Section Preview/
Specific Expectations
In this section, you will
explain corrosion as an
electrochemical process,
and describe some
techniques used to
prevent corrosion
evaluate the environmental
and social impact of some
common cells, including the
hydrogen fuel cell in electric
explain how electrolytic
processes are used in the
industrial production of
chlorine, and how this
element is used in the
purification of water
research and assess some
environmental, health, and
safety issues connected to
communicate your understanding of the following
terms: corrosion, galvanizing,
sacrificial anode, cathodic
protection, fuel cell,
chlor-alkali process
Issues Involving Electrochemistry
You have probably seen many examples of rusty objects, such as the one
in Figure 11.26. However, you may not realize that rusting costs many
billions of dollars per year in prevention, maintenance, and replacement
costs. You will now learn more about rusting and about other issues
involving electrochemistry.
Figure 11.26
Rust is a common sight on iron objects.
Rusting is an example of corrosion, which is a spontaneous redox
reaction of materials with substances in their environment. Figure 11.27
shows an example of the hazards that result from corrosion.
Many metals are fairly easily oxidized. The atmosphere contains a
powerful oxidizing agent: oxygen. Because metals are constantly in
contact with oxygen, they are vulnerable to corrosion. In fact, the term
“corrosion” is sometimes defined as the oxidation of metals exposed
to the environment. In North America, about 20% to 25% of iron and
steel production is used to replace objects that have been damaged or
destroyed by corrosion. However, not all corrosion is harmful. For
example, the green layer formed by the corrosion of a copper roof is
considered attractive by many people.
Figure 11.27 On April 28, 1988,
this Aloha Airlines aircraft was
flying at an altitude of 7300 m when
a large part of the upper fuselage
ripped off. This accident was
caused by undetected corrosion
damage. The pilot showed tremendous skill in landing the plane.
546 MHR • Unit 5 Electrochemistry
Rust is a hydrated iron(III) oxide, Fe2O3 • xH2O. The electrochemical
formation of rust occurs in small galvanic cells on the surface of a piece
of iron, as shown in Figure 11.28. In each small cell, iron acts as the
anode. The cathode is inert, and may be an impurity that exists in the iron
or is deposited onto it. For example, the cathode could be a piece of soot
that has been deposited onto the iron surface from the air.
Water, in the form of rain, is needed for rusting to occur. Carbon
dioxide in the air dissolves in water to form carbonic acid, H2CO3(aq) .
This weak acid partially dissociates into ions. Thus, the carbonic acid is
an electrolyte for the corrosion process. Other electrolytes, such as road
salt, may also be involved. The circuit is completed by the iron itself,
which conducts electrons from the anode to the cathode.
Figure 11.28 The rusting of iron involves the reaction of iron, oxygen, and water in a
naturally occurring galvanic cell on the exposed surface of the metal. There may be many
of these small cells on the surface of the same piece of iron.
The rusting process is complex, and the equations may be written in
various ways. A simplified description of the half-reactions and the
overall cell reaction is given here.
Oxidation half-reaction (occurs at the anode):
Fe(s) → Fe2+(aq) + 2e−
Reduction half-reaction (occurs at the cathode):
O2(g) + 2H2O() + 4e− → 4OH−(aq)
Multiply the oxidation half-reaction by two and add the half-reactions
to obtain the overall cell reaction.
2Fe(s) + O2(g) + 2H2O() → 2Fe2+(aq) + 4OH−(aq)
There is no barrier in the cell, so nothing stops the dissolved Fe2+ and
OH− ions from mixing. The iron(II) ions produced at the anode and the
hydroxide ions produced at the cathode react to form a precipitate of
iron(II) hydroxide, Fe(OH)2 . Therefore, the overall cell reaction could
be written as follows.
2Fe(s) + O2(g) + 2H2O() → 2Fe(OH)2(s)
The iron(II) hydroxide undergoes further oxidation by reaction with the
oxygen in the air to form iron(III) hydroxide.
4Fe(OH)2(s) + O2(g) + 2H2O() → 4Fe(OH)3(s)
Chapter 11 Cells and Batteries • MHR
Iron(III) hydroxide readily breaks down to form hydrated iron(III) oxide,
Fe2O3 • xH2O, more commonly known as rust. As noted earlier, the “x”
signifies a variable number of water molecules per formula unit.
2Fe(OH)3(s) → Fe2O3 • 3H2O(s)
Both Fe(OH)3 and Fe2O3 • xH2O are reddish-brown, or “rust coloured.”
A rust deposit may contain a mixture of these compounds.
Fortunately, not all metals corrode to the same extent as iron. Many
metals do corrode in air to form a surface coating of metal oxide.
However, in many cases, the oxide layer adheres, or sticks firmly, to the
metal surface. This layer protects the metal from further corrosion. For
example, aluminum, chromium, and magnesium are readily oxidized in
air to form their oxides, Al2O3 , Cr2O3 , and MgO. Unless the oxide layer is
broken by a cut or a scratch, the layer prevents further corrosion. In contrast, rust easily flakes off from the surface of an iron object and provides
little protection against further corrosion.
Corrosion Prevention
Corrosion, and especially the corrosion of iron, can be very destructive.
For this reason, a great deal of effort goes into corrosion prevention. The
simplest method of preventing corrosion is to paint an iron object. The
protective coating of paint prevents air and water from reaching the metal
surface. Other effective protective layers include grease, oil, plastic, or a
metal that is more resistant to corrosion than iron. For example, a layer
of chromium protects bumpers and metal trim on cars. An enamel coating
is often used to protect metal plates, pots, and pans. Enamel is a shiny,
hard, and very unreactive type of glass that can be melted onto a metal
surface. A protective layer is effective as long as it completely covers the
iron object. If a hole or scratch breaks the layer, the metal underneath
can corrode.
It is also possible to protect iron against corrosion by forming an alloy
with a different metal. Stainless steel is an alloy of iron that contains at
least 10% chromium, by mass, in addition to small quantities of carbon
and occasionally metals such as nickel. Stainless steel is much more
resistant to corrosion than pure iron. Therefore, stainless steel is often
used for cutlery, taps, and various other applications where rust-resistance
is important. However, chromium is much more expensive than iron. As
a result, stainless steel is too expensive for use in large-scale applications,
such as building bridges.
Galvanizing is a process in which iron is covered with a protective
layer of zinc. Galvanized iron is often used to make metal buckets and
chain-link fences. Galvanizing protects iron in two ways. First, the zinc
acts as a protective layer. If this layer is broken, the iron is exposed to
air and water. When this happens, however, the iron is still protected.
Zinc is more easily oxidized than iron. Therefore, zinc, not iron, becomes
the anode in the galvanic cell. The zinc metal is oxidized to zinc ions. In
this situation, zinc is known as a sacrificial anode, because it is destroyed
(sacrificed) to protect the iron. Iron acts as the cathode when zinc is present. Thus, iron does not undergo oxidation until all the zinc has reacted.
548 MHR • Unit 5 Electrochemistry
Cathodic protection is another method of preventing rusting, as
shown in Figure 11.29. As in galvanizing, a more reactive metal is
attached to the iron object. This reactive metal acts as a sacrificial anode,
and the iron becomes the cathode of a galvanic cell. Unlike galvanizing,
the metal used in cathodic protection does not completely cover the iron.
Because the sacrificial anode is slowly destroyed by oxidation, it must be
replaced periodically.
If iron is covered with a protective layer of a metal that is less reactive
than iron, there can be unfortunate results. A “tin” can is actually a steel
can coated with a thin layer of tin. While the tin layer remains intact, it
provides effective protection against rusting. If the tin layer is broken or
scratched, however, the iron in the steel corrodes faster in contact with
the tin than the iron would on its own. Since tin is less reactive than
iron, tin acts as a cathode in each galvanic cell on the surface of the can.
Therefore, the tin provides a large area of available cathodes for the small
galvanic cells involved in the rusting process. Iron acts as the anode of
each cell, which is its normal role when rusting.
Sometimes, the rusting of iron is promoted accidentally. For example,
by connecting an iron pipe to a copper pipe in a plumbing system, an
inexperienced plumber could accidentally speed up the corrosion of the
iron pipe. Copper is less reactive than iron. Therefore, copper acts as the
cathode and iron as the anode in numerous small galvanic cells at the
intersection of the two pipes.
Build on your understanding of corrosion by completing the following
practice problems.
Figure 11.29 Magnesium, zinc,
or aluminum blocks are attached to
ships’ hulls, oil and gas pipelines,
underground iron pipes, and
gasoline storage tanks. These
reactive metals provide cathodic
protection by acting as a sacrificial
Practice Problems
25. (a) Use the two half-reactions for the rusting process, and a table
of standard reduction potentials. Determine the standard cell
potential for this reaction.
(b) Do you think that your calculated value is the actual cell potential
for each of the small galvanic cells on the surface of a rusting iron
object? Explain.
26. Explain why aluminum provides cathodic protection to an iron
27. In the year 2000, Transport Canada reported that thousands of cars
sold in the Atlantic Provinces between 1989 and 1999 had corroded
engine cradle mounts. Failure of these mounts can cause the steering
shaft to separate from the car. The manufacturer recalled the cars
so that repairs could be made, where necessary. The same cars
were sold across the country. Why do you think that the corrosion
problems showed up in the Atlantic Provinces?
28. (a) Use a table of standard reduction potentials to determine whether
elemental oxygen, O2(g), is a better oxidizing agent under acidic
conditions or basic conditions.
(b) From your answer to part (a), do you think that acid rain promotes
or helps prevent the rusting of iron?
Chapter 11 Cells and Batteries • MHR
Automobile Engines
The internal combustion engine found in most automobiles uses
gasoline as a fuel. Unfortunately, this type of engine produces pollutants,
such as carbon dioxide (CO2 ), nitrogen oxides (NOx ), and volatile
organic compounds (VOCs). These pollutants contribute to health and
environmental problems, such as smog and the greenhouse effect. In
addition, the internal combustion engine is very inefficient. It converts
only about 25% of the chemical energy of the fuel into the kinetic
energy of the car. Electric cars, such as those shown in Figure 11.30,
may provide a more efficient and less harmful alternative.
Figure 11.30 Many electric cars
are currently under development.
In fact, the idea of electric cars is
not new. In the early days of the
automobile, electric cars were more
common than gasoline-powered
cars. The production of electric cars
peaked in 1912, but then completely
stopped in the 1930s.
Manufacturers and researchers have attempted to power
electric cars with rechargeable batteries, such as modified lead-acid
and nickel-cadmium batteries. However, rechargeable batteries run
down fairly quickly. The distance driven before recharging a battery
may be 250 km or less. The battery must then be recharged from an
external electrical source. Recharging the lead-acid battery of an electric
car takes several hours. Cars based on a version of the nickel-cadmium
battery can be recharged in only fifteen minutes. However, recharging
the batteries of an electric car is still inconvenient.
A new type of power supply for electric cars eliminates the need for
recharging. A fuel cell is a battery that produces electricity while reactants
are supplied continuously from an external source. Because reactants
continuously flow into the cell, a fuel cell is also known as a flow battery.
Unlike the fuel supply of a more conventional battery, the fuel supply in a
fuel cell is unlimited. As in the combustion of gasoline in a conventional
engine, the overall reaction in a fuel cell is the oxidation of a fuel by oxygen.
The space shuttle uses a fuel cell as a source of energy. This cell
depends on the oxidation of hydrogen by oxygen to form water. The
fuel cell operates under basic conditions, so it is sometimes referred to
as an alkaline fuel cell. Figure 11.31, on the next page, shows the design
of the cell. The half-reactions and the overall reaction are as follows.
Reduction (occurs at the cathode): O2(g) + 2H2O() + 4e− → 4OH−(aq)
Oxidation (occurs at the anode): H2(g) + 2OH−(aq) → 2H2O() + 2e−
Multiply the oxidation half-reaction by 2, add the two half-reactions, and
simplify to obtain the overall cell reaction.
2H2(g) + O2(g) → 2H2O()
Notice that the overall equation is the same as the equation for the
combustion of hydrogen. The combustion of hydrogen is an exothermic
reaction. In the fuel cell, this reaction produces energy in the form of
electricity, rather than heat.
550 MHR • Unit 5 Electrochemistry
inert carbon electrode
containing Ni
inert carbon electrode
containing Ni and NiO
hot KOH solution
Figure 11.31 In an alkaline fuel cell, the half-reactions do not include solid conductors of
electrons. Therefore, the cell has two inert electrodes.
The hydrogen fuel cell produces water vapour, which does not contribute
to smog formation or to the greenhouse effect. This product makes the
hydrogen fuel cell an attractive energy source for cars. Also, the hydrogen
fuel cell is much more efficient than the internal combustion engine. A
hydrogen fuel cell converts about 80% of the chemical energy of the fuel
into the kinetic energy of the car.
There is a possible problem with the hydrogen fuel cell. The cell
requires hydrogen fuel. Unfortunately, uncombined hydrogen is not found
naturally on Earth. Most hydrogen is produced from hydrocarbon fuels,
such as petroleum or methane. These manufacturing processes may
contribute significantly to pollution problems. However, hydrogen can
also be produced by the electrolysis of water. If a source such as solar
energy or hydroelectricity is used to power the electrolysis, the overall
quantity of pollution is low.
The following practice problems will allow you to test your
understanding of fuel cells.
Practice Problems
29. Calculate E˚cell for a hydrogen fuel cell.
30. In one type of fuel cell, methane is oxidized by oxygen to form
carbon dioxide and water.
(a) Write the equation for the overall cell reaction.
(b) Write the two half-reactions, assuming acidic conditions.
31. Reactions that occur in fuel cells can be thought of as being
“flameless combustion reactions.” Explain why.
32. If a hydrogen fuel cell produces an electric current of 0.600 A for
120 min, what mass of hydrogen is consumed by the cell?
Chapter 11 Cells and Batteries • MHR