Larry Brown
Tom Holme
www.cengage.com/chemistry/brown
Chapter 13
Electrochemistry
Jacqueline Bennett • SUNY Oneonta
Chapter Objectives
• Describe at least three types of corrosion and identify
chemical reactions responsible for corrosion.
• Define oxidation and reduction.
• Write and balance half-reactions for simple redox processes.
• Describe the differences between galvanic and electrolytic
cells.
• Use standard reduction potentials to calculate cell potentials
under both standard and nonstandard conditions.
2
Chapter Objectives
• Use standard reduction potentials to predict the spontaneous
direction of a redox reaction.
• Calculate the amount of metal plated, the amount of current
needed, or the time required for an electrolysis process.
• Distinguish between primary and secondary batteries.
• Describe the chemistry of some common battery types and
explain why each type of battery is suitable for a particular
application.
• Describe at least three common techniques for preventing
corrosion.
3
Corrosion
• Corrosion is the degradation of metals by chemical reactions
with the environment.
• Uniform corrosion occurs evenly over a large portion of the
surface area of a metal.
• Galvanic corrosion occurs when two different metals
contact each other in the presence of an appropriate
electrolyte.
• Crevice corrosion occurs when two pieces of metal touch
each other, leaving a small gap or crevice between the
metals.
4
Corrosion
• Different metals corrode differently.
• Aluminum has a greater tendency to corrode than iron, but
corrosion of aluminum is not problematic compared to
iron.
• The aluminum oxide corrosion product forms a
protective layer on the surface of aluminum metal.
• The iron oxide corrosion product flakes off the surface
of iron, exposing fresh iron to corrosion.
5
Corrosion
• Corrosion occurs in a variety of forms. The chain shows
uniform corrosion. The grill cover shows crevice corrosion
where the handle is attached.
6
Oxidation-Reduction Reactions and Galvanic Cells
• Special conditions must be present before iron reacts with
oxygen to form iron(III) oxide.
• Rust formation is a slow process, so the basics of
electrochemistry must be investigated using more easily
observed reactions.
• Reactions that transfer electrons between reactants are
known as oxidation-reduction or redox reactions.
• Oxidation is the loss of electrons from some chemical
species.
• Reduction is the gain of electrons to some chemical
species.
7
Oxidation-Reduction and Half-Reactions
• For an oxidation-reduction reaction to occur, one reactant
must be oxidized and one reactant must be reduced.
• Oxidation cannot occur without reduction.
• When copper wire is placed in a silver nitrate solution, a redox
reaction occurs.
• A reaction is observed to occur because the solution
changes color and crystals form on the surface of the
copper wire.
8
Oxidation-Reduction and Half-Reactions
• When a clean copper wire is placed into a colorless solution
of silver nitrate, it is quickly apparent that a chemical reaction
occurs.
9
Oxidation-Reduction and Half-Reactions
• The solution’s blue color is indicative of Cu2+ ions in solution.
• Cu2+ is formed when a copper atom loses two electrons.
• The copper metal is oxidized.
Cu(s) 
 Cu2+ (aq) + 2e
• The crystals forming on the surface of the copper wire are
silver metal.
• Silver is formed when a silver cation gains an electron.
• The silver cation is reduced.

Ag (aq) + 1e 
 Ag(s)
+
10
Oxidation-Reduction and Half-Reactions
• For the reaction between silver cation and copper metal, two
half-reactions are written.
• One for the oxidation of copper and one for the reduction
of silver.
• Neither half-reaction can occur without the other.
• The half-reactions as written indicate that Ag+ only accepts
one electron whereas Cu loses two electrons.
• The electron transfer must balance, so the reduction halfreaction is multiplied by 2.
Cu(s) 
 Cu 2+ (aq) + 2e
2Ag + (aq) + 2e  
 2Ag(s)
11
Oxidation-Reduction and Half-Reactions
• Add the two half-reactions together, the electrons cancel out,
leaving the net ionic equation for the redox reaction.
Cu(s) 
 Cu 2+ (aq) + 2e 
2Ag + (aq) + 2e  
 2Ag(s)
Cu(s) + 2Ag + (aq) 
 Cu 2+ (aq) + 2Ag(s)
12
Oxidation-Reduction and Half-Reactions
• The species undergoing oxidation is referred to as a reducing
agent.
• The Cu was oxidized and is the reducing agent.
• The Cu facilitated the reduction of Ag+ by losing electrons.
• The species undergoing reduction is referred to as an
oxidizing agent.
• The Ag+ was reduced and is the oxidizing agent.
• The Ag+ facilitated the oxidation of Cu by gaining
electrons.
13
Building a Galvanic Cell
• A galvanic cell is any electrochemical cell in which a
spontaneous chemical reaction can be used to generate an
electric current.
• The name electrochemistry comes from the observation of
electric currents in galvanic cells.
• To harness electricity from a galvanic cell, each half-reaction
is prepared separately in half-cells.
• Cu metal immersed in Cu2+ solution is one half-cell.
• Ag metal immersed in Ag+ solution is the second half-cell.
14
Building a Galvanic Cell
• Current flows by the migration of ions in solution.
• To transfer current between the half-cells, a salt bridge is
used.
• The salt bridge contains a strong electrolyte that allows
either cations or anions to migrate into the solution where
they are needed to maintain charge neutrality.
• A metal wire cannot transport ions and cannot be used.
15
Building a Galvanic Cell
• For a salt bridge composed of NH4Cl:
• NH4+ will flow into the Ag+ beaker to offset the removal of
Ag+ from solution.
• Cl– will flow into the Cu2+ beaker to offset the production of
Cu2+ in solution.
• The circuit is completed by connecting wires to each metal
strip.
• A voltage potential of 0.46 V will be measured for the
described cell.
16
Building a Galvanic Cell
• A salt bridge is crucial to a galvanic cell. The salt bridge
allows ions to flow between each half-cell, completing the
circuit.
17
Terminology for Galvanic Cells
• Electrodes are the electrically conducting sites at which either
oxidation or reduction occurs.
• The electrode where oxidation occurs is the anode.
• The electrode where reduction occurs is the cathode.
• Cell notation - a shorthand notation for the specific chemistry
of an electrochemical cell.
• Cell notation lists the metals and ions involved in the
reaction.
• A vertical line, |, denotes a phase boundary.
• A double vertical line, ||, denotes a salt bridge.
• The anode is written on the left, the cathode on the right.
18
Terminology for Galvanic Cells
• General form of cell notation
anode | anode electrolyte || cathode electrolyte | cathode
• For the previous example of copper and silver
Cu(s)| Cu2+ (aq) (1 M)|| Ag+ (aq) (1 M)| Ag
• The electrolyte concentration is also given.
• An electrochemical cell is at its standard state when the electrolyte
concentrations are 1 M.
• For half-cells that generate or consume a gas, a partial pressure of
1 atm is required for the standard state.
19
Atomic Perspective on Galvanic Cells
• Before half-cells are connected by a salt bridge, a small build
up of charge occurs for each half-cell at the interface between
the electrode and the electrolyte.
• At the anode, some oxidation occurs and cations dissolve
into solution, leaving a negative charge on the anode.
• At the cathode, some reduction occurs and cations are
removed from solution, leaving a positive charge on the
cathode.
20
Atomic Perspective on Galvanic Cells
• An equilibrium can be described for each half-cell, the halfreaction equilibrium.
• Not an oxidation-reduction equilibrium.
• The build up of charge on the electrode means there is
potential for electrical work.
• This potential is the cell potential, or electromotive force
(EMF).
• EMF is related to the maximum work obtainable from an
electrochemical cell.
• wmax = qE
• q is the charge, E is the cell potential.
21
Atomic Perspective on Galvanic Cells
• Without a salt bridge to
close the circuit, local
charges build up around
both electrodes. Neither
electrode reaction can
proceed to any significant
extent, so no cell voltage
can be measured.
22
Galvanic Corrosion and Uniform Corrosion
• Metals in contact with a solution establish an oxidation halfreaction equilibrium.
• If the solution contains a substance that can undergo
reduction, a redox reaction may occur.
• For two metals in contact, such as a tin-plated steel can,
exposure to air and moisture results in rapid corrosion.
• The half-reaction equilibrium for the tin facilitates the
process by which iron is oxidized.
• This is an example of galvanic corrosion.
23
Galvanic Corrosion and Uniform Corrosion
• A “tin can” is usually tin-plated steel. If the tin coating is
scratched to expose the underlying steel, iron in the steel will
corrode rapidly.
24
Galvanic Corrosion and Uniform Corrosion
• Metal not in contact with another metal can also corrode.
• A nonmetal is involved in the second half-cell.
• For the corrosion of iron, iron is one half-cell and oxygen
dissolved in water is the second half-cell.
• The electrode for the oxygen half-cell is the iron itself.
• Dissolved salts facilitate the corrosion reaction.
• This is an example of uniform corrosion.
25
Cell Potentials
• The relative corrosivities of various plated steels can be
expressed as cell potential.
• A voltmeter measures the size of the electrical potential and
also its polarity - the locations of the negative charge
(negative pole) and the positive charge (positive pole).
• An electric potential has a fixed polarity and voltage.
• Reversing the poles of a battery with respect to a
voltmeter changes the sign on the measured voltage but
does not influence the electrochemical reaction in the
battery.
26
Measuring Cell Potential
• When a voltmeter is connected to the previously described
copper/silver cell, a potential of 0.462 V is measured.
• Connecting the copper half-cell to a reducing iron(III)/iron(II)
half-cell, a cell potential of 0.434 V is measured.
• Connecting the iron(III)/iron(II) half-cell to the silver half-cell
results in a cell potential of 0.028.
• For the three cell potentials measured, the fact that 0.462 V =
0.434 V + 0.028 V suggests two things:
• The behavior of cell potentials is akin to state functions.
• If a specific standard electrode is chosen, comparison to
all other electrodes will result in a practical system for
determining cell potential.
27
Measuring Cell Potential
• Measurement of standard cell voltages for various
combinations of the same half-reactions suggests that a
characteristic potential can be associated with a particular
half-reaction.
28
Measuring Cell Potential
• The standard hydrogen electrode
(SHE) is the choice for the standard
component in cell potential
measurements.
• The cell is constructed of a platinum
wire or foil as the electrode.
• The electrode is immersed in a
1 M HCl solution through which
H2 gas with a pressure of 1 atm
is bubbled.
• The SHE has been chosen as the
reference point for the scale of
standard reduction potentials, and
assigned a potential of exactly zero
volts.
29
Measuring Cell Potential
• For the Standard Hydrogen Electrode
• The half-reaction for the SHE is: 2 H+(aq) + 2 e–  H2(g).
• The half-cell notation is: Pt(s) | H2(g, 1 atm) | H+ (1 M).
• The half-cell is assigned a potential of exactly zero volts.
• The cell potential is attributed to the other half-reaction.
30
Measuring Cell Potential
• For some galvanic cells, the SHE acts as the anode and for
other galvanic cells, the SHE acts as the cathode.
• The anode is the site of oxidation, releasing electrons and
creating a negatively charged electrode.
• If the anode is connected to the positive terminal of the
voltmeter, a negative potential is measured.
• The cathode is the side of reduction, consuming electrons
and creating a positively charged electrode.
• If the cathode is connected to the positive terminal of the
voltmeter, a positive potential is measured.
31
Measuring Cell Potential
• When the SHE is always connected to the positive terminal,
the sign of the potential tells us the direction of the redox
reaction.
• When the potential is negative, the SHE is the anode, and
H2 is oxidized to H+(aq).
• When the potential is positive, the SHE is the cathode,
and H+(aq) is reduced to H2.
32
Measuring Cell Potential
• Just like a commercial battery, a galvanic cell has a fixed polarity.
Electrons flow through the external circuit from the anode to the
cathode. Reversing the leads of the voltmeter changes the sign of
the reading, but does not influence the flow of current.
33
Standard Reduction Potentials
• To compare the oxidation-reduction trends of species used in
electrochemistry, all half-cell potentials are written as
reductions.
• A table of standard reduction potentials lists the potential
of any half-reaction when connected to a SHE.
• All materials are 1 M for aqueous species and 1 atm
partial pressure for gases.
34
Standard Reduction Potentials
• Standard reduction potentials for several half-reactions
involved in the cells discussed in the text.
35
Standard Reduction Potentials
• Although the half-reactions are listed as reductions in the
table, one half-reaction in any electrochemical cell must be an
oxidation and, therefore, reversed from what appears in the
table.
• The cell potential sign must be changed when writing the
half-reaction as an oxidation.
• Some half-reactions have positive potentials, whereas others
have negative potentials.
36
Standard Reduction Potentials
• All potentials are measured with a SHE connected to the
negative terminal.
• If the voltage is positive, the SHE is the anode, the
oxidation site.
• A positive standard reduction potential means the halfreaction proceeds as written (reduction occurs).
• If the voltage is negative, the SHE is the cathode, the
reduction site.
• A negative standard reduction potential means the halfreaction proceeds as an oxidation.
37
Standard Reduction Potentials
• The tendency for the chemicals involved in a half-reaction to
be an oxidation or reduction depends on the value of the
reduction potential.
• A large, positive value for the standard reduction potential
implies the substance is reduced readily and a good
oxidizing agent.
• A large, negative value for the standard reduction potential
implies the substance is oxidized readily and a good
reducing agent.
38
Standard Reduction Potentials
• For a galvanic cell, the half-reaction with the more positive
reduction potential will be the cathode.
• The half-reaction with the more negative reduction
potential will be the anode.
• The standard reduction potential for any pair of half-reactions,
E˚cell, is calculated from the standard reduction potentials for
the cathode and anode.
o
cell
E
E
o
red
E
o
ox
• E˚red is the standard reduction potential for the cathode
and E˚ox is the standard reduction potential for the anode.
39
Standard Reduction Potentials
• For standard reduction potentials arranged horizontally, the
anode and cathode for a galvanic cell can easily be
determined. The reduction potential furthest to the left is the
anode.
40
Example Problem 13.1
• Using standard reduction potentials, identify the anode and
the cathode and determine the cell potential for a galvanic cell
composed of copper and iron. Assume standard conditions.
• Confirm that the potential of the following galvanic cell is
0.462 V:
2+
+
Cu(s) | Cu (1 M) || Ag (1 M) | Ag(s)
41
Nonstandard Conditions
• The cell potential at nonstandard conditions is calculated
using the Nernst equation.
RT
EE 
ln Q
nF
o
• Q is the reaction quotient, F is the Faraday constant, and n
is the number of electrons transferred in the reaction.
• F = 96,485 J V-1 mol-1 or 96,485 C mol-1

42
Example Problem 13.2
• Assume that you have a cell that has an iron(II) concentration
of 0.015 M and an H+ concentration of 1.0  10-3 M. The cell
temperature is 38ºC, and the pressure of hydrogen gas is
maintained at 0.04 atm. What would the cell potential be
under these conditions?
43
Cell Potentials and Free Energy
• Corrosion is a spontaneous process and has a negative
Gibbs free energy change.
• The Gibbs free energy change for an electrochemical reaction
can be calculated from the standard reduction potential.
Go  nFE o
• n is the number of electrons transferred and F is Faraday’s
constant.

• The minus sign in required because a galvanic cell has a
positive cell potential, spontaneously generates electrical
work, and thus must have a negative G value.
44
Example Problem 13.3
• Suppose that we wish to study the possible galvanic corrosion
between zinc and chromium, so we set up the following cell:
Cr(s) | Cr 2+ (aq) || Zn2+ (aq) | Zn(s)
• What is the chemical reaction that takes place and what is the
standard free energy change for that reaction?
45
Equilibrium Constants
• The cell potential can be used to calculate the equilibrium
constant for an electrochemical reaction.
RT
E 
ln K
nF
o
• n is the number of electrons transferred, R is the universal
gas law constant, and F is Faraday’s constant.

46
Equilibrium Constants
• The relationship between the cell potential and the equilibrium
constant can be re-written in terms of the common (base 10)
log.
2.303RT
E 
log K
nF
o
• The equation can be simplified for reaction carried out at

standard temperature,
25°C (298 K).
0.0592 V
E 
log K
n
o
47
Equilibrium Constants
• The equilibrium
constant increases
as the cell potential
and number of
electrons transferred
increases.
• The different lines
correspond to
reactions
involving 1, 2, or
3 electrons
transferred.
48
Batteries
• A battery is a cell or series of cells that generate an electrical
current.
• Batteries are the means by which we harness the
electrical work of a galvanic cell and use it productively.
49
Primary Cells
• Single-use batteries that cannot be recharged are primary
cells, or primary batteries.
• The most prevalent type of primary cell is the alkaline
battery.
• An alkaline battery has a zinc electrode at which oxidation
occurs.

Zn(s) + 2OH (aq) 
 Zn(OH)2 (s) + 2e

• The cathode is derived from manganese(IV) oxide.
2MnO2 (s) + H2O(l ) + 2e 
 Mn2O3 (s) + 2OH (aq)
50
Primary Cells
Zn(s) + 2 MnO2 (s) + H2O(l ) 
 Zn(OH)2 (s) + Mn2O3 (s)
• The chemistry of an
alkaline dry cell battery.
The net reaction is shown
above.
• The alkaline battery is
termed a dry cell
because the KOH
electrolyte is in the form
of a paste or gel.
51
Primary Cells
• Mercury batteries are another type of primary cell and are quite
small. They are used for medical devices like pacemakers.
• Zinc is the anode.
Zn(s) + 2OH (aq) 
 Zn(OH)2 (s) + 2e
• Mercury(II) oxide is the cathode.
HgO(s) + H2O(l ) + 2e 
 Hg(l ) + 2OH (aq)
• The mercury battery
(also called a zincmercuric oxide cell) has
a voltage output that is
extremely stable over
long times.
52
Primary Cells
• Zinc-air batteries are also primary cells.
• Zinc is the anode.
Zn(s) + 2OH (aq) 
 Zn(OH)2 (s) + 2e
• Oxygen reacts at the cathode.
1
O2 (g) + H 2O(l ) + 4e 
 2OH  (aq)
2
• In a zinc-air battery, one
of the reactants is
oxygen from the
surrounding air. As a
result, these batteries
can offer a very attractive
energy density.
53
Secondary Cells
• Rechargeable batteries are secondary cell or secondary
batteries.
• Nickel-cadmium or “ni-cad” batteries are an example of
secondary cells.
• The anode for a ni-cad battery is cadmium.
Cd(s) + 2OH (aq) 
 Cd(OH)2 (s) + 2e
• The complex cathode reaction can be represented as
NiO(OH)(s) + H2O(l ) + e 
 Ni(OH)2 (s) + OH (aq)
54
Secondary Cells
• Important design features of a nickel-cadmium battery are
shown to the left.
55
Secondary Cells
• Nickel-metal-hydride batteries have become popular as
rechargeable cells. The design is quite similar to the Ni-Cd
cell, but nickel-metal-hydride cells are less prone to memory
effects.
56
Secondary Cells
• Nickel-metal-hydride batteries are secondary cells.
• The anode for a nickel-metal-hydride battery is M, some
metal or metal alloy.
MH(s) + OH (aq) 
 M(s) + H2O(l ) + e
• The complex cathode reaction can be represented as
NiO(OH)(s) + H2O(l ) + e 
 Ni(OH)2 (s) + OH (aq)
57
Secondary Cells
• The lead-acid storage battery found in cars is a secondary
cell.
• The anode for a lead-acid battery is lead metal.
Pb(s) + HSO4 (aq) 
 PbSO4 (s) + H+ (aq) + 2e
• The cathode for a lead-acid battery is lead oxide.
PbO2 (s) + 3H+ (aq) + HSO4 (aq) + 2e 
 PbSO4 (s) + 2H2O(l )
• The lead-acid storage
battery consists of Pb
anodes alternating with
PbO2 cathodes, all
immersed in sulfuric acid.
58
Fuel Cells
• A fuel cell is a voltaic cell in which the reactants can be
supplied continuously and the products of the cell reaction are
continuously removed.
• Most common type is based on the reaction of hydrogen and
oxygen to produce water.
2H2  O2  2H2O
• Oxygen is reduced at the cathode.


O2  4H  4e  2H2O
• Hydrogen is oxidized at the anode.
H2  2H  2e
59
Limitations of Batteries
• Corrosion is a major cause for the loss of performance in
batteries.
• Protective plating of materials used in batteries is an
attempt to limit the performance-diminishing effects of
corrosion on batteries.
60
Electrolysis
• Electrolysis is the process of passing an electric current
through an ionic solution or molten salt to produce a chemical
reaction.
• Electrolytic cells are divided into two categories based on the
nature of the electrodes used.
• Passive electrolysis: the electrodes are chemically inert
materials that simply provide a path for electrons.
• Active electrolysis: the electrodes are part of the
electrolytic reaction.
61
Electrolysis and Polarity
• Electrolysis changes the polarity of the electrodes in a
system.
• For reduction, electrons are forced to the cathode. The
cathode becomes the negative electrode.
• For oxidation, electrons are pulled from the anode. The
anode becomes the positive electrode.
• In electrolysis, an
external source of current
drives a redox reaction
that would otherwise not
be spontaneous. The
flow of ions through the
solution completes the
circuit.
62
Passive Electrolysis in Refining Aluminum
• Electrolysis provides the means to overcome the
nonspontaneous reaction to separate aluminum from its
oxide.
• The Hall-Heroult refining process uses carbon electrodes
as inert sites for passive electrolysis.
• The Hall-Heroult process
involves the
electrolytic refining of
aluminum from Al2O3 to
produce aluminum metal
and oxygen gas.
63
Active Electrolysis and Electroplating
• The process of depositing a thin coat of metal on another
metal by using electrolysis is electroplating.
• In some cases, the thin coating is cosmetic, or to provide
some vital functionality for the coated piece, such as
corrosion resistance or desirable conductive properties.
• Silver is plated onto electrical devices because silver is a
good conductor and resistant to corrosion.
• The solution from which silver is plated contains CN–(aq)
ions, which form a complex with Ag+. The need for
uniform coatings makes this an important step.
64
Active Electrolysis and Electroplating
• The object being electroplated is the cathode.
• Anode
Ag(s) + 2CN (aq) 
 Ag(CN)2 (aq) + e
• Cathode
Ag(CN)2 (aq) + e 
 Ag(s) + 2CN (aq)
• Opposite reactions at the anode and cathode are common for
electroplating operations.
• Silver is transferred from the anode to the cathode, coating the
cathode in a thin layer of silver.
• The zero cell potential is not critical since an external current
drives electrolysis.
65
Active Electrolysis and Electroplating
• Barrel plating is often used to apply coatings to small parts.
66
Electrolysis and Stoichiometry
• For electroplating, it can be vitally important to use carefully
controlled amounts of materials.
• Controlling the flow of electrons (current) in an
electroplating operation provides a method to accurately
limit the amount of material deposited.
• Electroplating is often used to prevent galvanic corrosion
in an electrical apparatus in places where different metals
come into contact with one another.
67
Current and Charge
• When current is measured in an electric circuit, the
observation is the flow of charge for a period of time.
• The unit of current, the ampere (A), is defined as one
coulomb per second: 1 A = 1 C s-1.
• If a known current flows through a circuit for a known time,
the charge can be easily calculated.
Charge = current  time
Q I t
68
Current and Charge
• Using Faraday’s constant, F = 96,485 C mol-1 and the
calculated charge, the number of moles of electrons that pass
through the circuit can be calculated.
• If the number of electrons required to reduce each metal
cation is known, the number of moles of material plated can
be calculated.
• Electricity use is often measured in terms of power. The SI
unit for power is the watt (1 watt = 1 J s-1)
• Electrical utilities normally determine consumption in kilowatthours, kWh (1 kWh = 3.60  106 J)
69
Example Problem 13.4
• In a process called flash electroplating, a current of
2.5 x 103 A passes through an electrolytic cell for
5.00 minutes. How many moles of electrons are
driven through the cell?
70
Example Problem 13.5
• Suppose that a batch of parts is plated with copper in an
electrolytic bath running at 0.15 V and 15.0 A for exactly 2
hours. What is the energy cost of this process if the electric
utility charges the company $0.0500 per kWh?
71
Calculations Using Masses of Substances in Electrolysis
• A knowledge of current, how long the current flows,
stoichiometry, and the number of electrons required to reduce
a metal cation are used to answer the following questions.
• How much material is plated given a specific current for an
allotted time or electrical energy expenditure?
• How long must a given current to pass through the cell to
yield a desired mass of plated material?
72
Example Problem 13.6
• An electrolysis cell that deposits gold (from Au+(aq)) operates
for 15.0 minutes at a current of 2.30 A. What mass of gold is
deposited?
73
Example Problem 13.7
• Suppose that you have a part that requires tin coating. You’ve
calculated that you need to deposit 3.60 g of tin to achieve an
adequate coating. If your electrolysis cell (using Sn2+) runs at
2.00 A, how long must you operate the cell to obtain the
desired coating?
74
Corrosion Prevention
• The following observations can be used to reduce corrosion in
a number of ways.
• Corrosion is a pervasive reaction, with a large, negative
free energy change.
• It is possible to predict what materials will corrode and use
this information to protect a material such as iron.
• Some materials, like aluminum, corrode readily, but the
product, in this case Al2O3, forms a protective layer that
eliminates further corrosion.
75
Coatings
• Applying a protective coating to a material is the most
common way of protecting against corrosion.
• A coating can be applied with electroplating or painting.
• The coating protects the underlying material from
exposure to water and oxygen.
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Coatings
• Rust inhibitors can be added to paint to further inhibit
corrosion.
• Many of the common inhibitors contain the following ions:
phosphate, borosilicate, chromate, or phosphosilicate.
• All of these ions, as part of a paint coating, form
compounds with oxidized iron that inhibit further rust
formation.
• This process is called passivation.
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Cathodic Protection
• Some materials are more easily oxidized than iron, which
provides a way to construct galvanic corrosion conditions
intended to protect the iron.
• Mg has a reduction potential more negative than Fe.
• When combined, Mg will oxidize and Fe will be reduced.
• Magnesium can be used to prevent iron corrosion.
• The piece of magnesium is called a sacrificial anode.
• Connecting magnesium to iron forces iron to be the
cathode, preventing iron from oxidizing.
• This process is called cathodic protection.
• The sacrificial anode must be replaced periodically to be
effective.
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Cathodic Protection
• Sacrificial anodes are one effective method of corrosion
prevention. The anode is preferentially oxidized relative to the
protected metal.
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Preventing Corrosion in space
• Corrosion prevention is a concern for NASA.
• Corrosion prevention at the launch site is a major concern.
• Batteries used in the International Space Station (ISS)
must be prevented from corroding in the ISS’s earth-like
atmosphere.
• The corrosivity of a planet’s atmosphere must be known
before a craft can land safely.
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