Download Chapter 17: Electrochemistry

17 Electrochemistry
CHAPTER
Chapter Preview
Sections
17.1 Electrolysis: Chemistry from
Electricity
MiniLab 17.1 Electrolysis
17.2 Galvanic Cells: Electricity from
Chemistry
MiniLab 17.2 The Lemon with
Potential
ChemLab
Oxidation-Reduction and
Electrochemical Cells
582
Lightning in a Beaker?
Y
ou are familiar with this spectacular site of
electrical current. Did you know that electrical current also can flow through solutions
in a beaker? Silver plated tableware, chrome
trim on cars, and batteries are examples of
processes that take advantage of current flowing
in solutions.
Start-up Activities
What I Already Know
A Lemon Battery?
You can purchase a handy package of portable power at
any convenience store—a battery. You also can craft a
working battery from a lemon. How are these power
sources alike?
Safety Precautions
Use caution with electricity.
Materials
• lemon pieces
• zinc metal strip
• copper metal strip
• voltmeter with leads
Procedure
1. Insert the zinc and copper strips into the lemon,
about 2 cm apart from each other.
2. Attach the black lead from a voltmeter to the zinc
and the red lead to the copper. Read and record the
potential difference (voltage) from the voltmeter.
3. Remove one of the metals from the lemon and
observe what happens to the potential difference on
the voltmeter.
Analysis
What is the purpose of the zinc and copper metals? What
is the purpose of the lemon?
Review the following concepts
before studying this chapter.
Chapter 4: electrolytes
Chapter 7: role of electrons in
chemical bonding; atomic structure
Chapter 9: how metals conduct
electricity
Chapter 16: oxidation; reduction;
redox reactions
Reading Chemistry
As you look through the chapter, pay
attention to some of the graphics
and photos. Read the captions that
accompany them. Make a list of
some of the ways the chemistry of
electricity is used in the world of
technology or business. Keep this list
in mind while reading the chapter.
Preview this chapter’s content and
activities at chemistryca.com
583
SECTION
17.1
SECTION PREVIEW
Objectives
✓ Explain how a nonspontaneous redox
reaction can be driven
forward during
electrolysis.
✓ Relate the movement of charge
through an electrolytic
cell to the chemical
reactions that occur.
✓ Apply the principles
of electrolysis to its
applications such as
chemical synthesis,
refining, plating, and
cleaning.
Review Vocabulary
Reduction: reaction in
which an element
gains one or more
electrons.
New Vocabulary
electrical current
electrolysis
cathode
anode
electrolytic cell
cation
anion
584
Chapter 17
Electrolysis: Chemistry
from Electricity
S
Cl
H3N
ometimes, unexpected
results in scientific
Pt
research, such as GalH3N
Cl
vani’s frog, lead to important
discoveries. In 1964, a group of
researchers studying how electricity affects the growth of bacteria found that bacterial cells
stopped dividing if they were
subjected to an electric current.
This was an important finding
because agents that inhibit cell
division have the potential to
work as cancer treatments. Cancer is a disease in which body cells divide
uncontrollably. Upon doing more careful studies, the researchers found that
it wasn’t the electricity that was preventing the cells from dividing. Rather, a
compound was made when platinum from the electrodes used to provide the
current took part in a reaction caused by electrical energy. This compound
was named cisplatin, and today it is a major medicine used in chemotherapy.
Redox Revisited
You have learned how oxidation and reduction always occur simultaneously. Think about the chemistry of corrosion you studied in Chapter 16.
When iron metal reacts with oxygen, a redox reaction creates rust, iron
oxide. Electrons are always transferred when a redox reaction occurs. In
the rust reaction, electrons are transferred from the reducing agent, iron,
to the oxidizing agent, oxygen.
Suppose you could separate the oxidation and reduction parts of a
redox reaction and cause the electrons to flow through a wire. The flow of
electrons in a particular direction is called an electrical current. In other
words, you are using a redox reaction to produce an electrical current.
This is what occurs in a battery—one form of an electrochemical cell in
which chemical energy is converted to electrical energy. You can reverse
the process and use a current to cause a redox reaction to occur.
In this section, you will examine electrolysis, the process in which electrical energy is used to cause a non-spontaneous chemical reaction to
occur. In the second part of the chapter, you will learn about the process
that is the reverse of electrolysis—electrochemical reactions that can be
used to produce electricity.
Electrochemistry
Electrolysis
Figure 17.1
Not too long after Volta’s invention of the first electrochemical cell,
the British chemist Humphry Davy built a cell of his own and used it
to pass electricity through molten salts. An electrochemical cell consists
of two electrodes and a liquid electrolyte. One electrode, the cathode,
brings electrons to the chemically reacting ions or atoms in the liquid;
the other electrode, the anode, takes electrons away, Figure 17.1. The
electrons act as chemical reagents at the electrode surface. The liquid
electrolyte acts as the chemical reaction medium.
In Davy’s electrolysis of molten NaCl, sodium ions were reduced to
metallic sodium at the cathode. The oxidation of chloride ions to chlorine gas occurred at the second electrode, the anode. Each half-reaction
in the electrolysis of molten sodium chloride is shown.
2Cl⫺(l) ˇ Cl2(g) ⫹ 2e⫺
2Na⫹(l) ⫹ 2e⫺ ˇ 2Na(l)
When the equations for the two half-reactions are combined, the equation for the overall reaction can be written as follows.
2Na⫹(l) ⫹ 2Cl⫺(l) ˇ 2Na(l) ⫹ Cl2(g)
Davy discovered several elements in this way, beginning in 1807.
After releasing purified potassium metal from potassium hydroxide, it
took him only a year to produce magnesium, strontium, barium, and
calcium. Fewer than 30 elements had been isolated by 1800, but by
1850, more than 50 were known. Most of these new elements were isolated using electrolysis. Figure 17.2 shows the modern commercial
electrolysis of molten rock salt. Rock salt is sodium chloride, NaCl. In
this process, pure sodium metal and chlorine gas are produced.
Cl2 gas
Liquid Na
metal
Oxidation and Reduction
You can remember that
reduction always occurs at the
cathode and oxidation always
occurs at the anode by studying this diagram.
Cl2 output
Inlet for
NaCl
Molten
NaCl
Figure 17.2
Production of Sodium and Chlorine
A device called a Downs cell is used for the
electrolysis of molten sodium chloride. As an
electrical current is passed through the cell,
liquid sodium forms at the circular iron cathode. Because liquid sodium is less dense than
molten sodium chloride, the sodium floats to
the surface and is collected. Chlorine gas
forms at the carbon anode and is collected
at the top. An iron screen separates the two
electrodes.
Na outlet
Iron
cathode
Iron
screen
Carbon
anode
Electricity source
17.1
Electrolysis: Chemistry from Electricity
585
The Electrolysis Process
Electrolysis takes place in a type of electrochemical cell called an electrolytic cell, in which a source of electricity, such as a battery, is added to
an external circuit connecting the electrodes. The electrolysis process
occurs when the electrons are transferred between the electronic conductors—the metal electrodes—and the ions or atoms at the electrode surfaces. In the liquid, the charge is conducted by ions such as the Na⫹ and
Cl⫺ in the molten rock salt. Of course, ions contain electrons too, but
these electrons are held tightly to individual ions. Ions can conduct current through a liquid only when they move through the liquid. This type
of conduction is called electrolytic conduction. A positive ion is known as
a cation. A negative ion is called an anion. Notice in Figure 17.3 that the
cations, Na⫹ ions, move toward the cathode. The negative anions, Cl⫺
ions, move toward the anode.
What happens when the moving ions reach an electrode surface? If the
electrodes are inert, which means that they don’t react chemically with the
ions in the solution, then only electron transfer will take place at the electrodes. Electrons are being pumped from the battery toward the cathode,
where reduction will occur. At the cathode, the ion that reacts is the one
that most readily reacts with electrons. In Figure 17.3, Na⫹ and Cl⫺ are
both present in the liquid near the surface of the cathode. Na⫹ accepts
electrons more readily, so each Na⫹ cation gains one electron, being
reduced to the metal.
At the anode, electrons are transferred from the ion that most easily
gives them up to the anode. In this case, Cl⫺ holds onto its electrons more
loosely, so each Cl⫺ anion loses an electron and is oxidized to a chlorine
atom. Chlorine atoms then combine to form Cl2 molecules. The electrons
released by the chloride ions flow through the external circuit to the battery and are recycled to the cathode, where they continue the reduction
reaction. Because only as many electrons are available at the cathode as
are removed at the anode, the reduction process at the cathode must
electrolysis:
electricus (L)
involving
electricity
lytikos (GK) to
split
Compounds are
split into their
component elements during electrolysis.
Battery
e–
Figure 17.3
Electrolytic Cell
Electrolysis, the splitting of compounds by electricity, occurs when two electrodes, an anode
and a cathode, are inserted into a liquid electrolyte such as molten sodium chloride and
connected to a source of electrical energy such
as a battery. When electrical current flows into
the electrolytic cell, chemical reactions occur.
Anions and cations conduct the current by
moving freely through the liquid. In the external circuit, electrons move out of the anode,
through the battery, and into the cathode.
e–
Anode
Cathode
Cation
Na+
Cl–
Anion
Molten NaCl
Oxidation
Reduction
–
–
+
–
Cl2(g) + 2e
2Cl
2Na + 2e
2Na(l)
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Chapter 17
Electrochemistry
Electrolysis
Electrolysis is the process in which an electric current is used to produce a chemical change. In this MiniLab, you will use the current from
a 9-V battery to produce chemical changes in a simple electrolytic cell.
Procedure
1. Pour about 200 mL of 0.5M
copper(II) sulfate into a 250mL beaker.
2. Use wires with alligator clips to
attach two 4-inch pencil leads
(actually graphite, not lead) to
the terminals of a 9-V battery.
3. Put the pencil leads into the
copper(II) sulfate solution,
keeping them as far from each
other as possible.
4. Observe the reactions that
occur at the two pencil leads
for five minutes. Record your
observations, including the
polarities (⫹ and ⫺) of the
leads.
Analysis
1. Describe the reaction that
occurs at the positive electrode,
the anode, of the electrolytic
cell. Write the equation for the
oxidation reaction.
2. Describe the reaction that
occurs at the negative electrode, the cathode. Write the
equation for the reduction
reaction.
3. Describe how you might use an
electrolytic cell to silver plate
an iron spoon.
1
always occur together with the oxidation process at the anode. Charge
transfer at the two electrodes must exactly balance because, just as in the
redox reaction, the liquid and its contents must always remain electrically
neutral. Therefore, in an electrolytic cell, the overall result of the two electrolysis processes is to carry out a balanced redox reaction, even though
the two half-reactions take place at different locations.
Producing Chemicals by Electrolysis
In the electrolysis of molten sodium chloride, a redox reaction taking
place in an electrolytic cell can be used to generate chemicals that are
important commercially. Because electrolysis consumes large amounts of
energy when it is carried out on a commercial scale, many companies that
use this process have made their homes in locations where electric power
is inexpensive. The abundant hydroelectric power available from Niagara
Falls has made that area of New York State a prime location for companies that use electrolysis. One commercial electrolytic process that is carried out in this area is the electrolysis of rock salt solutions to produce
chlorine, hydrogen, and sodium hydroxide. The overall change taking
place in this process is a redox as well as a substitution reaction.
2NaCl(aq) ⫹ 2H2O(l) ˇ Cl2(g) ⫹ H2(g) ⫹ 2NaOH(aq)
How is electrolysis of a sodium chloride solution different from electrolysis of molten sodium chloride? In molten NaCl, the only ions present
are Na⫹ and Cl⫺. What ions are present in an aqueous solution of rock
salt? Recall that water dissociates slightly to form H⫹ and OH⫺ ions.
17.1
Electrolysis: Chemistry from Electricity
587
Therefore, a rock salt solution contains Na⫹, H⫹, Cl⫺, and OH⫺ ions. At
the anode, the chloride ions lose electrons more easily than the other ions
present, just as they did in the electrolysis of molten rock salt. This oxidation forms chlorine gas that can be used for making PVC plastic and
other consumer products. At the cathode, the reaction that occurs in
molten salt doesn’t occur. Instead, H⫹ ions are easier to reduce than Na⫹,
Cl⫺, or OH⫺ ions, so hydrogen ions pick up electrons from the cathode
and are reduced to form hydrogen gas. Hydrogen is used in industrial
processes such as the catalytic hydrogenation of vegetable oils to form
margarine. The Na⫹ and OH⫺ ions are left dissolved in water after the
electrolysis process has removed H⫹ and Cl⫺ ions. This is a solution of the
base sodium hydroxide, NaOH. Sodium hydroxide is an important industrial and household chemical.
As you have seen, electrochemical transformation of a simple salt solution has produced three valuable products. Each can be sold to pay for the
electrical energy that must be invested to make them, with a little left over
for a profit to the manufacturer.
Applications of Electrolysis
Electrolysis has numerous useful applications in addition to the generation of chemical substances. The process can also be used to purify metals
from ores, coat surfaces with metal, and purify contaminated water. The
applications range from the world of art to the world of heavy industry.
+
Figure 17.4
Hall-Héroult Method of Producing
Aluminum
A Hall-Héroult electrolytic cell is used
to produce aluminum metal. It is made
of a steel shell lined with carbon that
forms the cathode. Anodes of carbon
hang down into the solution of aluminum oxide dissolved in cryolite. 䊳
Molten aluminum is drawn
off from the bottom of the
electrolytic cell, where it
accumulates during the
electrolysis process. 䊳
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Chapter 17
Electrochemistry
Carbon anode
–
Carbon-lining
cathode
Electrolyte
Al2O3 in
Na3AlF6(l)
Molten Al
Al outlet
Figure 17.5
Recycling Aluminum
䊴 Recycling of aluminum provides a relatively cheap source of this
metal. Currently, 60 percent of aluminum beverage cans are
recycled in the United States.
Refining Ores
Just as sodium can be produced from melted NaCl by electrolysis, many
metals are separated from their ores, or refined, using electrolysis. Today,
the metal produced in largest quantity by electrolysis is aluminum from
aluminum oxide. First, bauxite ore is heated, driving off the water and
leaving aluminum oxide, Al2O3. Pure aluminum oxide melts at about
2000°C, so cryolite, Na3AlF6, is added to lower the melting point to about
1000°C. The molten salt solution is placed in a large electrolytic cell lined
with carbon, which acts as the cathode during electrolysis. Large carbon
anodes are dipped into the molten salt to complete the cell, as shown in
Figure 17.4.
The electron transfers cause complex chemical reactions during electrolysis of molten Al2O3 / Na3AlF6, but the net reactions are simple. At the
cathode, electrons reduce aluminum ions to aluminum metal, which is
molten at this temperature. At the anodes, oxide ions lose electrons to
form oxygen. The oxygen then combines with the carbon anodes to produce carbon dioxide. The carbon anodes must be replaced periodically
because they are gradually used up. The overall redox reaction follows.
2Al2O3(s) ⫹ 3C(s) ˇ 4Al(l) ⫹ 3CO2(g)
This electrolytic method for producing aluminum is called the HallHéroult process because it was developed simultaneously by Charles
Martin Hall of the United States and Paul Héroult of France in 1886.
Aluminum was so rare before this process was developed that it was more
expensive than silver or gold. Today, about 10 million metric tons of aluminum are produced per year worldwide using this process.
The Hall-Héroult process is expensive and consumes large amounts of
electrical energy. Recycling aluminum metal prevents some of the expense
of producing new aluminum by electrolysis. A lot of waste aluminum is
available from discarded aluminum containers, and the energy invested in
the original refining of aluminum can be saved if metallic aluminum can
be melted down, Figure 17.5. It takes only about seven percent as much
energy to make new aluminum cans from old ones as it does to make new
cans from aluminum ore.
17.1
䊱 Scrap aluminum is melted
in a furnace and then
re-formed into other
products.
Electrolysis: Chemistry from Electricity
589
CHEMISTRY
&TECHNOLOGY
Copper Ore to Wire
It would be hard to imagine life without many of
the common metals that are used today. Copper, for
example, is used in many of the pots and pans in
your kitchen, the cooling coils in the air-conditioning
system, and some of the pennies you carry in your
pocket. More importantly, copper is the metal of
choice for most of the electrical wiring in your appliances, homes, and cars because of its good conductivity and low cost. How does copper get from rocks to
the finished wire?
Copper can be found in the ground as the free
metal and was used about 5000 B.C. In Roman times,
much of the copper was obtained from the island of
Cyprus, whose name means “copper.” Coppersmiths
soon learned that copper could be shaped by exposing
it to a slow, softening heat in a process called annealing. Annealing soon led to the development of the
smelting process. By 3000 B.C., smiths were adept at
the metallurgical processes of hammering, annealing,
oxidation and reduction, melting, alloying, and
removing impurities. Unfortunately, these processes
could produce only small quantities of copper. Largescale production was not mastered until modern furnaces and rolling techniques were developed.
1. Mining
Copper can be found as the free metal in many
parts of the world. However, most copper is
found as CuFeS2 (chalcopyrite), Cu2S (chalcocite), and CuS (covellite). The ores are typically
surface mined and ground into powders.
2. Ore Enrichment
Because each ore contains only from one percent to ten percent copper, the copper ore must
be concentrated by the flotation process. A frothing agent such as pine oil is mixed with the powdered copper ore, and air is blown in to froth the
mixture. Because copper ores are hydrophobic
(not wetted by water), the copper and iron sulfides cling to the oil and float to the top, where
they can be continuously removed.
Copper mine, Arizona
Ore enrichment
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Chapter 17
Electrochemistry
SO2. The blistered surface of the copper is due to
the escaping gas. The blister copper can be drawn
off the bottom and cast into large blocks.
5. Purification
Roasting
3. Roasting
The preparation of copper metal involves
roasting the ore with oxygen to convert the
metallic sulfides to metallic oxides. Usually, both
iron and copper are present in the mix.
2Cu2S(s) ⫹ 3O2(g) ˇ 2Cu 2O(s) ⫹ 2SO2(g)
2FeS(s) ⫹ 3O2(g) ˇ 2FeO(s) ⫹ 2SO2(g)
4. Smelting
The copper and iron oxides are smelted by
mixing and heating them with SiO2, air, and limestone. The result is dense blister copper, lightweight iron(II) calcium silicate slag, and gaseous
The copper can be purified by electrolysis to
99.95 percent purity. The large blocks of blister
copper are used as anodes suspended in a solution
of aqueous copper(II) sulfate. Pure copper is used
as the cathode. During electrolysis, copper is oxidized at the anode, moves through the solution as
Cu2⫹ ions, and is deposited on the cathode. Waste
products left
after the dissolution of the
anode produce a sludge
on the bottom
of the electrolysis vessel. The
sludge, which
is often rich in
silver and
gold, can be
recovered for
profit.
Electrolysis
–
+
Smelting
Slag
removal
Air
in
Iron(II) calcium
silicate slag
Pure
copper
cathode
CuSO4
Blister
copper
anode
Purification
Molten
copper
Copper
removal
Heat
17.1
Electrolysis: Chemistry from Electricity
591
6. Wire Bar Production
The electrolytic copper is
cast into wire bars ranging
from 60 to 227 kg. The wire
bars are heated to 700-850°C
and rolled without further
reheating into rods approximately 1 cm in diameter.
Wire bar production
7. Drawing Wire
8. Coatings
The 1-cm rods of copper are drawn through
successively smaller dies until the desired size of
wire is reached. The dies must be made of exceptionally hard materials because of the tremendous
amount of wear from drawing the wire. The dies
typically are made of tungsten carbide or diamond.
The finished wire may be coated with plastic,
enamel, or another metal to help protect it from
moisture and oxidation, or it may be left bare.
Drawing wire
DISCUSSING THE TECHNOLOGY
1. Thinking Critically Why is it important that
the slag formed during smelting be lightweight?
2. Acquiring Information Copper can be made
into many different forms and alloys. Research
and list several of the forms and uses for each.
592
Chapter 17
Electrochemistry
3. Hypothesizing How might the gold and silver
be removed from the sludge formed beneath
the anode?
Electroplating
You have learned that many metals can be protected from corrosion by
plating them with other metals. Zinc coatings are often used to keep iron
from rusting. Metal garbage cans are galvanized by dipping them into
molten zinc. This process produces an uneven, lumpy coating both inside
and outside of the can. That’s OK for a garbage can, but the lumpy surface wouldn’t look good under the snazzy red paint job on a new sports
car. That’s why automobile manufacturers electroplate zinc onto the steel
used for car bodies. This process involves dragging a sheet of steel across
the surface of an electrolyte in an electrolytic cell. The process produces a
thin (cost-saving), uniform (smooth and clean) coating of zinc on only
one side of the sheet of steel, saving half the coating cost. Because only
one side of the car body is exposed to the corrosive effects of water and
salt, it is not worth the cost of coating the inside. Figure 17.6 shows other
uses of electroplating.
In zinc electroplating, zinc ions are reduced to zinc atoms at the surface of the metal object to be coated, which becomes the cathode in an
electrolytic cell. At the anode, which is made of zinc, atoms of zinc are
oxidized to ions. The electrolyte solution contains dissolved zinc salt.
Reduction of silver
ions onto cheaper
metals forms silverplate. The object to be
plated is made the
cathode. At the pure
silver anode, oxidation
of silver metal to silver
ions replaces the silver
ions removed from
the solution by plating
at the cathode. 䊳
e–
Lab
See page 871 in
Appendix F for
Removing Electroplating
Figure 17.6
Electroplating
Chromium is often electroplated onto a softer metal to
improve its hardness, stability,
and appearance. Chrome
bumpers and trim can be found
on many vintage cars. 䊲
e–
–
+
Battery
–
NO3
Anode
Ag+
Ag + + e –
Ag
Cathode
Ag
Ag + + e –
䊴 In the early 1980s, because of high inflation of
the U.S. currency and a worldwide shortage of
copper, the cost of producing a copper penny
became nearly equal to the value of the copper
metal itself. The U.S. Mint was instructed to
begin making pennies from a cheaper metal,
such as zinc. Knowing that the American public
would not accept dusty-gray zinc pennies, the
Mint began to produce pennies by electroplating
zinc disks with a copper coat. Many people think
that pennies are still made of pure copper.
17.1
Electrolysis: Chemistry from Electricity
593
Chemistry
Manufacturing a Hit CD
The sound from your new compact disc (CD) is
so clear and crisp that it seems as if the musicians
are in the same room. The clarity of the sound is
possible thanks to the chemical process of metallic
depositing and electroplating used in the manufacturing of CDs.
Data pickup A CD is a collection of binary data
in a long, continuous spiral that runs from the
inside edge to the outside. Nothing touches the
information stored on the CD except laser light,
which is reflected from the CD and read by the CD
player computer as a binary signal. This signal is
then transformed by a tiny computer into audio
signals, which are amplified to produce the sound
heard from your speakers.
Master copy production Your CDs are stamped
from a master copy called a stamper master. But
how is this stamper made? The musical data are
first transferred onto a glass disc using a highpowered laser that etches small pits in the glass as
shown here. This glass master copy then contains
all of the musical information in binary data form.
The master is coated with a dilute solution of silver diammine complex—[Ag(NH3)2]⫹—followed
Electrodeposition The silver coating forms the
surface onto which a thin layer of nickel is electrodeposited to make a nickel-coated disc called the
mother disc.
Ni2⫹(aq) ⫹ 2e⫺ ˇ Ni(s)
A second coating of nickel is then electrodeposited
onto the mother disc. This nickel layer, the stamper master, is peeled off and used to stamp the
data impression onto melted polycarbonate plastic
discs. The polycarbonate disc now has all of the
pits found on the original glass master disc etched
by the laser. Because polycarbonate is clear, it is
vacuum coated with a thin aluminum film to produce the reflective layer required by the laser. This
delicate aluminum layer is covered with a protective layer of polycarbonate to prevent aluminum
oxidation and marring of the data surface. The
back of the CD can now be covered with information in the form of art and lettering.
Stamping characteristics Each nickel stamper
master can make about 20 000 copies before it
wears out. At this point, the nickel stamper can be
recycled into an aqueous nickel solution and used
to make more nickel stampers. A hit CD may go
through as many as 50 nickel stampers plated onto
the mother disc. CD-ROMs for your computer are
made in the same way.
Exploring Further
Etched pits, magnified 1200⫻
by a solution of formaldehyde that acts as a reducing agent for the silver. The result is a redox reaction
that deposits a thin silver mirror that plates the
etched disc. This mirrored master disc is produced
in the following reaction.
2Ag⫹ ⫹ HCHO ⫹ H2O ˇ 2Ag ⫹ HCOOH ⫹ 2H⫹
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Chapter 17
Electrochemistry
1. Hypothesizing Why do you think nickel is
used for stamping the polycarbonate discs?
2. Acquiring Information Polycarbonate is the
base material used for CDs. Research its properties and some of its other major uses.
To find out more about the technology behind the
production of compact discs, visit the Chemistry
Web site at chemistryca.com
The thickness of the zinc metal coating can be controlled exactly by controlling the total charge (number of electrons) used to plate the object.
Only that portion or side of the object that is immersed in the cell electrolyte receives a zinc coat.
Because coatings adhere best to a chemically clean foundation, objects
to be coated are usually degreased, cleaned with soap, and then treated
with a corrosive fluid to remove any dirt on their surface before electroplating. Then, the object is immersed in an electrolyte containing the salt
of the metal to be deposited. Because the object acts as the cathode of the
electrolytic cell, it must be a conductor. Metal objects are electroplated
most often because metals usually are excellent conductors. The anode is
made of the same metal that is being plated so that it will replenish the
metal ions in the electrolyte that are removed as the plating proceeds. The
net effect is that when an electric current is passed through the electroplating bath, metal is transferred from the anode and distributed over the
cathode. Eventually, the entire object becomes coated with a thin film of
the desired metal.
corrosion:
com (L)
thoroughly
rodere (L) to
gnaw
Metals wear away
gradually during
corrosion, as if
they were being
eaten.
Electrolytic Cleaning
Electrolysis can be used to clean objects by pulling ionic dirt away from
them. The process has been used to restore some of the many metal artifacts taken from the shipwrecked cruise ship Titanic, which sank in the
northern Atlantic Ocean in 1912, Figure 17.7. Coatings of salts containing
chloride ions, which came from the seawater, were removed by electrolysis. The electrolysis cell for this cleaning process includes a cathode that is
the object itself, a stainless steel anode, and an alkaline electrolyte. When
an electric current is run through the cell, the chloride ions are drawn
out. Hydrogen gas forms and bubbles out, helping to loosen corrosion
Figure 17.7
products. Among the objects that have been recovered are a porthole, a
When That Great Ship
chandelier, and buttons from the uniforms of crew members.
Went Down
Chemistry played a key role in
restoring many items taken
from the wreck of the Titanic.
Electrolysis was used to clean
and stabilize many metal artifacts, and electrophoresis was
used to remove corrosion from
bank notes, leather, and
objects such as these casserole
dishes. Chemicals that attract
and hold metal atoms or ions
were used to remove iron
stains from delicate objects
made of organic materials such
as newspapers, textiles, and
letters. The study of objects
from the ship may help scientists compile information for
long-term storage and containment under seawater.
17.1
Electrolysis: Chemistry from Electricity
595
Meet Harvey Morser,
Metal Plater
A seasonal highlight in Reno, Nevada, is a celebration
called Hot August Nights. On display are thousands of
beautifully restored antique cars. “Muscle cars”—like a
1934 Ford sedan with a $12 000 custom paint job and
immaculately plated chrome—draw admiring crowds.
Harvey Morser watches with pride as the cars sporting his
shop’s work on grilles and door handles parade by.
On the Job
Mr. Morser, will you tell us what you do
on the job?
Although I’m now the owner of Western
Metal Finishing, I still go out on the shop
floor and do plating. All iron metal needs
some type of protective coating to keep it from
rusting. There are different types of applications
that you can use. Cadmium plating is probably the
best, but it’s very toxic. Hardchrome plating also
releases fumes into the atmosphere. So, because of
environmental concerns, electroless nickel is the
favored process. It’s a chemically applied nickel
plating done without electricity. In my opinion,
this process is twice as effective because it goes on
easily and consistently, unlike hard plating, which
you have to apply and then grind back down.
Chemists have developed an electroless nickel that
has the same Rockwell factors (the hardness factors) as hardchroming. I know that a major heavyequipment manufacturer converted from hardchrome over to electroless nickel and saved something like $3 million the first year.
596
Chapter 17
Electrochemistry
What kinds of metal products do you
plate?
One of my major accounts is for the metal
straps on hearing protectors that the Navy
uses on aircraft carriers. At this plant, we
also plate things like computer chassis and covers
for stereo systems.
Why does the electroless nickel process
produce a more even coating?
Visualize a flat, square plate. When you put
a hook in one edge of it and hang it in the
tank to hardplate it, the electrical current
will reach the corners first, then travel down the
side edges, and finally spread into the center of the
plate. The plating will go on in the same way,
building up probably twice as fast on the edges as
in the center. So the edges might have eight tenthousandths of an inch of plating, whereas the
center might only have four ten-thousandths. With
electroless nickel, a metal part will plate perfectly
evenly because it’s put on chemically. Electroless
Personal Insights
If someone came to you just out of high
school, would you give him or her a
chance on the job?
In a heartbeat. I believe people need to get
their schooling, but they also need to have
common sense. Every employee here gets
on-the-job training, just like I did. What I look for
in a prospective employee is honesty and dependability, plus an ability and desire to learn.
Is this a stressful business to be in?
nickel has the highest corrosion resistance next to
cadmium plating, so these plated parts resist corrosion as well as being uniform.
Early Influences
What training did you have in metal
plating?
All my training came on the job. I started
out polishing the metal prior to plating.
That’s a tough and dirty job, but an important one. Plating duplicates a surface, so it has to be
polished like a mirror. Otherwise, plating will magnify even a tiny pit. In those early days, I had to
work a side job in a bowling alley as a pin chaser,
unsticking the balls and unjamming the pins.
Along the way, I picked up carpentry, welding, and
electrical skills, which have come in very handy in
doing the maintenance at the metal plant.
Absolutely! I carry a pager all the time and
even take a cellular phone out on the lake
when I go fishing. There’s always the possibility of an industrial accident. Earthquakes don’t
announce that they are on their way. If I see the
numbers 1-5 on my pager, I know it’s all clear. Seeing five zeros is what I dread.
What appeals to you about the plating
business?
I was always turned on by the fact that I
could take something that looked terrible,
like an old car part, and make it look gorgeous. I want even the most modest plating job to
look good, even when it’s on a part that probably
won’t be visible after it’s installed.
How did you work your way up to owning the plant?
CAREER
It all boiled down to learning quickly, not
whining, and working hard. Plating is
tough and heavy work. The heat on the
lines is terrible, with the humidity and steam off
the tanks. All the tanks are running about 160°, so
in the summertime it will be 104° on the line. The
joke here is that we don’t charge extra for the steam
bath. Nine years ago, I became the owner of the
business. To me, that’s the Great American Dream.
CONNECTION
These jobs also involve working with metals.
Metallurgical Technician Two-year training
program
Mining Engineer Bachelor’s degree in engineering
Scrap Metal Processing Worker On-the-job
training after high school
17.1
Electrolysis: Chemistry from Electricity
597
Figure 17.8
Electrophoresis
Electrophoresis is a valuable laboratory tool
used to separate and identify large charged
particles such as DNA and proteins. Negatively charged particles move toward the
anode. Positively charged particles move
toward the cathode.
Electrophoresis is another electrochemical process that was used to
restore some of the ceramic and organic artifacts from the Titanic. Electrophoresis involves placing an artifact in an electrolyte solution between
positive and negative electrodes and applying a current. The current
breaks up salts, dirt, and other particles as their charged components
migrate to the electrodes. Electrophoresis is also used in laboratories to
separate and identify large molecules, Figure 17.8.
Electrolysis of Toxic Wastes
For more practice with solving
problems, see Supplemental
Practice Problems,
Appendix B.
The plating baths used in the various applications of electrolysis often
contain toxic materials or produce toxic by-products. After bath solutions
have been used for a period of time, they must be changed and the toxic
contents disposed of in a safe manner. Remarkably, electrolysis offers one
of the safest and most thorough means of cleaning up toxic metalcontaining wastes. When the bath solution is subjected to electrolysis, the
toxic metal ions are reduced to free metal at the cathode. The metal can
then be recycled or disposed of safely.
SECTION REVIEW
Understanding Concepts
Applying Chemistry
1. Draw and label the parts of an electrolytic cell.
2. Of what value is electroplating? Why is the
process used?
3. How is electrolysis used for cleaning objects?
5. Magnesium from the Sea Magnesium in seawater is found mostly as Mg(OH)2, which can
be converted to MgCl2 by reacting it with HCl.
Magnesium metal can then be purified by electrolysis of molten MgCl2.
a) What reaction takes place at the cathode
during electrolysis?
b) What reaction takes place at the anode during electrolysis?
c) Write an equation for the net reaction that
occurs.
Thinking Critically
4. Applying Concepts What effect would electroplating steel jewelry with gold have on the rate
of corrosion of the jewelry?
598
Chapter 17
Electrochemistry
chemistryca.com/self_check_quiz
SECTION
Galvanic Cells: Electricity
from Chemistry
17.2
S
SECTION PREVIEW
uppose it’s nearing the end of
half-time during the seventh game
of the NBA championship finals,
and your favorite team is leading by
just two points. Suddenly . . . the electricity goes out. You might find that
battery-powered radio you haven’t
used for months, but its batteries are
dead. You can’t go buy more batteries
because the electricity is out all over
town. Must you resign yourself to missing the end of this game?
Objectives
✓ Relate the construction of a galvanic cell
to how it functions to
produce a voltage and
an electrical current.
✓ Trace the movement of electrons in a
galvanic cell.
Electrochemical Cells
You don’t have to miss the end of the NBA game because you have at
hand all the ingredients for making a battery that will power your radio.
All you need are several lemons or pieces of fruit (or even glasses of fruit
juice), two different kinds of metal (a penny and a steel nail will do), and
some pieces of wire for connecting everything together. Once you have
those items, you need only a little knowledge of electrochemistry, and
you’ll soon be listening to your team go all the way.
How can it be that simple? The energy in oxidation-reduction reactions
can be harnessed to do useful work, if listening to an NBA game can be
called work. A battery is the tool that makes this possible.
When Luigi Galvani used two dissimilar metals to produce an electrical
current that stimulated the nerve in the frog leg, he didn’t know that he
had invented the first battery. Batteries are electrochemical cells. The battery made out of fruit, which allowed you to hear the end of the basketball
game, is such a cell. In a battery, the two halves of a spontaneous redox
reaction are separated and made to transfer electrons through a wire.
✓ Relate chemistry in
a redox reaction to
separate reactions
occurring at electrodes in a galvanic
cell.
Review Vocabulary
Electrolysis: process
in which electrical
energy is used to
cause a non-spontaneous chemical reaction to occur.
New Vocabulary
potential difference
voltage
galvanic cell
The Lemon Battery
How does the lemon battery produce electrical energy? The lemon itself
is a container for a solution of electrolyte—the lemon juice. As you know,
lemon juice is sour; that is, it is acidic. The hydrogen ions from partially
dissociated citric acid give it a sour taste and also provide the ions for
conduction of charge through the lemon battery. The two dissimilar metal
strips are the electrodes at which an oxidation reaction and a reduction
reaction take place to provide the battery’s power source.
17.2
Galvanic Cells: Electricity from Chemistry
599
The Lemon with Potential
Lemons are good for more than just making lemonade. By adding
some metal strips, lemons have other “potential” uses. In this MiniLab,
you will investigate the interactions of zinc and aluminum with lead
when the metals are placed in a lemon.
2
Procedure
1. Gently knead a lemon without
breaking the skin. Make two
slits about 1 cm in depth on
opposite sides of the lemon.
2. Insert a strip of zinc in one of
the slits and a strip of lead in
the other slit.
3. Connect an alligator clip wire to
each of the metal strips, touching or connecting the other end
of each wire to the poles of a
voltmeter. If the voltmeter gives
no reading, reverse the wires.
4. Read and record the voltage.
5. Repeat steps 1 through 4 with
strips of lead and aluminum,
making a new slit for the
Analysis
1. What causes the potential difference between the zinc and
lead strips?
2. Why is the potential difference
greater when aluminum is substituted for zinc?
3. If strips of zinc and magnesium,
rather than zinc and lead, were
used in the MiniLab, would the
reaction that occurs at the zinc
strip be the same? Explain.
In the lemon battery shown in Figure 17.9, a different chemical reaction
occurs at each of the metal-strip electrodes. The electrode made of the
metal that is more easily oxidized becomes the anode—the electrode at
which the oxidation reaction occurs. The second electrode becomes the
cathode, and a reduction reaction proceeds at its surface. The substance in
a lemon that is most easily reduced is the abundant hydrogen ion of the
electrolyte. When these two reactions occur together, in the same cell, they
combine to produce a spontaneous redox reaction. This type of reaction is
represented by the equation below, where M is the metal that is oxidized.
M ⫹ 2H⫹ ˇ M 2⫹ ⫹ H2
This spontaneous reaction generates the cell voltage of the battery by producing a different electrical potential at each electrode.
Figure 17.9
Chapter 17
Lemon battery
0
0.2
0.4
0.6
0.8
1
Lemon Battery
A battery can be made by
inserting iron and copper
strips into a lemon and connecting them with a conducting wire in an external circuit.
Electrons travel through the
wire by metallic conduction
and through the lemon by
electrolytic conduction.
600
aluminum and lightly buffing
the metal with fine steel wool
immediately before inserting to
remove the oxide coating.
CAUTION: Discard the lemon.
Do not use for food.
mA
Iron
strip
Electrochemistry
Copper
strip
Carolin
a
Voltmeter
Potential Difference
Electrons in the metal electrodes of the lemon battery move through
the external circuit as a current and can do useful work. The portable
radio you connected to a lemon battery to listen to the game needs power
to work, and it gets this power from the electrons flowing through the
wires from the anode to the cathode of the battery. The chemical reaction
at the anode gives off electrons, which enter the metal and then flow
through the external part of the circuit connecting the anode to the cathode. At the cathode, the electrons are used up in a reduction reaction. Just
as adding water to a container raises the level of the water, adding electrons builds up a negative potential at the anode. This electrical potential
is often described as a force, or a pressure of electrons produced by raising
the level of the electron sea, Figure 17.10.
Why do the electrons travel in one direction and not in the reverse? The
electron pressure at the cathode is kept low by the reduction reaction, and
the electrons flow from a region of high pressure (negative potential at the
anode) to a region of low pressure (positive potential at the cathode). This
potential difference between the electrodes in the lemon battery causes an
electrical current to flow. If there is no potential difference between the
electrodes, no current will flow. The size of the current depends upon the
size of the potential difference. As the electrons move from a region of
more negative potential to a region of more positive potential, they lose
energy, so the discharge of a lemon battery is a spontaneous process.
Potential energy stored in chemical bonds is released as electrical energy
and, finally, as heat. An electrical potential difference is called voltage and
is expressed in units of volts in honor of Alessandro Volta.
Wire
e–
Electric current
e–
More
positive
potential
Zn
2e –
2+
2H+
2e
–
e–
More
negative
potential
Zn
Sea of electrons
H2
Copper cathode
(reduction occurs)
2H+ + 2e –
H2
Electrolyte
(lemon juice)
Electron
pressure
Zinc anode
(oxidation occurs)
Zn
Zn2+ + 2e –
The first battery, made
by Alessandro Volta in
1800, was a simple
apparatus. It consisted
of alternating layers of
silver and zinc separated by pieces of paper
soaked in salt water.
These layers were
repeated over and over
to form a tall pile.
Wires and an acid as
the necessary electrolyte completed the
apparatus, which was
called a voltaic pile.
Electron pump
e – (battery)
ee–– e –
–
e– –e –
e
e – – M+ M
e e
+
–
e – –e – M
e –ee–
e
–
e M
e– e–
X– M+
Potential
Sea of electrons
Metal cathode
Metal anode
Electrolyte (ionic conductors)
Figure 17.10
Potential Difference
䊱 In this model of a lemon battery, the level of the electron sea is
raised or lowered by the chemical reactions at the electrode surfaces,
creating a potential difference across the battery. A spontaneous
oxidation reaction raises the electron pressure (potential) at the
anode, and a spontaneous reduction reaction reduces the pressure at
the cathode. The “sea level” in the lemon juice is uniform throughout and is intermediate between the levels at the two electrodes.
17.2
䊱 Because the redox reactions
that take place during electrolysis are not spontaneous,
a battery is needed to pump
electrons from an area of low
potential to one of high
potential.
Galvanic Cells: Electricity from Chemistry
601
Table 17.1 Ease of Oxidation of Common Metals
Easily oxidized
Not easily oxidized
Chemists at the University of California at
Irvine have made the
world’s smallest galvanic cell. It is too small to
be seen without an
electron microscope
and much smaller than
most human cells. The
galvanic cell consists of
two mounds each of
copper and silver
attached to a graphite
surface. Although it
probably will never be
used as a practical battery, it may allow scientists to study redox
reactions at the atomic
level.
602
Chapter 17
Li
K
Ca
Na
Mg
Al
Mn
Zn
Cr
Fe
Ni
Sn
Pb
Cu
Ag
Hg
Pt
Au
Li ˇ Li+ + e–
K ˇ K+ + e–
Ca ˇ Ca2+ + 2e–
Na ˇ Na+ + e–
Mg ˇ Mg2+ +2e–
Al ˇ Al3+ + 3e–
Mn ˇ Mn2+ + 2e–
Zn ˇ Zn2+ + 2e–
Cr ˇ Cr3+ + 3e–
Fe ˇ Fe2+ + 2e–
Ni ˇ Ni2+ + 2e–
Sn ˇ Sn2+ + 2e–
Pb ˇ Pb2+ + 2e–
Cu ˇ Cu2+ + 2e–
Ag ˇ Ag+ + e–
Hg ˇ Hg2+ + 2e–
Pt ˇ Pt2+ + 2e–
Au ˇ Au3+ + 3e–
Gold
Iron is readily oxidized partly because the transfer of electrons from
iron to an oxidizing agent releases a large amount of energy. You learned
that other metals also are oxidized in corrosion reactions. However, different substances release different amounts of energy when they become oxidized, and this fact may be used to construct a table such as Table 17.1. It
may be used as a general guide to the ease with which a substance will
lose electrons. By examining this table, you can see why copper, gold, and
silver are the metals most commonly used in jewelry. All three are hard to
oxidize and are, thus, resistant to corrosion.
Galvanic Cells
In the lemon battery, a redox reaction occurs spontaneously to produce
a separation of charge at the two electrodes. The reaction begins as soon
as the two electrodes are connected by a conductor so that current can
flow. An electrochemical cell in which an oxidation-reduction reaction
occurs spontaneously to produce a potential difference is called a galvanic
cell. In a galvanic cell, chemical energy is converted into electrical energy.
Galvanic cells are sometimes called voltaic cells; both terms refer to the
same device. A galvanic cell that has been packaged as a portable power
source is often called a battery.
Sometimes, the chemical change taking place in a galvanic cell can be
seen easily, such as in the simple magnesium-copper galvanic cell shown in
Figure 17.11. Because magnesium is more easily oxidized than copper, the
magnesium loses electrons and becomes oxidized, forming Mg2⫹ ions. The
potential of the magnesium anode becomes more negative because of the
Electrochemistry
increased electrical pressure from the released electrons. At the same time,
the Cu2⫹ ions pick up electrons from the copper electrode and are reduced
to copper metal. The potential of the copper electrode becomes more positive because electrical pressure is lowered as electrons are removed from
the cathode. If a wire is connected between the electrodes, current flows
from the magnesium electrode to the copper electrode, and the voltmeter
in the external circuit reads a voltage of 2.696 V. The energy released during discharge of the cell can be used to power a device such as a radio by
connecting the wire from the electrodes through the radio. The overall
reaction in the copper-magnesium cell is a redox reaction.
Cu2⫹(aq) ⫹ 2e⫺ ˇ Cu(s)
Mg(s) ˇ Mg2⫹(aq) ⫹ 2e⫺
Oxidation half-reaction
Reduction half-reaction
Mg(s) ⫹ Cu2⫹(aq) ˇ Mg2⫹(aq) ⫹ Cu(s)
Net redox reaction
Figure 17.11
Magnesium-Copper Galvanic Cell
A piece of magnesium metal is placed in a
beaker containing a solution of magnesium sulfate, and a piece of copper metal
is placed in a beaker containing a solution
of copper(II) sulfate. The two beakers are
connected via a salt bridge, which is a
porous barrier containing a salt solution;
this prevents the two solutions from mixing but permits the movement of ions
from one side of the cell to the other. An
external circuit containing a voltmeter
connects the two metal electrodes. 䊳
e–
0
0.5
1
e–
1.5
2
2.5
Salt bridge
3
Magnesium
anode
Copper
cathode
mA
Cl – Na+
Voltmeter
SO42–
SO42–
Mg2+
Mg
Mg2+ + 2e –
Oxidation
Cu2+ Cu
Cu2+ + 2e –
Cu
Reduction
Mg
䊴 As the blue copper ions are reduced, they move out of
the solution and are deposited as copper metal on the
copper strip. As a result, the blue copper solution
becomes less and less blue-colored. Positively charged
sodium and negatively charged chloride ions move out
of the salt solution in the salt bridge to restore charge
balance in the beakers.
17.2
Galvanic Cells: Electricity from Chemistry
603
Figure 17.12
Magnesium-Copper Redox
Reaction
When magnesium metal is
added to a blue solution of
CuSO4 , both the magnesium
metal and the blue color
disappear. 䊳
The more easily oxidized magnesium forms colorless
Mg2⫹ ions, which dissolve in the solution. The blue
copper(II) ions are reduced to the red-brown copper
metal that can be seen at the bottom of the beaker. 䊳
spontaneous:
sponte (L) of free
will
A spontaneous
reaction arises
from the inherent
qualities of the
reactants and usually occurs with no
external input of
energy.
604
Chapter 17
The same overall redox reaction occurs if the magnesium metal is placed
directly into a solution of copper sulfate, Figure 17.12. However, this is not
a galvanic cell because the electrons do not flow through an external circuit. Instead, the electrons move directly from the magnesium metal to the
copper ions, forming copper metal. This is a way to make copper metal
from copper ions, but it is not a way to make electrical power.
You can see that for every spontaneous redox reaction, you theoretically
can construct a galvanic cell that can capture the energy released by the
reaction. The amount of energy released depends upon two properties of
the cell: the amount of material that is present and the potential difference between the electrodes. The more material there is in the electrode,
the more electrons it can produce during the course of the reaction. The
potential difference depends upon the nature of the reaction that takes
place; that is, it corresponds to the relative positions of the two substances
in a table such as Table 17.1. The farther apart the two substances are in
the table, the greater the potential difference between the electrodes, and
the greater the energy delivered by each electron that flows through the
external wire.
How do you know which substance will be oxidized and which reduced
in any cell? Look back at Table 17.1. Experimental chemists such as
Humphry Davy and his student Michael Faraday did many experiments
from which this type of table could be made. The table is used today to
predict the outcome of new experiments. For example, in a Zn-Cu galvanic cell, zinc will be oxidized and copper reduced. Because zinc is more easily oxidized than copper, electrons will flow from zinc to copper.
A cell voltage should register on the voltmeter shown in Figure 17.11
because a potential difference exists between the magnesium and copper
electrodes. What function does the salt bridge serve? As the half-reactions
continue, magnesium ions are released into the solution at the anode, and
copper ions are removed at the cathode. Ions must be free to move
between the electrodes to neutralize positive charge (Mg2⫹ cations) created at the anode and negative charge (anions) left over at the cathode. The
solution of ions in the salt bridge allows ionic conduction to complete the
electrical circuit and prevent a buildup of excess charge at the electrodes.
Electrochemistry
Mg
anode
e–
Figure 17.13
e–
Cu
cathode
Cl–
Na+
Salt
bridge
SO42–
SO42–
Mg
Mg
Mg2+ + 2e –
Cu
Cu2+ + 2e –
Batteries Perform Work
When a simple galvanic cell does useful
work, it is called a battery. If the external
circuit is connected with a wire, electrons
flow from the site of oxidation at the magnesium strip and through the LED to the
surface of the copper strip, where reduction of Cu2⫹ ions takes place. The voltage
pushes electrons through the LED, causing
it to light up.
Cu
Useful work may be done if the voltmeter is replaced by wires connecting the galvanic cell to a lightbulb. Then, electrical energy will be converted into light energy, a useful process in a dark room. In Figure 17.13,
wires connect the cell to a light source with a low-voltage requirement,
called a light-emitting diode or LED. If the circuit to the cell is complete,
the LED lights up, showing that the cell is doing useful work. With time,
the light intensity will fade. Why doesn’t it stay lit indefinitely? Eventually,
all of the magnesium in the anode becomes oxidized. The capacity of the
battery has been exceeded, the magnesium is gone and, if there is no electrode, there can be no cell.
Figure 17.14
Better and Better Batteries
Although the galvanic cell made from magnesium and copper can do
useful work, it isn’t something you’d want to bring along on a camping
trip. The wet solutions could be sloppy, the glass could break easily, and
the capacity is limited. Fortunately, scientists have developed much better
batteries that are smaller, lighter, provide higher voltages,
and last longer. Figure 17.14 shows an assortment of commonly used batteries. Experimental batteries no thicker
than a sheet of paper have already been developed. And,
although you might think batteries always have to be
made of metal and acids, some batteries of the future
may be made of microorganisms that use the energy in
sugar to make electricity. A living fuel cell has been developed that someday could be used to power an automobile
for up to 15 miles on two pounds of sugar.
How are batteries designed? The farther apart two metals are in Table 17.1, the larger the voltage of a battery
that can be constructed from them. If you wanted to
make a high-voltage battery to power your radio, you
would choose metals that are far apart in the table. A copper penny with an iron nail will yield a larger voltage than a
penny with a piece of nickel because copper is farther away from
iron in the table than it is from nickel.
17.2
Modern Batteries
Modern batteries come in a
wide variety of sizes, shapes,
and strengths. Each type of
battery serves a different
purpose.
Galvanic Cells: Electricity from Chemistry
605
Oxidation-Reduction and
Electrochemical Cells
Redox reactions involve the loss and gain of
electrons. By separating the oxidation process
from the reduction process and connecting them
electrically through an external circuit, many
spontaneous redox reactions can be utilized to
produce an electrical potential and an electrical
current. Devices that perform these functions are
called electrochemical cells. In this ChemLab, you
will investigate a redox reaction and use it to construct an electrochemical cell.
Problem
How may a spontaneous redox reaction be
used to construct an electrochemical cell?
Objectives
Observe a simple oxidation-reduction reaction.
Relate the reaction to the oxidation tendencies
of the reactants.
Utilize the reaction to construct an electrochemical cell that can operate electrical devices.
•
•
•
Materials
craft-stick support with V-cut and slit cut
25-mm (flat diameter) dialysis tubing (15 cm in
length)
magnesium ribbon (10 cm in length)
magnesium ribbon (1 cm in length)
copper foil (10 cm ⫻ 1 cm strip)
copper foil (1 cm ⫻ 2 mm piece)
metric ruler
250-mL beaker
10 ⫻ 100 mm test tubes (2)
wire leads with alligator clips (2)
DC voltmeter with a 2-V or 3-V scale
606
Chapter 17
Electrochemistry
flashlight bulb for 2 AAA batteries
9-V transistor radio
0.5M sodium chloride solution
0.5M copper(II) chloride solution
0.1M magnesium chloride
Safety Precautions
Wear an apron and safety goggles. Rinse the
solutions down the drain with large amounts of tap
water. Wash your hands after performing the lab.
1. Soak the dialysis tubing in tap water for about
ten minutes while you complete steps 2 and 3.
Tie two knots near one end of the tubing, and
open the other end by sliding the material
between your fingers.
2. Pour a small amount of the copper(II) chloride solution into a 10 ⫻ 100 mm test tube,
and drop a 1-cm length of magnesium ribbon
into the solution. Observe the system for
about one minute, and record your observations in a data table similar to the table in
Data and Observations. Pour the solution
down the drain and discard the piece of magnesium in a wastebasket.
3. Repeat step 2 using magnesium chloride solution and the small piece of copper foil.
Mg ribbon
Craft stick support
Cu foil
CuCl2
solution
in dialysis
tubing
NaCl
solution
Knots
4. Pour copper(II) chloride solution into the
open end of the tubing to a depth of 6 cm to
8 cm, and insert the strip of copper foil. Slide
the top of the tubing and copper strip into the
V-cut in the stick as shown below. Suspend
the tubing in the beaker, as shown.
5. Slide the length of magnesium ribbon into the
slit cut in the craft stick, as shown.
magnesium, and dispose of them according to
the instructions of your teacher.
Craft stick support
Mg – Cu2+
Data and Observations
Cu – Mg2+
V-cut for
tubing and Cu
Slit cut for Mg
6. Pour about 200 mL of sodium chloride solution into the beaker.
7. Connect leads to the pieces of copper and
magnesium, and touch the leads to a DC voltmeter—the lead from the Cu electrode to the
⫹ terminal and the lead from the Mg electrode to the ⫺ terminal. Read and record the
potential difference, or voltage.
8. Cooperate with other lab groups in the following way to light the bulb and to operate the
transistor radio. The flashlight bulb requires a
voltage of about 3V, and the radio requires a
voltage of about 9V. Connect your electrochemical cells in series (copper to magnesium)
to provide the desired voltages. CAUTION: Be
sure to connect your cells to the battery terminals of the radio in the correct polarity. Connected in series, the voltages are additive; for
example, five 2-V cells in series yield a voltage
of 10V. Such combinations of electrochemical
cells are called batteries.
9. Disassemble your cell, observing the pieces of
copper and magnesium and recording your
observations. Rinse the pieces of copper and
Voltage
Pieces of Cu and Mg
1. Interpreting Data Write the balanced equation for the single-replacement reaction
between magnesium and copper(II) chloride
that occurred in step 2. Which metallic element, Cu or Mg, has the greater tendency (or
oxidation potential) to lose electrons?
2. Relating Concepts In an electrochemical cell,
oxidation occurs at the anode, and reduction
occurs at the cathode. Which metal was the
anode and which was the cathode? Write the
equations for the half-reactions.
1. When an electrochemical cell is used to operate an electrical device, in which direction do
the electrons move in the external circuit?
2. Is it possible to construct an electrochemical
cell in which lead is the anode and lithium is
the cathode? Explain.
17.2
Galvanic Cells: Electricity from Chemistry
607
If you have metal fillings in your teeth, you
may have received an
electric shock after biting down on a bit of
aluminum foil. The aluminum foil acts as an
anode, and the silvermercury amalgam, an
alloy, in the filling acts
as a cathode. The saliva
in your mouth serves as
an electrolyte. The galvanic cell in your
mouth is short-circuited
when the foil comes
into contact with the
filling, causing a weak
current to flow
between the electrodes. This current is
detected as pain by the
nerve of the tooth.
Although the term battery usually refers to a series of galvanic cells connected together, some batteries have only one such cell. Other batteries may
have a dozen or more cells. When you put a battery into a flashlight, radio,
or CD player, you complete the electrical circuit of the galvanic cell(s),
providing a path for the electrons to flow through as they move from the
reducing agent (the site of oxidation) to the oxidizing agent (the site of
reduction). The most powerful batteries combine strong oxidizing agents
and strong reducing agents to give the largest possible potential difference.
But those agents aren’t necessarily safe, convenient, or economical to use. To
get a higher voltage from a cell type with a relatively small potential difference, several of the cells can be connected in series, as Figure 17.15 shows.
FM
AM
Alligator
clip
Plain wire
(without
washer)
Figure 17.15
A More Powerful Lemon Battery
One lemon cell wouldn’t provide enough voltage to power a
transistor radio, but several connected together in series would.
This means that the positive terminal of one cell is connected to
the negative terminal of the next, and so on. The electrodes in
this battery are steel washers and copper wire. The total voltage
of the battery is the sum of all the voltages of the individual cells.
Carbon-Zinc Dry Cell
Whenever you put two or more common D batteries into a flashlight,
you are connecting them in series. They have to be placed in the correct
order so that electrons flow through both cells. These relatively inexpensive batteries are carbon-zinc galvanic cells, and they come in several
types, including standard, heavy-duty, and alkaline. This type of battery is
often called a dry cell because there is no aqueous electrolyte solution; a
semisolid paste serves that role. Examine the cutaway view of the carbonzinc battery in Figure 17.16 to see if you can locate the parts of the galvanic cell it contains.
608
Chapter 17
Electrochemistry
Figure 17.16
Carbon-Zinc Dry Cell
A standard D battery is shown both
whole and cut in half to reveal
the structure of the carbonzinc dry cell. Beneath the outside paper cover of the battery is a cylinder casing made
of zinc. The zinc serves as the
anode and will be oxidized in
the redox reaction. 䊲
Carbon cathode
surrounded by
black MnO2 -carbon
paste
Porous spacer
NH4Cl-ZnCl2
electrolyte
Zinc shell (anode)
䊱 The carbon rod in the center of the cylinder—surrounded by a moist, black paste
of manganese(IV) oxide (MnO2) and carbon black—acts as a cathode. Ammonium chloride (NH4Cl) and zinc chloride
(ZnCl2) serve as electrolytes. Alkaline batteries contain potassium hydroxide
(KOH) in place of the ammonium chloride electrolyte, and they maintain a high
voltage for a longer period of time.
What is missing in this galvanic cell? Notice that the circuit is not complete, so the electrons that are produced at the zinc cylinder have no
external conductor through which to travel to the carbon. This is by
design and is not a defect in the battery. The circuit will be complete
when the battery is placed in something designed to be powered by it,
such as a flashlight. When the flashlight is turned on, the redox reaction
starts. Electrons travel out of the zinc casing into a piece of metal built
into the flashlight. There, they travel through a bulb, causing it to light up.
The electrons then reenter the battery at the top and move down through
the carbon rod and into the black paste, where they take part in the
reduction reaction.
The flow of electrons from the zinc cylinder through the electrical circuits of an appliance and back into the battery provides the electricity
needed to power a flashlight, radio, CD player, toy, clock, or other item.
When electrons leave the casing, zinc metal is oxidized.
Zn ˇ Zn2⫹ ⫹ 2e⫺
The reactions in the carbon rod and the paste are much more complex,
but one major reduction that takes place is that of manganese in manganese(IV) oxide. In this reaction, the oxidation number of manganese is
reduced from 4⫹ to 3⫹.
2MnO2 ⫹ H2O ⫹ 2e⫺ ˇ Mn2O3 ⫹ 2OH⫺
Adding the two half-reactions together gives the major redox reaction taking place in a carbon-zinc dry cell.
Zn ⫹ 2MnO2 ⫹ H2O ˇ Zn2⫹ ⫹ Mn2O3 ⫹ 2OH⫺
17.2
Galvanic Cells: Electricity from Chemistry
609
HEALTH
CONNECTION
Lithium Batteries in Pacemakers
It’s always frustrating to have batteries go dead just when it seems you need
them most. However, imagine needing a battery upon which your life depends.
What materials could it be made from? Would it last long enough?
Heart stimulation Consider for a moment that, over a
period of time, you began to feel light-headed, dizzy,
weak, or fatigued. It could be that the chambers of
your heart are not beating rhythmically or fast enough.
You might be a candidate for a heart pacemaker. This
device, which is inserted inside the body, monitors the
heart’s activity. When necessary, the pacemaker supplies the electrical impulses needed to stimulate the
heart. In order to be most effective, the batteries in a
pacemaker need to be fully powered for long periods of
time and survive in the hostile, saline environment of
the human body without breaking down.
Powerful electrodes One type of battery used in pacemakers is a cell based on
lithium and sulfuryl chloride. Lithium is a popular choice for battery anodes
because of its strong tendency to be oxidized. Lithium is oxidized during the
reaction, and the sulfur in sulfuryl chloride is reduced. The unbalanced halfreactions are given below.
oxidation
Li(s) ˇ Li⫹(aq) ⫹ e⫺
SOCl2(l) ⫹ 4e⫺ ˇ Cl⫺(aq) ⫹ S(s) ⫹ SO2(g) reduction
Battery characteristics Lithium is the least dense of all nongaseous elements,
with a density of only 0.534 g/mL. The lightweight lithium contributes little
weight to the small, detachable battery pack, which is a circular disk only about
5 cm by 1 cm in size. The pacemaker battery pack is implanted under the
patient’s collarbone and has a life expectancy of about seven to ten years,
depending upon how often it is needed to stimulate the heart.
Lithium hazards Lithium also presents some potential problems when used in a
wet environment. The tremendous activity of lithium makes it dangerously reactive with a variety of compounds, including
water. If the battery pack is not adequately
sealed against the body’s corrosive fluids, the
Connecting to Chemistry
resulting exothermic reaction of lithium in
water accompanied by the production of H2
1. Interpreting Write 2. Interpreting Write
gas could have serious consequences.
the balanced overall
the balanced overall
As battery technology continues to
reaction of lithium
reaction of lithium
advance, medical applications requiring
with sulfuryl chlowith water. What is
portable, long-lasting power supplies will
ride.
the oxidizing agent?
become more and more common.
The reducing agent?
610
Chapter 17
Electrochemistry
Automobile Lead Storage Battery
The most common type of battery used in cars is a lead-acid, 12-volt
storage battery. It contains six 2-volt cells connected in series. Although
much larger than carbon-zinc batteries and relatively heavy, this type of
battery is durable, supplies a large current, and can be recharged. When
you turn your key in the ignition, it is the battery that supplies electricity
to start the car. It also provides energy for any demands not met by the
car’s alternator, such as running the radio or using the lights when the
engine is off. Leaving on the lights or radio for too long with the engine
off can make the battery go dead because it is the engine that recharges
the battery as the car runs.
Each galvanic cell in a lead-acid battery has two electrodes—one made
of a lead(IV) oxide (PbO2) plate and the other of spongy lead metal, as
Figure 17.17 shows. In each cell, lead metal is oxidized as lead(IV) oxide is
reduced. The lead metal is oxidized to Pb2⫹ ions as it releases two electrons at the anode. The Pb4⫹ ions in lead oxide gain two electrons, forming Pb2⫹ ions at the cathode. The Pb2⫹ ions combine with SO42⫺ ions
from the dissociated sulfuric acid in the electrolyte solution to form
lead(II) sulfate at each electrode. Thus, the net reaction that takes place
when a lead-acid battery is discharged results in the formation of lead sulfate at both of the electrodes.
PbO2 ⫹ Pb ⫹ 2H2SO4 ˇ 2PbSO4 ⫹ 2H2O
⫹
Cathode
Anode
⫺
Cell
connector
Cell spacer
Pb
PbO2
Figure 17.17
Lead Storage Battery
䊴 A lead storage battery is not a dry cell because it contains several connected cells filled with an aqueous
solution of sulfuric acid, which serves as the electrolyte.
The electrodes are alternating plates of lead metal and
lead(IV) oxide. The case surrounding the battery is
hard plastic. It holds the cells in place and acts as an
insulator because it does not conduct electricity itself.
This helps keep the electricity inside the battery.
Cell with
electrolyte
A car with a dead battery can still
be started. Electricity from a second
car is used to jump-start the car,
bypassing the dead battery. 䊳
17.2
Galvanic Cells: Electricity from Chemistry
611
Nicad Rechargeable Batteries
The nickel-cadmium, or nicad, cell is a common storage
battery that can usually be discharged and recharged
more than 500 times. These batteries are used
in calculators, cordless power tools and vacuum cleaners, and rechargeable electric toothbrushes and
shavers. Once nicad batteries have been spent, disposal presents a problem because cadmium is toxic. Nicads can be
recycled, but the process is expensive. Although rechargeable
batteries containing less toxic metals are being developed,
none have been found that can sustain a constant rate of discharge as well as the nicad.
Cap
Vent ball
1. The nicad
galvanic cell
has cadmium
anodes,
hydrated
nickel oxide
cathodes,
and KOH as
the electrolyte. The
electrodes
are arranged
in jelly-roll
fashion.
2. In the redox reaction that takes place during discharge, nickel oxide is reduced at the cathode, and
cadmium is oxidized at the anode.
Cover
Seal
Cd ⫹ NiO ˇ CdO ⫹ Ni
3. The electrolysis reaction that takes place when an
external source of electricity is used to recharge
the cell is the reverse of the discharge reaction.
CdO ⫹ Ni ˇ Cd ⫹ NiO
Core
Positive tab
KOH electrolyte
Pressed powdered
cadmium anode
BATTERY
BATTERY
Separators
Insulating
washer
612
Chapter 17
Electrochemistry
Can
charged before use.
5. Nicad batteries are not suitable for
Nickel oxide
cathode
”Jelly roll”
4. Newly purchased nicad batteries must be
devices that are left idle for long stretches—such as smoke detectors, cameras,
and flashlights—because they will lose
about one percent of their charge daily
even when not being used.
Thinking Critically
1. What are the equations for the oxidation and reduction
half-reactions that
occur during
recharging of a nicad
battery?
2. What might be an
environmental
advantage to using
nicad batteries?
The reaction that occurs during discharge of a lead-acid battery is
spontaneous and requires no energy input. The reverse reaction, which
recharges the battery, is not spontaneous and requires an input of electricity from the car’s alternator. Current enters the battery and provides energy for the reaction in which lead sulfate and water are converted into
lead(IV) oxide, lead metal, and sulfuric acid.
2PbSO4 ⫹ 2H2O ˇ PbO2 ⫹ Pb ⫹ 2H2SO4
Sulfuric acid is corrosive. It is important to be careful when working
around a car battery, as well as disposing of it properly when it finally
goes dead for good. These batteries can usually be discharged and
recharged numerous times and last about three to five years.
Better Batteries for Electric Cars
At the end of the 19th century, most cars were powered by steam or by
electric batteries, although today most cars have internal combustion
engines that are powered by gasoline. Electric cars, Figure 17.18, could
help reduce our dependence on fossil fuels, cause less pollution, and be
more economical in the long run, but they have several disadvantages,
such as high initial cost, limited driving range, low speed, and long
recharge time. They also present a disposal problem because cadmium is a
toxic metal.
These disadvantages would disappear if a battery that is cheap enough,
powerful enough, and safe enough for running an electric car could be
developed. Two new experimental types of batteries for use in electric
cars show early promise as candidates. One is a rechargeable, nickelmetal hydride or NiMH battery. This type of battery is less toxic and has
a higher storage capacity than the batteries now used in electric cars.
Another experimental battery is a lithium battery with a water-based
electrolyte. Lithium is more easily oxidized than any other metal but has
a drawback that has limited its use in batteries: it explodes violently
when it comes into contact with water. Lithium is used in some batteries
Figure 17.18
Electric Cars
Most of the nickel-cadmium
batteries that are used in electric cars today power the car
for only 50 to 100 miles
before they run down and
need to be recharged, a
process that takes many
hours. In Randers, Denmark,
special parking spaces with
electric hookups are available
for battery-operated cars.
17.2
Galvanic Cells: Electricity from Chemistry
613
Hydrogen-Oxygen Fuel Cell
e–
Recall that the combustion of a fuel is a redox reaction in
which the fuel molecules are oxidized and oxygen is reduced
to form an oxide. For years, scientists have worked to find a
way to separate the oxidation and reduction reactions to
make them produce an electric current. The simplest fuel cell
involves the oxidation of the fuel hydrogen gas to form water.
Today, hydrogen-oxygen fuel cells are used to supply electricity to the space shuttle orbiters. The fuel cells have a weight
advantage over storage batteries, and the water produced during their operation can be used for drinking.
1. A simple hydrogen-oxygen fuel cell differs in
e–
O2 in
H2 in
Platinum
(anode)
Platinum
(cathode)
Electrolyte
solution
5. Water vapor—steam—is produced in
two major ways from a galvanic cell: the electrodes are made of an inert material that
doesn’t react during the process, and hydrogen and oxygen gas are fed in continuously.
the fuel cell, as up to 75 percent of the
chemical energy is converted into electricity. The steam can be condensed
and used for drinking water.
2H2(g) ⫹ O2(g) ˇ 2H2O(g) ⫹ energy
2. Hydrogen is fed
onto an electrode
on one side of the
fuel cell, and oxygen is fed onto an
electrode on the
other side.
Hot potassium hydroxide
electrolyte (approx. 200°C)
Hydrogen
H2
Anode
O
yyy
yyyyy
yyy
yyyyy
yyy
yyyyy
yyy
yyyyy
yyy
yyyyy
yyy
yyyyy
yyy
yyyyy
yyyyy
yyy
yyyyy
yyy
yyy
yyy
yyyyy
yyy
yyy
yyyyy
yyyyyy
yyy
yyy
yyy
yyy
yyy
yyy
yyy
yyy
yyy
Porous carbon electrodes
e–
4. The electrons lost by hydrogen molecules, which are oxidized at the anode,
flow out of the fuel cell, through a circuit, and then back into the fuel cell at
the cathode, where oxygen is reduced.
614
Chapter 17
6. If more inexpensive
and longer-lasting fuel
cells can be developed,
they may someday produce electricity in
power plants.
Waste steam
3. Concentrated
KOH serves as the
electrolyte in the
fuel cell.
Oxygen
Electrochemistry
2
Cathode
e–
Thinking Critically
1. What causes electrons to flow from
hydrogen to oxygen
in a fuel cell?
2. If fuel cells are about
75 percent efficient,
what happens to the
rest of the potential
energy?
Figure 17.19
ANODE
e⫺
CATHODE
e⫺
Electrolyte
Electron
Lithium atom
e⫺
e⫺
Lithium ion
Aqueous Lithium Battery
How can a lithium battery have an aqueous
electrolyte? Two facets of the construction of
this new battery keep the lithium metal from
reacting with water. First, the lithium is in the
form of individual atoms embedded in a
material such as manganese(IV) oxide, rather
than as a solid metal. Second, the electrolyte
is full of dissolved lithium salts, so the lithium
ions that are produced travel to the site of
reduction without reacting with water.
Manganese(IV)
oxide
to power camcorders, but they require an expensive, nonaqueous electrolyte. Figure 17.19 shows the construction of the experimental aqueous
lithium battery. This battery is less toxic and will probably be cheaper to
manufacture than the nickel-cadmium batteries used in most electric
cars in operation today.
Connecting Ideas
In your study of electrochemistry, you have seen how chemical reactions in batteries can be used to generate electricity. However, most of the
electricity you use comes from another chemical source—the fossil fuels
petroleum, natural gas, and coal. Although inorganic chemicals are usually used to fuel batteries, fossil fuels are a major source of a large group of
chemicals, the carbon-containing organic compounds. In the next chapter, you’ll learn that organic chemicals also provide us with most medicines, dyes, plastics, and textiles.
SECTION REVIEW
Understanding Concepts
For more practice
with solving problems,
see Supplemental Practice
Problems, Appendix B.
a) Use Table 17.1 to predict which metal will be
reduced and which will be oxidized.
b) Write an equation for the net redox reaction
that occurs. HINT: Cu2⫹ is formed.
c) Is this system a galvanic cell? Explain.
1. Describe the movement of electrons in a galvanic cell.
2. Draw a diagram of a simple galvanic cell.
3. How are zinc-carbon and lead-acid batteries
different?
Applying Chemistry
Thinking Critically
5. Dry Cells A dry cell cannot really be dry.
Explain why.
4. Using a Table A piece of copper metal is placed
in a 1M solution of silver nitrate (AgNO3).
chemistryca.com/self_check_quiz
17.2
Galvanic Cells: Electricity from Chemistry
615
CHAPTER 17 ASSESSMENT
REVIEWING MAIN IDEAS
17.1 Electrolysis: Chemistry from Electricity
■
■
■
■
■
An electrolytic cell is a chemical system that
uses an electric current to drive a nonspontaneous redox reaction. Electrolysis is
the process that takes place in such a cell.
An electrical current is the flow of charged
particles such as electrons.
Reduction takes place at the cathode in an
electrolytic cell.
Oxidation takes place at the anode in an electrolytic cell.
Electrolysis can be used to produce compounds, separate metals from ores, clean
metal objects, and plate metal coatings onto
objects.
■
■
A galvanic cell is a chemical system that produces an electric current through a spontaneous redox reaction.
Batteries contain one or more galvanic cells.
Vocabulary
For each of the following terms, write a sentence that shows
your understanding of its meaning.
anion
anode
cathode
cation
electrical current
electrolysis
electrolytic cell
galvanic cell
potential difference
voltage
17.2 Galvanic Cells: Electricity from Chemistry
■
A potential difference between two substances
is a measure of the tendency of electrons to
flow from one to the other.
UNDERSTANDING CONCEPTS
1. What is the function of the salt bridge in a galvanic cell?
2. What is the difference between an electrolytic
cell and a galvanic cell?
3. What is a galvanic cell?
4. What happens to the case of a carbon-zinc dry
cell as the cell is used to produce an electric
current?
5. Why is the electrolyte necessary in both galvanic and electrolytic cells?
6. What products are formed from the electrolysis of an aqueous solution of rock salt?
7. What is the function of the acid in the leadacid storage battery used in cars?
8. By what process can chlorine gas be prepared
commercially?
616
Chapter 17
Electrochemistry
APPLYING CONCEPTS
9. What will happen if a rod made of aluminum
is used to stir a solution of iron(II) nitrate?
10. Can a solution of copper(II) sulfate be stored
in a container made of nickel metal? Explain.
11. How would gold-electroplated jewelry compare to jewelry made of solid gold in terms of
price, appearance, and durability?
12. Tests conducted on different types of common
commercial batteries involved measuring the
voltage drop over time during simulated nonstop use of a motorized toy. Based on the following graph of data obtained from testing
rechargeable, alkaline, and heavy-duty batteries, which battery type would be best to use if
you wanted to run the toy for a long period of
time? Which battery type goes dead abruptly?
chemistryca.com/vocabulary_puzzlemaker
CHAPTER 17 ASSESSMENT
1.50
THINKING CRITICALLY
Volts
1.40
1.30
Making Predictions
20. MiniLab 1 Would it be possible to plate a silver spoon or a gold spoon with copper?
Rechargeable
1.20
1.10
Heavyduty
1.00
Alkaline
0.90
0
1
2
3
Hours
4
5
Drawing Conclusions
21. MiniLab 2 Why should the metal pieces used
as electrodes in the lemon battery be cleaned
with steel wool?
6
13. What would you expect to see if you placed
a) a strip of copper metal into a solution of
zinc sulfate?
b) a strip of zinc metal into a solution of copper sulfate?
14. How could lead be removed from drinking
water by electrolysis?
Everyday Chemistry
15. The information on a CD stamper master is the
reversal of that on the original glass master cut
by the recording laser. Explain why this reversal
is necessary.
Chemistry and Technology
16. Should the anode or the cathode be made of
pure copper in an electrolytic cell designed for
refining copper metal? Explain.
Health Connection
17. What happens to lithium metal if it comes into
contact with water?
How it Works
18. What are the advantages and disadvantages of
using rechargeable batteries instead of conventional types?
How it Works
19. Write the equations for the two half-reactions
that take place in a hydrogen-oxygen fuel cell.
chemistryca.com/chapter_test
Relating Cause and Effect
22. ChemLab How would your result in step 7 of
the ChemLab have been different if
a) a piece of zinc were used instead of the
piece of magnesium?
b) silver were used instead of copper?
Making Decisions
23. What factors must be considered in designing
or selecting batteries for the following applications?
a) flashlight
c) pacemaker
b) hearing aid
d) toy car
Making Predictions
24. If a strong tendency to be oxidized were the only
consideration, what metals other than lithium
might be used to power a cardiac pacemaker?
Forming a Hypothesis
25. Why was electrophoresis rather than electrolysis used to restore ceramic and organic artifacts
from the Titanic?
CUMULATIVE REVIEW
26. Draw Lewis electron dot structures for the ions
listed. (Chapter 2)
b) Cl⫺
c) OH⫺
d) O2⫺
a) Ca2⫹
27. List the names and symbols of all of the noble
gases. (Chapter 8)
28. Which requires more energy: boiling 100 g of
water or melting 100 g of ice? Explain.
(Chapter 10)
Chapter 17
Assessment
617
CHAPTER 17 ASSESSMENT
29. A 0.543-g piece of magnesium reacts with
excess oxygen to form magnesium oxide in the
reaction
2Mg(s) ⫹ O2(g) → 2MgO(s).
How much oxygen reacts? What is the mass
of magnesium oxide, MgO, produced? (Chapter 12)
30. What are the mass percents of magnesium and
oxygen in magnesium oxide in the problem
above? (Chapter 12)
31. Compare the hydronium ion concentrations in
two aqueous solutions that have pH values of 9
and 11. (Chapter 14)
SKILL REVIEW
32. Interpreting Scientific Illustrations A process
called cathodic protection is sometimes used
to protect a buried steel pipeline from corrosion. In this process, the pipeline is connected
to a more active metal such as magnesium,
which is corroded preferentially before the
iron. The diagram below illustrates how the
two metals are connected and shows the reactions that take place.
a) What acts as the cathode in this process?
What acts as the anode?
b) What is the oxidizing agent?
c) Write a short summary describing how the
magnesium is preferentially corroded.
Magnesium rod
Moist
soil
O2 ⫹ 2H2O ⫹ 4e⫺
4OH⫺
Steel pipe
Mg
618
Chapter 17
Mg2⫹ ⫹ 2e⫺
Electrochemistry
33. Making Scientific Illustrations Draw a diagram of a galvanic cell in which the reaction is
Ni(s) ⫹ 2Ag⫹(aq) ˇ Ni2⫹(aq) ⫹ 2Ag(s).
Label the cathode and the anode. Show the
ions present in both compartments, and indicate the direction of electron flow in the external circuit.
WRITING IN CHEMISTRY
34. Write an article about the development and
uses for the Daniell cell, an early battery made
in 1836 by John Frederic Daniell of Great
Britain. Find out how it improved on the Volta
cell and whether or not this type of battery is
used much today.
PROBLEM SOLVING
35. How many grams of aluminum oxide are
needed to produce 10 million metric tons
(1 metric ton ⫽ 1000 kg) of aluminum each
year using the Hall-Héroult process? Assume
100 percent yield.
36. In one type of fuel cell, methane gas (CH4) is
“burned” electrochemically to produce electricity:
CH4(g) ⫹ H2O(g) → CO(g) ⫹ 3H2(g) ⫹ energy
a) Is this a galvanic or an electrolytic cell?
b) What acts as the oxidizing reagent?
c) What acts as the reducing reagent?
d) If 224 L of natural gas are burned in this
fuel cell at STP, how many moles of carbon
monoxide and hydrogen gases are produced?
37. What will happen to your gold ring if you leave
it sitting in a solution of iron(II) chloride
(FeCl2) at room temperature?
38. What will happen to a copper bracelet that
remains in contact with a solution of silver
nitrate (AgNO3) for several hours at room
temperature?
Standardized Test Practice
1. Electrolysis is
a) the process of using electrical currents to
speed up redox reactions.
b) the process of using electrical currents to
speed up any reaction.
c) the process of using electrical currents to
start reactions that do not occur.
d) the process of using electrical currents to
stop reaction from occurring.
2. The part of an electrochemical cell that carries
electrons to a reacting ion is a(n)
a) cathode.
b) anode.
c) electrode.
d) electrolyte.
3. Which of the following processes is an application of electrolysis?
a) electroplating
b) cleaning toxic wastes
c) separating metals from ores
d) all of the above
4. Why does electric current flow between the
two electrodes of a lemon battery?
a) The two electrodes have a potential difference maintaining a flow of electrons from
an anode to a cathode.
b) The two electrodes have a potential difference maintaining a flow of electrons from a
cathode to an anode.
c) A battery attached to the two electrodes
creates a potential equivalence between the
two different metals.
d) A battery attached to the two electrodes
creates a potential difference between the
two different metals.
5. An electrical potential difference is also called
a) oxidation.
b) reduction.
Standard Reduction Potentials
Metal
Standard Reduction Potential
Li
⫺3.0401
Al
⫺1.662
Cu
0.521
Ag
0.7996
Au
1.498
ing oxidation potential. Which metal will
most easily corrode?
a) Li
c) Cu
b) Al
d) Au
7. What would happen if a lemon battery were
constructed using a copper strip and a gold
strip?
a) Electrons would flow from the copper strip,
where reduction would occur, to the gold
strip, where oxidation would occur.
b) Electrons would flow from the copper strip,
where oxidation would occur, to the gold
strip, where reduction would occur.
c) Electrons would flow from the gold strip,
where reduction would occur, to the copper
strip, where reduction would occur.
d) Electrons would flow from the gold strip,
where oxidation would occur, to the copperstrip, where reduction would occur.
8. A Galvanic cell packaged and sold as a
portable power source is called a(n)
a) voltaic cell.
b) electrode.
c) battery.
d) electrochemical cell.
c) corrosion.
d) voltage.
Test Taking Tip
Focus
Use the table to answer questions 6 and 7.
6. The table of standard reduction potential of
substances lists common metals in order of
increasing reduction potential and decreas-
chemistryca.com/standardized_test
When you take a test, pay no attention to
anyone other than the proctor. If students near you
are talking, move to a different seat. If someone
other than the proctor talks to you during a test,
don’t respond. Not only is talking a distraction, but
the proctor may think that you are cheating. Don’t
take the chance. Focus on the test, and nothing else.
Standardized Test Practice
619