ENGINEERING CHEMISTRY

CORROSION AND ITS CONTROL

UNIT-3

DEPARTMENT OF CHEMISTRY

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 3. Corrosion and its Control
3.1 Introduction
The deterioration or degradation of a material due to a reaction with its environment
is called as Corrosion. It leads to a change in physical properties of the material.
Corrosion causes to weakening of a material. It is of various types like shattering of a
metal or cracking of a polymer. The Material that undergoes corrosion can either be
a metals or polymer or ceramics or composites. In the field of engineering metals are
generally used as fabrication or construction materials. The metal structures
deteriorate if they are not properly maintained.
3.1.1. Why do metals corrode?
When the environments are chemically unstable metal corrosion takes place. Only
metals that have a positive E° value according to the electrochemical series, like gold,
silver, platinum, etc. Some metals are unstable with their environments and form a
more stable state. There are certain metals that form a protective film on their
surfaces that prevent corrosion. Corrosion can prevented by using metals that form
protective films which are expensive. So, simple methods like painting are used
for corrosion control.
Definition
"Corrosion is the gradual destruction or deterioration of metals or alloys by the
chemical or electrochemical reaction with its environment"
3.2 Causes of Corrosion
Metals occur in two different forms.
1. The native state.
2. The combined state.
3.2.1. Native state
An element is said to exist in the native state in its elementary form. They occur
in an uncombined state which is non–reactive towards the environment. Noble metals
are good examples that exist as such in the earth’s crust. They posses very good
corrosion resistance and are less reactive, elements such as gold, platinum, etc., are
found to be in the native state.
3.2.2. Combined State
An element is said to exist in the combined state when it exists in nature as a
compound. Reactive elements occur in nature in the combined state. Expect noble
metals, all other metals are reactive and react with environment to form a stable
compound as their oxides or sulphides or chlorides or carbonates. They exist in the
form of stable compounds called ores and minerals. A good example is, when iron is
exposed to air that is moist it gets corroded and a layer of reddish brown substance
called as rust is formed on the surface. Rust is a chemically hydrated form of iron (III)
oxide, Fe2O3. xH2O. Copper metal when exposed to moist air gets coated forming a
greenish white powdery substance which is copper carbonate.

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3.3. How does corrosion occurs?
A metal is extracted from its metallic compound or ore. During the extraction an
ore is reduced to its metallic state by applying energy in the form of various processes.
When a metal exists in its pure metallic state it is highly unstable. It is considered to
be in a high energy state or excited state. So when a metal is extracted from their ore
the reverse process begins and form metallic compound which reaches the lower
energy state or it is said to be thermodynamically stable. When a metal is used in
various forms they are exposed to different environments. Here the metallic surface
starts to decay or get converted to a more stable compound.

Corrosion - Oxidation
Metal Metallic Compound + Energy
Metallurgy -Reduction

Corroded metals are thermodynamically more stable than pure metals but due to
corrosion, useful properties like malleability, ductility, hardness, luster and electrical
conductivity are lost.
3.4. Effects of corrosion
Some effects of corrosion are
• It reduces the metal thickness which leads to a loss of mechanical strength and
structural failure.
• Corrosion is hazardous and creates injuries to people arising from structural.
• Cost of the product is reduced due to deterioration of appearance.
• Fluids when transferred get contaminated in the vessels and during flow in pipes.
• Blockage of pipes due to solid corrosion products.
• Expense of equipment replacement is high.
• Plant failure occurs.
3.5. Classification or theories of corrosion
Based on the environment, corrosion is classified into
• Dry or Chemical corrosion.
• Wet or Electrochemical corrosion.
3.5.1 Dry or Chemical Corrosion
Dry corrosion is due to the attack of metal surfaces by the atmospheric gases
such as oxygen, hydrogen sulphide, sulphur dioxide, nitrogen, etc.,
There are 3 main types of dry corrosion.
1. Oxidation corrosion (or) corrosion by oxygen.
2. Corrosion by hydrogen
3. Liquid–metal corrosion.
3.5.1.1. Oxidation corrosion (or) corrosion by oxygen
Oxidation corrosion occurs when there is a direct action of oxygen at low or
high temperature on metals in the absence of moisture. At normal temperature there
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is a slight attack on the metal. But alkalies and alkaline earth metals are rapidly
oxidized at low temperature and almost all metals except the noble metals are oxidized
at high temperatures.
Reactions in oxidation corrosion
2M → 2Mn+ + 2ne- (Loss of electrons)
n/2 02 +2ne- → n02- (Gain of electrons)
Overall reaction
2M + n/2O2 → 2Mn+ + nO2- (Metal oxide)
Mechanism
Oxidation first occurs at the surface of the metal and the resulting metal oxide
forms a barrier which restricts further oxidation. For oxidation to continue the metal
must diffuse outwards through the metal oxide layer or oxygen must diffuse inwards
through the scale to the underlying metal. The outward diffusion of the metal is more
rapid than inward diffusion of oxygen, since the metal ion is much smaller than
oxygen molecule, thus there is higher mobility. (Fig. 3.1.)

Fig.3.1. Mechanism and photographic image of oxidation corrosion

Nature of oxide film
There are three types of oxide film formed on the metal surface
Stable oxide layer
A stable oxide layer is a fine grained structure which gets adsorbed tight on to
the metal surface. Such a layer is highly impervious in nature and stops further
oxygen attack through diffusion. Such a film behaves as a protective coating and no
further corrosion can develop.
Example: Oxides of Cu, Al, Sn, Pb, etc.,
• Unstable oxide layer
Unstable oxide layer is produced on to the surface of the noble metals, which
decomposes back into metal and oxygen.
Metal oxide Metal + Oxygen
Example: Oxides of Pt, Ag, etc.,

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 • Volatile oxide layer
The oxide layer volatilizes or forms a gas and disappears immediately after it is
formed, which causes further corrosion.
Example: Molybdenum oxide (MoO3) is volatile.

Pilling–Bedworth ratio
The Pilling-Bedworth ratio or P-B ratio is the ratio of the volume of the metal
oxide formed to the volume of the metal consumed. On the basis of the Pilling
bedworth ratio a metal can be judged whether it is likely to form a protective oxide
layer.
• Porus (or) Non–protective oxide film
If the volume of the oxide layer formed is less than the volume of the metal
consumed, the oxide layer is considered to be porous and non–protective. A good
example is alkali and alkaline earth metals such as Na, Ca, etc.
Non - Porus (or) Protective oxide film
If the volume of the oxide layer formed is greater than the volume of the metal
consumed, the oxide layer is considered to be non–porous and protective. The oxides
of heavy metals such as Pb, Sn, are good examples.
3.5.1.2. Corrosion by hydrogen
• Hydrogen embrittlement
Hydrogen embrittlement is a process by which various metals become brittle
and fracture on exposure to hydrogen. Hydrogen embrittlement is the introduction of
hydrogen into metals forming cracks in the metal.
When a metal such as Fe react with H2S at ordinary temperature it causes the
formation of atomic hydrogen.
Fe + H2S → FeS + 2H
This atomic hydrogen diffuses into the metal immediately and recombines to
form molecular hydrogen.
H + H → H2 
These molecular hydrogen gas formed penetrates into the voids present and
develops a high pressure which causes cracks and blisters on metal. ( Fig.3.2.)

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Fig.3.2. Mechanism and Photographic image of Hydrogen embrittlement
• Decarburisation (at higher temperature)
When steel is exposed to hydrogen at very high temperatures hydrogen
penetrates into the alloy and combines with carbon to form tiny packets
of methane in the internal surfaces. This methane does not diffuse out of the metal
but gets collects in the voids at high pressure leading to cracks in the steel.
This process leads to decarburization of the steel which results in loss of strength and
ductility of steel.
H 2 ⎯Heat
⎯→
⎯ 2H
When steel is exposed to hydrogen environment atomic hydrogen readily
combines with the carbon present in steel to produce methane gas. (Fig.3.3.)
C + 4H ⎯
⎯→ CH 4 

Fig.3.3. Photographic image of decarburisation

3.5.1.3. Liquid–metal corrosion
Liquid metal corrosion also called as liquid metal embrittlement is caused when
a certain metal undergoes fracture when it is introduced into a liquid metals
environment. For example when liquid gallium is introduced on an aluminium sheet

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liquid metal embrittlement is observed. During liquid metal embrittlement cracking
occurs drastically leading to very high crack growth rates. These effects are seen even
in solid alloys when one of the metals is brought close to its melting point. For
example cadmium when operated at high temperature leads to solid metal
embrittlement. Embrittlement caused due to mercury spills cause corrosion. (Fig.3.4.)

Fig.3.4. Photographic image of Liquid–metal corrosion

3.5.2. WET OR ELECTRO–CHEMICAL CORROSION

Mechanism of wet corrosion
Wet corrosion of a metal takes place through electron transfer which involves
both oxidation and reduction. During oxidation a metal atoms loses electrons and the
surrounding environment undergoes reduction and gains electrons. The metal in
which there is loss of electrons is called as the anode and the other metal or liquid or
gas that gains electrons is called the cathode.
The reaction is as follows
At anode
At the anode oxidation or dissolution of the metal occurs
M⎯
⎯→ M2+ + 2e−
At cathode
At the cathode reduction of the solid or liquid or gas occurs. This depends on the
nature of the environment. There are two types of environment and they are

a. Acidic environment or Hydrogen evolution type corrosion

The metals that posses a anodic potential in the electrochemical series have got
a tendency to get dissolved in acidic solution with the evolution of hydrogen gas (Fig.
3.5.)

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Fig. 3.5. Hydrogen evolution type corrosion
For example when iron metal contacts with non–oxidizing acid like H2SO4
and HCl, H2 evolution occurs.
At anode
Iron undergoes dissolution to give Fe2+ ions with the liberation of electrons.
Fe ⎯
⎯→ Fe2+ + 2e− (Oxidation)
At cathode
The liberated electrons flow from anodic to cathodic part, where H– ions get
reduced to H2.
2H + + 2e− ⎯
⎯→ H 2  (reduction)
b. Neutral environment or Absorption of oxygen or Formation of hydroxide ion
type corrosion
If the corrosive environment is slightly alkaline or neutral. The product formed is
at the cathode that is hydroxide ions are formed at the cathode.
½O2 + 2e–+H2O ⎯ ⎯→ 2OH–
The metal ions from the anode and the non–metallic ions from the cathode diffuse
towards each other through a conducting medium to form a corrosion product
between the anode and the cathode. For example when iron is in contact with an
electrolyte in presence of oxygen, OH– ions are formed. A thin film of iron oxide is
coated on the surface. However, if the oxide film develops crack at the anode created
on the surface while the remaining part acts as cathode. (Fig. 3.6.)

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 Fig. 3.6. Absorption of oxygen type corrosion
At anode
Iron dissolves as Fe2+ with the liberation of electrons.
Fe ⎯⎯→ Fe2+ + 2e− (oxidation)
At cathode
The liberated electrons flow from anodic to cathodic part through metal, where
the electrons are taken up by the dissolved oxygen to form OH– ions.
½O2 + H2O + 2e– ⎯ ⎯→ 2OH–
Thus the net corrosion reaction is
Fe2+ + 2OH − ⎯
⎯→ Fe(OH)2 
If enough O2 is present Fe(OH)2 is easily oxidized to Fe(OH)3, a rust (Fe2O3, H2O).
4Fe(OH)2 + O2 + 2H2O− ⎯
⎯→ 4Fe(OH)3

Table 3.1 Difference between Chemical and Electro chemical Corrosion

Sl. Chemical Corrosion Electrochemical Corrosion
No.
1. It occurs only in dry condition. It occurs in presence of water or
electrolyte.
2. It is due to the attack of the metal It is due to the set up of a large number of
on the environment. cathodic and anodic areas.
3. Even a homogeneous metal surface Heterogeneous surface (or) Bimetallic
gets corroded. contact is the condition.
4. Corrosion products accumulate Corrosion occurs at the anode, while
where corrosion occurs. products formed elsewhere.
5. Chemical corrosion is self– Electrochemical corrosion is continuous
controlled process.
6. It follows adsorption mechanism It follows electrochemical reaction.
Example Example
Formation of mild scale on iron Rusting of iron in moist atmosphere.
surface

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3.5.2.1. Types of electrochemical corrosion
1. Galvanic corrosion
When two different metallic materials are connected galvanic corrosion occurs.
The material that is more anodic will corrode.
The corrosion rate depends strongly on
• The electric conductivity of the environment (Electrolyte or solution)
• The difference in corrosion potential between the galvanically coupled
materials
• The surface ratio between the electrodes.
For example when stainless steel is joined with steel, stainless tubes and
carbon-steel tube plates in a heat exchanger, the carbon steel suffers from galvanic
corrosion attack. When stainless steel is combined with graphite, the attack will be
on steel. Graphite gaskets should therefore be avoided.
Here, the metal with more negative electrode potential according to the
electrochemical series acts as anode and the metal with less negative electrode
potential acts as cathode.(Fig.3.7.)

Fig.3.7. Mechanism and Photographic image of Galvanic corrosion

Fig. 3.9 (a) represents Zn–Fe couple, in which zinc is more active since it has a
position that is higher in the emf series. Dissolves in preference to iron which is a less
active metal i.e., Zn acts as anode and undergoes corrosion and Fe acts as cathode.
Fig. 3.9 (b) represents Fe–Cu couple, in which iron is more active when compared
to Cu dissolves in preference to copper is less active i.e., Fe acts as anode and
undergoes corrosion and Cu acts as cathode.
Examples for Galvanic Corrosion
• Steel screw in a brass marine hardware corrodes
Due to galvanic corrosion, Iron which has a higher position in electrochemical
series becomes an anode and gets corroded, while brass which is lower in
electrochemical series acts as cathode and is not attacked.
• Bold and Nut made of the same metal is preferred
To avoid galvanic corrosion, it is made up of the same metal so it has no anodic
and cathodic part.

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Prevention
Galvanic corrosion can be reduced by insulating between the two metals.
2. Differential aeration (or) concentration cell corrosion
Corrosion of metals due to the uneven supply of air on the metal surface is
known as differential aerations corrosion. It is caused when metals are exposed to
different concentrations of oxygen or any electrolyte on the surface of the base metal.
The part of the metal exposed to high concentration of oxygen acts as a cathode and
part of the metal exposed to low concentration of oxygen acts as anode. Hence, the
region that is less oxygenated undergoes corrosion.
A good example is a motor pump where the metal pipe is partially immersed in water
or a conducting solution, it is also called water line corrosion.

Fig. 3.8. Differential aeration corrosion

If a metal is immersed partially into a solution, (Fig. 3.8.) the part of the metal
above the solution is considered to be more aerated and hence become cathodic. On
the other hand, the part of the metal inside the solution is considered to be less
aerated and thus, become anodic and suffers corrosion.
At anode (less aerated) corrosion occurs
M⎯
⎯→ M2+ + 2e−
At cathode (more aerated part) OH– ions are produced.
½O2 + H2O + 2e– ⎯ ⎯→ 2OH–
Examples for Differential aeration corrosion
a. Pitting or localized corrosion
b. Crevice corrosion.
c. Pipeline corrosion
d. Corrosion on wire fence.

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a. Pitting corrosion
Pitting corrosion occurs in an area where there is a formation of a hole and the
area around the metal is unattacked.

Fig. 3.9. Pitting corrosion

Example: drop of water, sand, dust, scale, etc on a metal.
Consider a railway track during a rainy day, a drop of water or aqueous solution
falls on the metal surface (Fig. 3.9). The area below the drop of water acts as an anode
since there is less oxygen concentration. This area suffers corrosion. The area that is
freely exposed to air acts as a cathode, due to high oxygen concentration.
The rate of corrosion will be more where the area of cathode is larger and the
area of anode is smaller. Therefore, more and more material is removed from the
same spot. Thus a small hole or pit is formed on the surface of the metal so it is
called as pitting corrosion.
At anode
At the anode, Iron is oxidized to Fe2+ ions.
Fe → Fe2+ + 2e–
At cathode
At the cathode, oxygen is converted to OH– ions.
½O2 + H2O + 2e–→ 2OH–
Net reaction is
Fe2+ + 2OH–→ Fe(OH)2 → Fe(OH)3
b. Crevice corrosion
Crevice corrosion occurs in places where small volumes of solutions that are
stagnant below seals or in crevices, a good example is a nut and rivet head. Deposits
of dust, sand and corrosion products which cannot be removed can also lead to
crevice corrosion. It is seen particularly in metals which are unstable in the presence
of high Cl- and H+ ions concentrations.

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 Crevice corrosion takes place when alloys are exposed to a chloride rich media
gradual acidification take place inside the crevice that leads to corrosion. Galvanic
corrosion does not occur when crevice corrosion occurs. (Fig. 3.10.)

Fig.3.10. Mechanism and Photographic image of crevice corrosion

c. Pipeline corrosion

Fig. 3.11. Pipeline corrosion

Differential aeration corrosion may also occur in different parts of pipeline.
Buried pipelines or cables passing through the soil, from one type of soil to
another say, from clay that is less aerated to sand that is more aerated may get
corroded due to differential aeration.(Fig. 3.11.)
d. Corrosion on a wire fence
Corrosion on a wire fence occurs in areas where the wires inter cross with one
another. This area seems to be less aerated and acts as an anode. The rest of the
fence acts as a cathode. The part that acts as an anode results in corrosion.( Fig.
3.12.)

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Fig.3.12. Mechanism and Photographic image of corrosion on a wire fence
Other examples for differential aeration corrosion
• Corrosion occurring under metal washers, where oxygen cannot diffuse easily.
• Lead pipeline passing through clay to cinders undergo corrosion. Since the
pipeline under cinders is more aerated, it gets corroded easily.
3.6. FACTORS INFLUENCING THE RATE OF CORROSION
The rate and extent of corrosion mainly depends on
1. Nature of the metal.
2. Nature of the environment.
3.6.1 Nature of the metal
a. Position in emf series
Position of the metal in the emf series plays a major role in the extent of
corrosion. All metals in the anodic region above hydrogen in emf series have a
tendency to get corroded. Lower the reduction potential of the metal the greater is the
rate of corrosion. When two metals are in electrical contact metal that has high
negative reduction potential undergoes corrosion. So, the rate of corrosion
depends on the difference in position in the electrochemical series. Greater the
difference higher is the corrosion rate.
b. Relative areas of the anode and cathode
The rate of corrosion increases when the cathodic area is high. When the area
of the cathode is high the demand for electrons is also high which increases the rate
of corrosion.
c. Purity of the metal
Pure metals will not corrode. The presence of impurities in a metal creates
heterogeneity and thus acts like a galvanic cell. It sets up a distinct anode and
cathode in the metal. Higher the amount of impurity greater is the rate of corrosion
at the anode.
The effect of impurities on the rate of corrosion of zinc is given below.
% purity of zinc 99.999 99.99 99.95
Corrosion rate 1 2650 5000

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d. Over voltage or over potential
The over voltage of a metal plays an important role in corrosion. Since over
voltage of the corrosive environment is inversely proportional to corrosion rate. For
example the normal hydrogen over voltage of zinc metal when it dipped in 1 M H2SO4
is 0.7 volt. Here the rate of corrosion is very low. By adding small amount of impurity
like CuSO4 to H2SO4, the hydrogen over voltage goes down to 0.33 V. This results in
the corrosion of zinc metal.
e. Nature of the surface film
The nature of the oxide film formed on the metal surface decides the extent of
corrosion which can be decided by Pilling – Bedworth rule
i. In the case of alkaline earth metals such as Na, Ca, and alkali metals from
oxide, whose volume is less than the volume of the metal. Hence the oxide
film will be porous and bring about further corrosion.
ii. But in heavy metals Al, Cr, etc., form oxide, whose volume is greater than
that of the metal. Hence the oxide film will be non–porous and prevents
further corrosion.
f. Nature of the corrosion product
If the corrosion product dissolves in the medium the corrosion rate is faster.
Similarly, if the corrosion product is volatile (like MoO3 on Mo surface), the corrosion
rate will be faster.
3.6.2. Nature of the environment
a. Temperature
Temperature plays a major role in corrosion since it is directly proportional to
the rate of corrosion. The rate of a chemical reaction and formation of the product
increases with increase in temperature. So the corrosion rate increases with
temperature.
b. Humidity
Corrosion rate increases when the rate of humidity increases in the environment.
Moisture is a good solvent for oxygen to produce the electrolyte. This is required for a
corrosion cell.
c. Presence of corrosive gases
The acidic gases like, CO2, SO2, H2S and fumes of HCl, H2SO4, etc., produce
electrolytes, which are acidic and increases the electrochemical corrosion.
d. Presence of suspended particles
Powerful electrolyte particles like NaCl, (NH4)2, SO4 along with moisture
accelerate the electrochemical corrosion.
e. Effect of pH and electrode potential
The corrosion rate with respect to pH and electrode potential of metals is studied
with the help of a Pourbaix diagram.

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Pourbaix diagram

Fig. 3.13. Pourbaix diagram

The Pourbaix diagram for iron (Fig. 3.13.) shows that metallic iron consists of
two regions of corrosion, the passivity region and the immunity region. The Passivity
region indicates low corrosion rate, due to the formation of a protective metal oxide
layer on the metal. The Immunity region indicates no corrosion, since this region has
an internal cathodic protection. These two regions signify the release of iron into the
solution. In the diagram, 'Z' is the point where pH is 7 and the electrode potential is
–0.4 V. It is present in the corrosion zone. This clearly shows the iron rusts in water
under those conditions. In actual practice, it is observed to the true. The corrosion
rate can be changed by creating a shift in the point 'Z' into immunity or passivity
regions. The iron becomes immune when the potential is changed to about – 1.2 V
by applying external current. The rate of corrosion of iron can also be decreased by
moving the point to the passivity region on applying a positive potential. The diagram
clearly indicates that the rate of corrosion can also be decreased by increasing the pH
by adding an alkali.
Thus the rate of corrosion will be maximum when the corrosive environment is
acidic. i.e. pH is less than 7.

3.7. Corrosion Control

By cathodic protection
When a metal is in contact with an electrolyte, there is a high chance of it to
get corroded. Cathodic protection plays an important role to save these metals from
getting corroded. The principles involved in the cathodic protection makes a metal to
behave like a cathode.

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 The important cathodic protections are
(i) Sacrificial anodic protection.
(ii) Impressed current cathodic protection.

(i) Sacrificial anodic protection method

A Sacrificial Anode is a metal that is highly active. It prevents less active
material surface that are in the more anodic region in the electrochemical series from
corroding. The sacrificial anode is a metal alloy which placed in the more anodic
region in the electrochemical series which has a more negative electrochemical
potential than the metal that has to be protected. The sacrificial anode gets
consumed. The materials used for sacrificial anodes are zinc or magnesium or
aluminium alloys. Current flows from the sacrificial anode to the protected metal
which becomes more cathodic creating a galvanic cell. The oxidation of the protected
metal gets transferred to the sacrificial anode which gets corroded and thus protects
the metal structure.(Fig. 3.18.)

Fig. 3.18. Sacrificial anodic protection

Applications of sacrificial anodic protection
(a) This method is used for the protection of ships and boats. Sheets of Mg or
Zn are hung around the hull of the ship (Fig. 3.19.). Zn or Mg will act as anode
compared to iron (ship or boat is made of iron), so corrosion concentrates on Zn or
Mg. Since they are sacrificed in the process of saving iron, they are called sacrificial
anodes.

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Fig. 3.19. Sacrificial anodic protection
b. Protection of underground pipelines and cables from soil corrosion (Fig. 3.19.).
c. Insertion of Mg sheets into the domestic water boilers to prevent the formation
of rust (Fig. 3.19.).
d. Calcium metal is introduced to decrease engine corrosion.

(ii) Impressed current cathodic protection method (ICCP)

Since sacrificial anodes cannot deliver enough protection to large structures
impressed current cathodic protection method is used. It is used in larger structures
where electrolyte resistivity is high. It consists of an anode connected to a power
source. Anodes used may be tubular or solid rod shapes or continuous ribbon shaped.
They contain silicon cast iron, platinum, graphite and mixed metal oxide. They are
connected with a niobium coated wire. For huge pipelines such as those used in oil
wells and for water supply anodes are set in ground beds depending on the field
condition including current distribution requirements. The rectifier is connected to
the anode and the output is connected to the structure to be protected through
cathodic protection system. An optimised power has to be provided to give protection
to the target structure. The operating current output is often checked to make sure
there is proper supply. ICCP‘s are normally used in structures present in shores and
complex structures. (Fig. 3.20.).
Application of impressed current protection
Structures like tanks, pipelines, transmission line towers, underground water
pipe lines, oil pipe lines, ships etc., can be protected by this method.

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Fig. 3.20. Impressed current cathodic protection

Table 3.2 Comparison of Galvanic method (Sacrificial anode) and impressed

current Cathodic method

Sl.No. Sacrificial anodic method Impressed current method

1. No external power supply is External power supply must be
necessary. present.
2. This method requires periodical Here anodes are stable and do not
replacement of sacrificial anode. disintegrate.
3. Investment is low. Investment is more.
4. Soil and microbiological corrosion Soil and microbiological corrosion
effects are not taken into account. effects are taken into account.
5. This is most economical method This method is well suited for large
especially when short–term structures and long term operations.
protection is required.
6. This method is suitable when the But this method can be practiced
current requirement and the even if the current requirement and
resistivity of the electrolytes are the resistivity of the electrolytes are
relatively low. high.

3.8 PROTECTIVE COATINGS

Protective coatings are used to protect the metals from corrosion. Protective
coatings save the metal surface from the environment. They are also used in
decorative purpose. They have properties such as hardness, oxidation–resistance,
thermal insulating properties and electrical properties. They also protected surface.
3.8.1. PAINT
Paint is a mechanical dispersion of one or more finely divided pigments. They are
present in a medium that contains both the thinner and the vehicle. When paint is
applied on a metal surface, the thinner that is highly volatile evaporates but the
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vehicle undergoes slow oxidation and coats the pigmented film to the metal surface.
It is a liquid which after application to a metal surface forms a thin solid film. It
provides colour and texture to objects.
Definition
3.8.1.1. Characteristics of a good paint
1. It should spread easily on the metal surface.
2. It should have high hiding (covering) power.
3. It should not crack on drying.
4. It should adhere well to the surface.
5. The colour of the paint should be stable.
6. It should be a corrosion and water resistant.
7. It should give a glossy film.

3.8.1.2. Constituents and their functions of a paint.

1. Pigments
Pigments are granular solids in the paint. They give colour, toughness and
texture to the surface. There are two types of pigments, natural or synthetic. Pigments
make paint opaque and protects the surface from ultraviolet light. Some pigments are
given below.
Example
i. White pigments White lead (2PbCO3 – Pb(OH)2)
Lithophone (75% BaSO4 + 25% ZnS)
ii. Black pigments Lamp black, carbon black.
iii. Red pigments Venetian red (Fe2O3 and CaSO4
Indian red (Fe2O3).
iv. Blue pigments Prussian blue Fe4[Fe(CN06]3.
v. Green pigments Chromium oxide.

2. Vehicle (or) drying oil (or) Binder (or) resins

The Vehicle is a binder. It is component that forms a thin film. They
provide adhesion, gloss, durability, flexibility, and toughness. Alkyds, acrylics, vinyl-
acrylics, vinyl acetate, polyurethanes, polyesters, melamine resins, and oils act as
binders. They provide oxidative cross-linkages when they get dried. They undergo slow
oxidation and coat the pigments onto the metallic surface.
3. Thinners (or) solvents (or) Diluents
Thinners play a major role to dissolve the polymer and adjust the viscosity of the
paint. They are highly volatile, act as diluents and controls flow. They are carriers for
the non volatile components, such as heavy oils. These volatile substances have only
a temporary property, once the solvent evaporates the remaining paint gets fixed onto
the surface. Aliphatics, aromatics, alcohols, ketones and white spirit are good
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examples of thinners elements such as petroleum, esters, glycol ethers act as
thinners.
4. Extenders (or) fillers
Fillers are pigments that thicken the film. They support the structure and
increase the volume of the paint. Fillers are cheap and inert materials.
Talc, lime, barites, clay are good examples for fillers.
5. Driers
These are the substances, used to accelerate the processing of drying.
Functions
i. They act as oxygen – carries (or) catalysts.
ii. They provide oxygen, which is essential oxidation, polymerisation of drying
oil.
Examples: Metallic soaps, linoleates and resinates of Co, Mn and Pb.
6. Plasticisers
These are chemical added to the paint to provide elasticity to the film and to
prevent cracking of the film.
Examples : Triphenyl phosphate, tricresyl phoshate etc.,
7. Anti–skinning agents
These are chemicals added to the paint to prevent gelling and skinning of the
paint.
Example : Polyhydroxy phenol.
Pigment Volume Concentration (P.V.C)
It is an important property of a paint. The following equation is used to calculate
te P.V.C.
Volume of pigment in the paint
P.V .C. =
Volume of pigment + vehicle in the paint
Higher the volume of P.V.C lower will be the durability, adhesion, consistency of
the paint.
3.9.1. METALLIC COATINGS
3.9.1.1. Electroplating (or) Electro–deposition
Electroplating is the process in which the coating a metal such as Nickel or
Copper or Gold is deposited on the base metal such as iron or steel by passing a direct
current through an electrolytic solution such as corresponding soluble salt of the
coating metal. Electroplating is used to change the properties like abrasion, wear
resistance, corrosion protection, lubricating power and aesthetic qualities of a metal.

Construction
Pure nickel metal is made as anode. The object to be coated acts as cathode.
Both the anode and the cathode are immersed in a solution of Nickel sulphate, Nickel

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chloride and Borax mixture. Anode and cathode are connected to an external power
supply. The following conditions are maintained for Nickel plating.
Ni Pellets or pieces in Titanium
Anode
mesh
Nickel sulphate NiSO4.6H2O 240-300 g/l
Nickel chloride NiCl2.6H2O 30-90 g/l
Electrolyte
Boric acid H3BO3 30-45 g/l
Water
Coumarin, Saccharin, Benzene
Additives
sulphonamide
Operating Temperature: 105-150°F
pH 3.0-4.5
Cathode current density 20-100 mA cm-2
Brighteners p-toluene sulfonamide, benzene
sulphonic acid
Levelers, second class allyl sulfonic acid, formaldehyde
brighteners chloral hydrate

Working
In the anode, oxidation occurs where Ni atoms are converted in to Ni2+ ions
and dissolves in to the solution. The dissolved metal ions in the electrolyte solution
moves towards the cathode. These Ni2+ ions get reduced at the cathode and deposit
as Ni at the object. The rate at which the anode is dissolved is equal to the rate at
which the cathode is plated. The rate of plating also depends on the current flowing
through the circuit.
At anode: Ni → Ni2+ + 2e–

At Cathode: Ni2+ + 2e– → Ni

In order to get strong, adherent and smooth deposit certain additives (glue, gelatin,
etc.,) are added to the electrolytic bath. To improve the brightness of deposit,
brightening agents are added in the electrolytic bath.

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Application:
1. It used to improve the appearance of exterior auto parts, such as grills and
bumpers.
2. Nickel electroforming, is commonly used in the aerospace, textile and
communication industries.

Advantages:
1. It used to improve the physical properties such as resistance to wear, heat or
corrosion.
2. It serves the dual purpose of providing a bright, attractive finish as well as
imparting improved corrosion resistance

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ANNA UNIVERSITY QUESTIONS
PART A
1. What is corrosion? Mention its types.
The gradual deterioration or degradation of a material due to a reaction with its
environment is called as corrosion. The types are:
A) Dry or chemical corrosion
➢ Oxidation corrosion
➢ Corrosion by hydrogen
➢ Liquid-metal corrosion
B) Wet or electrochemical corrosion
➢ Galvanic corrosion
➢ Differential aeration.

2. What is Pilling-Bedworth ratio? State its importance.

The Pilling-Bedworth ratio or P-B ratio is the ratio of the volume of the metal
oxide formed to the volume of metal consumed.
P-B ratio = Volume of metal oxide formed
Volume of metal consumed
Importance of P-B ratio:
Through this ratio we can predict whether the metal is likely to form a
protective oxide layer.

3. Distinguish between chemical and electrochemical corrosion.

SI.NO CHEMICAL CORROSION ELECTROCHEMICAL CORROSION

It occurs in the presence of water or
1 It occurs only in dry condition
electrolyte.
It is due to the setup of a large
It is due to the direct chemical
2 number of cathodic and anodic
attack of metal on the environment
areas.
Even a homogeneous metal surface It requires a heterogeneous surface
3
gets corroded or bimetallic contact.
Corrosion products accumulate Corrosion occurs at the anode, while
4
where the corrosion occurs. products formed elsewhere.
5 It is self-controlled. It is a continuous process.
Follows electrochemical reaction.
Follows adsorption mechanism.
Eg: rusting of iron in moist
6 Eg: formation of mild scale on iron
atmosphere.
surface.

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4. Define galvanic corrosion.
The corrosion that occurs when two different metallic materials are connected with
each other in the presence of moisture or electrolyte. The material which is more
anodic will corrode.
Example: Steel screw in a brass marine hardware corrodes.

5. What is differential aeration corrosion?

Corrosion of metals due to the uneven supply of air on the metal surface is known as
differential aeration corrosion.
At anode: M M2+ + 2e-
At cathode: 1/2O2 + H2O +2e- 2OH-
Example : Pipeline corrosion and Pitting corrosion

6. Steel screw in brass hardware corrodes. Why?

This is due to the galvanic corrosion. Iron which has higher position in
electrochemical series becomes an anode and gets corroded, while brass which is
lower in electrochemical series act as cathode and gets not attacked.

7. What is pitting corrosion?

Pitting corrosion is a type of differential aeration corrosion. It occurs in an area where
there is a formation of a hole and the area around the metal is unattacked.
At anode: Fe Fe2+ +2e-
At cathode: 1/2 O2 + H2O + 2e- 2OH-
Net reaction: Fe2++ 2OH- Fe(OH)2 Fe(OH)3
Example: Metal area covered by water droplet.
8. What is pipeline corrosion?
Differential aeration corrosion may also occur in different parts of pipeline. Buried
pipelines or cables passing through the soil, from one type of soil to another say, from
clay that is less aerated to sand that is more aerated may get corroded due to
differential aeration.

9. Zinc is more readily corroded when coupled with Cu than with Pb. why?
This is due to the galvanic corrosion. When the two different metals like Cu and Zn
are in direct contact with each other then Cu (E0 = +0.34 V) acts a cathode and Zn
(E0 = - 0.76 V) acts as anode.
At anode: Zn Zn2+ + 2e-
Zn2++ 2OH- Zn(OH)2
At cathode: Cu + 2e
2+ - Cu
Thus due to the above reaction, Zn gets corroded.

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10. Small anodic area results in intense corrosion. Why?
The rate of corrosion is rapid and intense if anodic area is smaller and the cathodic
area is bigger because the demand for electron is more.

11. What is paint?

Paint is the mechanical dispersion of one or more finely divided pigments. They are
present in a medium that contains both the thinner and the vehicle. It’s a liquid which
provides colour and texture to objects.

12. Mention the factors that affect the rate of corrosion.

i) Nature of the metal
➢ Position in emf series
➢ Relative areas of the anode and cathode
➢ Purity of the metal
➢ Over voltage or over potential
➢ Nature of the surface film
➢ Nature of the corrosion product.
ii) Nature of the environment
➢ Temperature
➢ Presence of corrosive gases
➢ Humidity
➢ Presence of suspended particles
➢ Effect of pH and electrode potential.

13. What is the role of pigment in paints?

Pigments are granular solids in the paint. They give colour, toughness and texture to
the surface. Pigments make paint opaque and protect the surface from ultraviolet
light.

14. What is the function of driers in the paint?

i) These are used to accelerate the processing of drying.
ii) They act as oxygen carriers or catalyst.
iii) They provide oxygen, which is essential for oxidation and polymerization of drying
oil.

15. Define electroplating.

It is a process in which the coating metals such as copper or gold on the base metals
such as iron or steel by passing a direct current through an electrolytic solution
containing the soluble salt of the coating metal such as CuSO4 or AuCl3.
Example: Electroplating of copper over iron.

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16. What is electro less plating?
It is a technique of depositing a noble metal such as nickel from its salt solution on
a catalytically active surface of the metal or non metal, to be protected, by using a
suitable reducing agent. No electrical energy is used.
Metal ions + Reducing agents Metal + Oxidized products.

17. Explain the advantages of electro less plating over electroplating?

i) No electricity is required.
ii) Electro less plating on insulators and semiconductors can be easily carried out.
iii) Complicated parts can also be plated uniformly.
iv) Electro less coatings posses good mechanical chemical and magnetic properties.

PART B
1. Difference between Chemical and electrochemical corrosion. Mention any four
factors that affect electrochemical corrosion.
2. i. Describe the mechanism of electrochemical corrosion by hydrogen evolution
and oxygen absorption.
3. What is sacrificial anode? Mention its role in prevention of corrosion.
4. What is chemical corrosion? Explain the mechanism of oxidation corrosion.
5. How is corrosion prevented in ship hulls and transmission line towers. Explain.
6. Explain water–line corrosion.
7. What should be the nature of corrosion product to prevent further corrosion?
8. How is galvanic corrosion occur?
9. State and explain pilling–bedworth rule.
10. Deposition of oil or dust or metal surfaces for a long period is undesirable. Give
reasons.
11. Mention the possible means of rendering a metal cathodic to protect it from
corrosion.
12. Describe the mechanism of differential aeration corrosion taking pitting as
example.

13. Explain the important factors which influence the rate of corrosion of a metal.

14. Explain the electrochemical theory of corrosion with suitable example.

15. Write a brief note on
i. Cathodic protection by sacrificial anode.
ii. Differential aeration corrosion.
16. Discuss the mechanism of chemical and electrochemical corrosion with suitable
example.

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17. Write a note on pitting corrosion and cathodic protection.
18. Explain the principle of electrochemical corrosion with suitable example.
19. Describe the sacrificial anode methods of corrosion control.
20. What are factors, influencing chemical and electrochemical corrosion?
21. Write a note on application of inhibitors in corrosion control.
22. Define corrosion and explain the various factors influencing corrosion of a metal.

23. Enlist the different methods of corrosion control. Explain the cathodic protection
method in detail.
24. Explain the various types of corrosion.
25. Discuss the mechanism of electrochemical corrosion.
26. Discuss about the corrosion control methods.
27. What is cathodic protection? Explain the sacrificial anode and impressed
current techniques for the prevention of corrosion.
28. Explain the control of corrosion by the use of sacrificial anodes and by impressed
current cathodic protection.
29. Explain the factors connected with metal that effect corrosion.
30. What are cathodic and anodic protection for controlling corrosion? Discuss their
merits and demerits.
31. Describe corrosion control through impressed current method.
32. Explain differential aeration corrosion with an example.
33. Explain chemical corrosion. Also explain the intensity of corrosion varying with
the nature of oxide layer formation over metal.
34. What is differential aeration corrosion? Write its mechanism. Give any
illustrations that show differential corrosion.
35. What is paint? What are the consistents and functions of paints.
36. What are the requisites of a good paint?
37. Write detailed notes on pigments.
38. Describe briefly the important parameters involved in the electroplating.
39. What are the main objectives of electroplating. Give account of the method used
in electroplating of Nickel.
40. Explain the process of electrodeposition taking a suitable example.

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