6
d) Ag becomes Ag
3. Identify the oxidizing agent in each of the following reactions.
a) Cu2(aq) Mg(s) ˇ Cu(s) Mg2+(aq)
b) Fe2O3(s) 3CO(g) ˇ 2Fe(l) 3CO2(g)
chemistryca.com/vocabulary_puzzlemaker
oxidizing agent
reducing agent
reduction
4. What is the oxidizing agent in household
bleach?
5. Why does a photographic negative need to be
fixed?
6. In which direction do electrons move during a
redox reaction: from oxidizing agent to reducing agent or vice versa?
7. Why is aluminum metal used to remove tarnish from silver?
8. What chemical process do hibernating animals
use to stay warm?
9. Write the equation for the redox reaction that
occurs when a piece of iron metal is dipped in
a solution of copper(II) sulfate.
Chapter 16
Assessment
577
CHAPTER 16 ASSESSMENT
10. Identify the following as an oxidation reaction
or a reduction reaction.
Fe2 ˇ Fe3 e
Fe2
Fe3
e
APPLYING CONCEPTS
11. If galvanized nails, which have been coated
with zinc, are placed in a brown solution containing I2, the solution slowly turns colorless.
Adding a few drops of bleach to the colorless
solution results in a return of the brown color.
Explain what makes these changes occur.
12. List several ways in which a steel chain-link
fence could be treated to prevent corrosion.
13. When hydrogen peroxide is added to a colorless solution of potassium iodide, a red-brown
color appears. What substance is responsible
for the color?
14. Write the equation for a reaction that is not a
redox reaction. Are electrons transferred in this
reaction?
15. Indigo is one of the oldest known dyes. It has
been detected in cloth used to wrap mummies
that are more than 5000 years old. When cotton
jeans are dyed with indigo, they are dipped into
a yellow solution of indigo and sodium hydrosulfite, which is a good reducing agent. Within
minutes after being taken out of the solution,
the jeans turn blue. How can you explain this?
16. A shiny copper mirror can be formed on the
inside of a test tube in which the following
reaction takes place.
H2C=O(aq) Cu2(aq) 2OH(aq) ˇ
formaldehyde
Cu(s) HCOOH(aq) H2O(l)
formic acid
a) Identify the substance that is reduced during this reaction.
578
Chapter 16
Oxidation-Reduction Reactions
b) Identify the substance that is oxidized during this reaction.
c) What is the oxidizing agent in this reaction?
d) What is the reducing agent in this reaction?
17. Is oxygen a necessary reactant for an oxidation
reaction? Explain.
18. Sodium nitrite is often added to meat to inhibit the growth of microorganisms and to keep
the meat from spoiling. Under the acidic conditions in our stomachs, nitrites can be converted into potentially cancer-causing substances. Vitamin C can convert nitrite ions into
nitrogen monoxide gas and may help protect
us from the effects of these ions.
NO2(aq) ˇ NO(g)
a) Is the nitrite ion oxidized or reduced in this
reaction?
b) Does vitamin C act as an oxidizing agent or
a reducing agent?
19. Why is gold rather than copper used to coat
electrical connections in expensive electronic
equipment?
Everyday Chemistry
20. Just before World War I, a German chemist
named Fritz Haber developed a process for fixing atmospheric nitrogen into ammonia. The
ammonia produced this way can be converted
into ammonium nitrate, an important fertilizer and explosive.
3H2 N2 ˇ 2NH3
a) What element is oxidized during this reaction? What is reduced?
b) What is the oxidizing agent? What is the
reducing agent?
Physics Connection
21. By passing an electric current through water,
the water can be separated into its component
elements in the reverse of the reaction used to
power the main stage of the space shuttle.
2H2O(l) energy ˇ 2H2(g) O2(g)
a) Is this a redox reaction? If so, what element
is oxidized?
chemistryca.com/chapter_test
CHAPTER 16 ASSESSMENT
b) Where does the energy for this endothermic
reaction come from?
How It Works
22. If ethanol were less volatile, how might the
usefulness of a Breathalyzer test be affected?
Explain.
Chemistry and Technology
23. Why should a positive result from the luminol
test not be taken as proof of the presence of
blood?
THINKING CRITICALLY
Using a Table
24. The table below lists some of the most common compounds that are used as oxidizing
agents.
Common Oxidizing Agents
O2
K2Cr2O7
H2O2
HNO3
KMnO4
NaClO
Cl2
KClO3
a) Name each of the compounds in the table.
b) List at least one practical application, mentioned in this chapter or from a reference
book, of each of these oxidizing agents.
c) Make a similar table for common reducing
agents. Should any compounds be listed in
both tables?
Making Predictions
25. Hydrogen peroxide (H2O2) can be used to
restore white areas of paintings that have darkened from the reaction of lead paint pigments
with polluted air containing hydrogen sulfide
gas.
PbS 4H2O2 ˇ PbSO4 4H2O
(black)
(white)
Could hydrogen peroxide be used to remove
tarnish from silver objects? Would the reaction
have any undesirable effects?
Interpreting Data
26. ChemLab Write the balanced equation for the
reaction that caused the limewater to become
cloudy. Is this a redox reaction? Explain.
27. MiniLab 1 Why do you think corrosion seems
to occur mostly at the head and point of a nail?
Making Inferences
28. MiniLab 2 When a pile of orange ammonium
dichromate is ignited, it decomposes in an
exothermic reaction in which the green product and flames shoot upward like an erupting
volcano. (CAUTION: Do NOT perform this
reaction.)
(NH4)2Cr2O7(s) ˇ
Cr2O3(s) N2(g) 4H2O(g)
a) What is the reducing agent in this reaction?
The oxidizing agent?
b) How is this reaction similar to the Breathalyzer reaction?
CUMULATIVE REVIEW
29. Identify each of the following as a pure substance or a mixture. (Chapter 1)
a) petroleum
d) diamond
b) fruit juice
e) milk
c) smog
f) iron ore
30. List some characteristic properties of metals.
(Chapter 3)
31. Name each of the following ionic compounds.
(Chapter 5)
a) NaF
d) Na2Cr2O7
b) CaS
e) KCN
f) NH4Cl
c) Al(OH)3
32. How many grams of nitrogen are needed to
react completely with 346 g of hydrogen to
form ammonia by the Haber process?
(Chapter 12)
N2 3H2 ˇ 2NH3
Chapter 16
Assessment
579
CHAPTER 16 ASSESSMENT
33. Draw Lewis electron dot diagrams for each of
the following covalent molecules. (Chapter 9)
c) CH3CH3
a) CHCl3
b) CH3CH2OH
SKILL REVIEW
34. Designing an Experiment Do you think silver
will tarnish more quickly in clean air or in polluted air? Design an experiment to test your
hypothesis.
WRITING IN CHEMISTRY
35. Research the evidence that suggests that the
antioxidant properties of vitamin C may help
prevent cancer in people who take large doses
of this vitamin. Write a summary of your findings in which you propose how you would do
more tests to determine whether or not vitamin C has anticarcinogenic properties.
PROBLEM SOLVING
36. A flask filled with acid-washed steel wool is fitted with a long, thin glass tube in a rubber
stopper. When the flask is inverted so the tube
opening is in a beaker of colored water, the
water slowly begins to rise in the tube. Write a
summary of this experiment, as if you had performed it. Explain what makes the water rise.
Predict what portion of the flask will be filled
with water at the end of the experiment.
37. The patina coating on the Statue of Liberty has
preserved most of the copper metal in the statue. Some damage does occur wherever steel
rivets are in contact with copper and exposed
to water. Do library research to determine why
those sites are more susceptible to corrosion
than the rest of the statue. Write up your findings in a short report. Include a diagram or
make a poster showing the movement of electrons in the process.
580
Chapter 16
Oxidation-Reduction Reactions
38. Metallic lithium reacts vigorously with fluorine
gas to form lithium fluoride.
a) Write an equation for this process.
b) Is this an oxidation-reduction reaction?
c) If it is an oxidation-reduction, which element is oxidized? Which is reduced?
d) If 2.0 g of lithium are reacted with 0.1 L fluorine at STP, which reactant is limiting?
e) If 0.04 g of lithium fluoride is formed in
reaction in part d., what is the percent yield?
39. Identify the oxidizing reagent in each of the
following reactions.
a) C2H5OH(l) 3O2(g) → 2CO2(g) 3H2O(l)
b) CuO(s) H2(g) → Cu(s) H2O(l)
c) 2FeO(s) C(s) → 2Fe(s) CO2(g)
d) 2Fe2+(aq) Br2(l) → 2Fe3+(aq) 2 Br(aq)
40. When coal and other fossil fuels containing
sulfur are burned, sulfur is converted to sulfur
dioxide: S(s) O2(g) → SO2(g)
a) Is this an oxidation-reduction reaction?
b) If it is an oxidation-reduction, which element is oxidized? Which is reduced?
c) If 7.0 103 kg of fuels containing 3.5 percent sulfur are burned in a city on a given
day, how much SO2 will be emitted? Assume
that the sulfur reacts completely.
41. Sodium nitrite is formed when sodium nitrate
reacts with lead:
NaNO3(s) Pb(s) → NaNO2(s) PbO(s)
a) What is the oxidizing reagent in this reaction? What is the reducing reagent?
b) If 5.00 g of sodium nitrate is reacted with
an excess of lead, what mass of sodium
nitrite will form if the yield is 100 percent?
Standardized Test Practice
1. The term used to describe a chemical reaction
in which a substance losses electrons to another substance is
a) oxidation.
c) reduction.
b) redox.
d) corrosion.
Data for Elements in the Redox Reaction
Zn + HNO3 0 Zn(NO3)2 + NO2 + H2O
Element
Oxidation
Number
Complex ion of which
element is a part
Zn
0
none
2. The oxidation numbers of the elements in
CuSO4 are
a) Cu 2, S 6, O 2
b) Cu 3, S 5, O 2
c) Cu 2, S 2, O 1
d) Cu 2, S 0, O 2
Zn in Zn(NO3)2
2
none
H in HNO3
1
none
H in H2O
?
none
N in HNO3
?
NO3
4
none
3. For the reaction X + Y → XY, the element
that will be reduced is the one that is
a) more reactive.
c) more electronegative.
b) more massive.
d) more radioactive.
N in Zn(NO3)2
?
NO3
2
NO3
O in NO2
?
none
O in Zn(NO3)2
?
NO3
4. The reaction between sodium iodine and
chlorine is:
2NaI(aq) Cl2(aq) → 2NaCl(aq) I2(aq)
The oxidation state of Na remains unchanged
because
a) Na is a spectator ion.
b) Na cannot be reduced.
c) Na is an uncombined element.
d) Na is a monatomic ion.
O in H2O
2
none
5. The reaction between nickel and copper(II)
chloride is:
Ni(s) CuCl2(aq) → Cu(s) + NiCl2(aq)
The half reactions for this redox reaction are
a) Ni → Ni 2 2e; Cl2 → 2Cl 2e
b) Ni → Ni e; Cu+ + e → Cu
c) Ni → Ni 2 2e; Cu2 2e → Cu
d) Ni → Ni 2 2e; 2C 2e → Cu
Interpreting Tables: Use the table in the next column to answer questions 6–8.
6. Which of the following elements forms a
monatomic ion that is a spectator in the redox
reaction?
a) Zn
c) N
b) O
d) H
chemistryca.com/standardized_test
N in NO2
O in HNO3
7. The oxidation number of N in Zn(NO3)2 is
a) 3.
a) 1.
b) 5.
d) 6.
8. The element that is oxidized in this reaction is
a) Zn.
c) N.
b) O.
d) H.
9. Why are redox reactions so common in everyday situations?
a) Nitrogen is an abundant reducing agent.
b) Nitrogen is an abundant oxidizing agent.
c) Oxygen is an abundant reducing agent.
d) Oxygen is an abundant oxidizing agent.
Test Taking Tip
Write It Down!
Most tests ask you a large number of questions in a small amount of time. Write
down your work whenever possible. Write out the
half-reactions for a redox problem, and make sure
they add up. Do math on paper, not in your head.
Underline and reread important facts in passages
and diagrams—don’t try to memorize them.
Standardized Test Practice
581