UNIT IV- CORROSION AND ITS CONTROL

CONTENTS

4.1 Causes of corrosion

4.2 Consequences of corrosion
4.3 Classification or theories of corrosion
4.4 Passivity
4.5 Factors influencing corrosion
4.6 Corrosion Control ( Protection Against Corrosion)

Applications
Questions
Assignments
References

Introduction :
Corrosion of metals is defined as the spontaneous destruction of metals in the course of their chemical ,
electrochemical or biochemical interaction with the environment.
Corrosion is an undesirable process. Due to corrosion there is limitation of progress in many areas.
Example : Rusting of iron – formation of a reddish brown scale or powder of hydrated iron oxide on its
surface.

4.1 Causes of corrosion:

In nature, metals occur in two different forms.

(1) Native State (2) Combined State

Native State: The metals exist as such in the earth crust then the metals are present in a native
state. Native state means free or uncombined state. These metals are non- reactive in nature. They are
noble metals which have very good corrosion resistance. Example: Au, Pt, Ag, etc.,

Combined State: Except noble metals, all other metals are highly reactive in nature which undergoes
reaction with their environment to form stable compounds called ores and minerals. This is the
combined state of metals. Example: Fe2O3, ZnO, PbS, CaCO3, etc.,

Metallic Corrosion: The metals are extracted from their metallic compounds (ores). During the extraction,
ores are reduced to their metallic states by applying energy in the form of various processes. In the pure
metallic state, the metals are unstable as they are considered in an excited state (higher energy state).
Therefore as soon as the metals are extracted from their ores, the reverse process begins and form
metallic compounds, which are thermodynamically stable (lower energy state). Hence, when metals are
used in various forms, they are exposed to the environment, the exposed metal surface begins to decay
(conversion to more stable compound). This is the basic reason for metallic corrosion.

Corrosion-Oxidation
Metal Metallic Compound +
Energy
Metallurgy-Reduction

Although corroded metal is thermodynamically more stable than pure metal, due to corrosion, useful
properties of a metal like malleability, ductility, hardness, luster and electrical conductivity are lost.

4.2 Consequences of corrosion:

The economic and social consequences of corrosion include

1. Enormous wastage of machineries and different types of metallic materials.
2. It leads to unpredictable machinery failure which may lead to dangerous situations or accidents
causing loss of life.
3. Due to formation of corrosion product over the machinery, the efficiency of the machine
decreases and frequent replacement of corroded equipments increases cost.
4. It may cause leakage of toxic chemicals and gases leading to the contamination of the environment
and health hazards.
5. Preventive maintenance like metallic coating or organic coating is required.
6. It may cause contamination of potable water and food stored in metallic containers.

4.3 Classification or theories of corrosion:

Based on the environment, corrosion is classified into

(i) Dry or Chemical Corrosion (ii) Wet or Electrochemical Corrosion

4.3.1 Dry or Chemical Corrosion:

This type of corrosion is due to the direct chemical attack of metal surfaces by the atmospheric gases
such as oxygen, halogen, hydrogen sulphide, sulphur dioxide, nitrogen or anhydrous inorganic liquid,
etc. Chemical corrosion is defined as the direct chemical attack of metals by the atmospheric gases
present in the environment.
Example: (a) Silver materials undergo chemical corrosion by Atmospheric H2S gas . (b) Iron metal
undergoes chemical corrosion by HCl gas.

Types of dry or chemical corrosion:

I) Corrosion by Oxygen or Oxidation corrosion
II) Corrosion by Hydrogen
 III) Liquid Metal Corrosion

(a) Corrosion by oxygen or oxidation corrosion:

Oxidation Corrosion is brought about by the direct attack of oxygen at low or high temperature on
metal surfaces in the absence of moisture. Alkali metals (Li, Na, K et c.,) and alkaline earth metals
(Mg, Ca, Sn, etc.,) are rapidly oxidized at low temperature. At high temperatures, almost all metals
(except Ag, Au and Pt) are oxidized. The reactions of oxidation corrosion are as follows:

Mechanism:
(i) Oxidation takes place at the surface of the metal forming metal ions M 2+
2+ -
M → M + 2e
2-
(ii) Oxygen is converted to oxide ion (O ) due to the transfer of electrons from metal (reduction).
n/ 2 O2 + 2n e- → n O2-
(iii) The overall reaction of oxide ions reacts with the metal ions to form metal oxide film.
2 Mn+ + nO2- → 2 M O
The Nature of the Oxide formed plays an important part in the oxidation corrosion process.
Metal + Oxygen → Metal oxide (corrosion product)

When oxidation starts, a thin layer of oxide is formed on the metal surface and the nature of this
film decides the further action. If the film is

a) Stable layer:
A Stable layer is fine grained in structure and can get adhered tightly to the parent metal surface.
Hence, such a layer can be of impervious nature (ie., which cuts-off penetration of attaching oxygen to
the underlying metal). Such a film behaves as a protective coating in nature, thereby shielding the
metal surface. The oxide films on Al, Sn, Pb, Cu, Pt, etc., are stable, tightly adhering and impervious
in nature.

b) Unstable oxide layer:

This is formed on the surface of noble metals such as Ag, Au, Pt. As the metallic state is more stable
than oxide, it decomposes back into the metal and oxygen. Hence, oxidation corrosion is not possible
with noble metals.

c) Volatile oxide layer:

The oxide layer film volatilizes as soon as it is formed. Hence, always a fresh metal surface is
available for further attack. This causes continuous corrosion. M oO 3 is volatile in nature.

d) Porous layer:
The layer has pores or cracks. In such a case, the atmospheric oxygen has access to the underlying
surface of metal, through the pores or cracks of the layer, thereby the corrosion continues unobstructed,
till the entire metal is completely converted into its oxide.

Pilling-Bedworth rule:

According to it “ an oxide is protective, if the volume of the oxide is at least as great as the volume of
the metal from which it is formed”. On the other hand, “ if the volume of the oxide is less than the
volume of metal, the oxide layer is porous (or non-continuous) and hence, non-protective, because it
cannot prevent the access of oxygen to the fresh metal surface below”.
Thus, alkali and alkaline earth metals ( like Li, K, Na, M g) form oxides of volume less than the
volume of metals. Consequently, the oxide layer faces stress and strains, thereby developing cracks
and pores in its structure. Porous oxide scale permits free access of oxygen to the underlying metal
surface (through cracks and pores) for fresh action and thus, corrosion continues non-stop.
Metals like Aluminium forms oxide, whose volume is greater than the volume of metal. Consequently,
an extremely tightly-adhering non-porous layer is formed. Due to the absence of any pores or cracks in
the oxide film, the rate of oxidation rapidly decreases to zero.

Pilling-Bedworth ratio = Volume of the metal oxide formed / Volume of the metal consumed

(b) Corrosion by hydrogen :

(i) Hydrogen Embrittlement:

Loss in ductility of a material in the presence of hydrogen is known as hydrogen embrittlement

Mechanism:
This type of corrosion occurs when a metal is exposed to a hydrogen environment. Iron liberates
atomic hydrogen with hydrogen sulphide in the following way.
Fe + H2S → FeS + 2H
Hydrogen diffuses into the metal matrix in this atomic form and gets collected in the voids present inside
the metal. Further, diffusion of atomic hydrogen makes them combine with each other and forms
hydrogen gas.
H + H → H2↑
Collection of these gases in the voids develops very high pressure, causing cracking or blistering of
metal.

(ii) Decarburisation:
The presence of carbon in steel gives sufficient strength to it. But when steel is exposed to a hydrogen
environment at high temperature, atomic hydrogen is formed.
heat
H2 2H
Atomic hydrogen reacts with the carbon of the steel and produces methane gas.

C + 4H → CH4

Hence, the carbon content in steel decreases. The process of decrease in carbon content in steel is
known as decarburization.
Collection of methane gas in the voids of steel develops high pressure, which causes cracking. Thus,
steel loses its strength.

Liquid metal corrosion:

This is due to chemical action of flowing liquid metal at high temperatures on solid metal or alloy. Such
corrosion occurs in devices used for nuclear power. The corrosion reaction involves either: (i)
dissolution of a solid metal by a liquid metal or (ii) internal penetration of the liquid metal into the
solid metal. Both these modes of corrosion cause weakening of the solid metal.
4.3.2 Wet or Electrochemical Corrosion

Electrochemical corrosion involves:

(i) The formation of anodic and cathodic areas or parts in contact with each other
(ii) Presence of a conducting medium
(iii) Corrosion of anodic areas only and
(iv) Formation of corrosion products somewhere between anodic and cathodic areas. This involves
flow of electron-current between the anodic and cathodic areas.

At anodic area oxidation reaction takes place ( liberation of free electrons), so anodic metal is destroyed
by either dissolving or assuming combined state (such as oxide, etc.). Hence corrosion always occurs at
anodic areas.
n+ -
M (metal) → M +ne
n+
M (metal ion) → Dissolves in solution → forms compounds such as oxide

At the cathodic area, a reduction reaction takes place (gain of electrons), usually cathode reactions do not
affect the cathode, since most metals cannot be further reduced. So at the cathodic part, dissolved
constituents in the conducting medium accept the electrons to form some ions like OH - and O2-.
Cathodic reaction consumes electrons with either (a) evolution of hydrogen or (b) absorption of oxygen,
depending on the nature of the corrosive environment.

Hydrogen Evolution Type:

All metals above hydrogen in the electrochemical series have a tendency to get dissolved in acidic solution
with simultaneous evolution of hydrogen. It occurs in an acidic environment. Consider the example of iron
At anode: Fe → Fe2+ + 2e-
These electrons flow through the metal, from anode to cathode, where H+ ions of acidic solution are
eliminated as hydrogen gas.
+ -
At cathode: 2 H + 2 e → H2 ↑
+ 2+
The overall reaction is: Fe + 2H → Fe + H2

Oxygen Absorption Type:

Rusting of iron in neutral aqueous solution of electrolytes (like NaCl solution) in the presence of
atmospheric oxygen is a common example of this type of corrosion. The surface of iron is usually coated
with a thin film of iron oxide. However, if this iron oxide film develops some cracks, anodic areas
are created on the surface; while the well metal parts act as cathodes
.
At Anode: Metal dissolves as ferrous ions with liberation of electrons.
2+ -
Fe → Fe + 2e
At Cathode: The liberated electrons are intercepted by the dissolved oxygen.
½ O2 + H2O + 2 e- → 2OH-
2+ -
The Fe ions and OH ions diffuse and when they meet, ferrous hydroxide is precipitated.
2+ -
Fe + 2OH → Fe(OH)2
(i) If enough oxygen is present, ferrous hydroxide is easily oxidized to ferric hydroxide.
4Fe(OH)2 + O2 + 2H2O → 4Fe(OH)3
(Yellow rust Fe2O3.H2O)
(ii) If the supply of oxygen is limited, the corrosion product may be even black anhydrous magnetite,
Fe3O4.

Differences between ( dry) chemical and ( wet) electrochemical corrosion:

Sl. Chemical Corrosion Electrochemical Corrosion

No.
1. It occurs in dry conditions. It occurs in the presence of moisture or
electrolyte.
2. It is due to the direct chemical attack of It is due to the formation of a large number
the metal by the environment. of anodic and cathodic areas.
3. Even a homogeneous metal surface Heterogeneous (bimetallic) surface alone gets
gets corroded. corroded.

4. Corrosion products accumulate at the Corrosion occurs at the anode while the
place of corrosion products are formed elsewhere.

5. It is a self controlled process. It is a continuous process.

6. It adopts an adsorption mechanism. It follows an electrochemical reaction.

7. Formation of a mild scale on an iron Rusting of iron in moist atmosphere is an

surface is an example. example.
TYPES OF ELECTROCHEMICAL CORROSION

The electrochemical corrosion is classified into the following two types: (i) Galvanic (or
Bimetallic) Corrosion
(ii) Differential aeration or concentration cell corrosion.

(i) Galvanic Corrosion:

When two dissimilar metals (eg., zinc and copper) are electrically connected and exposed to an
electrolyte, the metal higher in electrochemical series undergoes corrosion. In this process, the more
active metal ( with more negative electrode potential) acts as anode while the less active metal ( with
less negative electrode potential) acts as cathode.
In the above example, zinc (higher in electrochemical series) forms the anode and is attacked and
gets dissolved; whereas copper (lower in electrochemical series or more noble) acts as cathode.

Mechanism:
In acidic solution, the corrosion occurs by the hydrogen evolution process; while in neutral or slightly
alkaline solution, oxygen absorption occurs. The electron-current flows from the anode metal, zinc, to
the cathode metal, copper.
2+ -
Zn Zn + 2e (Oxidation)

Thus it is evident that the corrosion occurs at the anode metal, while the cathodic part is protected
from the attack.
Example: (i) Steel screws in a brass marine hardware (ii) Lead-antimony solder around copper wise;
(iii) a steel propeller shaft in bronze bearing ( iv) Steel pipe connected to copper plumbing.

(ii) Concentration Cell Corrosion:

It is due to electrochemical attack on the metal surface, exposed to an electrolyte of varying
concentrations or of varying aeration.
It occurs when one part of metal is exposed to a different air concentration from the other part. This
causes a difference in potential between differently aerated areas. It has been found experimentally that
poor-oxygenated parts are anodic.
 Examples:
i) The metal part immersed in water or in a conducting liquid is called water line corrosion.
(ii) The metal part is partially buried in soil.
Explanation: If a metal is partially immersed in a conducting solution the metal part above the solution
is more aerated and becomes cathodic. The metal part inside the solution is less aerated and thus
becomes anodic and suffers corrosion.

At anode:

Corrosion occurs (less aerated) M → M 2+ + 2e-

At cathode:

OH- ions are produced (more aerated) ½O2 + H2O → 2e- + 2OH-
Examples for this type of corrosion are
(i) Pitting or localized corrosion
(ii) Crevice corrosion
(iii) Pipeline corrosion
(iv) Corrosion on wire fence

Pitting Corrosion:

Pitting is a localized attack, which results in the formation of a hole around which the metal is relatively
unattacked.
The mechanism of this corrosion involves setting up of differential aeration or concentration Metal area
covered by a drop of water, dust, sand, scale etc. is the aeration or concentration. Pitting corrosion is
explained by considering a drop of water or brine solution (aqueous solution of NaCl) on a metal surface,
especially iron.
The area covered by the drop of salt solution has less oxygen and acts as anode. This area suffers
corrosion, the uncovered area acts as cathode due to high oxygen content.
It has been found that the rate of corrosion will be more when the area of the cathode is larger and the area
of the anode is smaller. Hence there is more material around the small anodic area resulting in the
formation hole or pit.

At anode: Fe is oxidized to Fe2+ and releases electrons.

Fe Fe2+ + 2e-

At cathode: Oxygen is converted to hydroxide ion

½ O2 + H2O + 2e- 2OH-

The net reaction is Fe + 2OH- ------------ Fe(OH)2

The above mechanisms can be confirmed by using a ferroxyl indicator (a mixture containing
phenolphthalein and potassium ferricyanide). Since OH- ions are formed at the cathode, this area imparts
pink colour with a phenolphthalein indicator. At the anode, iron is oxidized to Fe2+ which combines
with ferricyanide and shows blue colour.

Crevice corrosion:

If a crevice ( a crack forming a narrow opening) between metallic and non-metallic material is in contact
with a liquid, the crevice becomes anodic region and undergoes corrosion. Hence, oxygen supply to the
crevice is less. The exposed area has high oxygen supply and acts as a cathode.
Bolts, nuts, rivets, joints are examples for this type of corrosion.

Pipeline corrosion:

Buried pipelines or cables passing from one type of soil (clay less aerated) to another soil (sand more
aerated) may get corroded due to differential aeration.
Corrosion in wire fence:

A wire fence is one in which the areas where the wires cross (anodic ) are less aerated than the rest of the
fence (cathodic). Hence corrosion takes place at the wire crossing.
Corrosion occurring under metal washers and lead pipeline passing through clay to cinders(ash) are other
examples.

Soil corrosion :

The moisture and dissolved electrolytes in soil make it corrosive. Presence of microorganisms in soil
further leads to corrosion of the underground structures and pipelines. Thus soil corrosion is due to the
combined effect of the presence of
Acidity of soil
Moisture and electrolytes
Micro-organisms
Differential aeration
Passivity :
In an electrochemical series, metals are arranged top to bottom in increasing order of their reduction
electrode potentials. A metal high in the series is more anodic and undergoes corrosion faster than the
metal below it. For example, Zn corrodes faster than Fe which in turn corrodes faster than Sn, because Zn
is placed high above followed by Fe which in turn is followed by Sn in the electrochemical series.
However there are a number of exceptions. For eg. Ti is placed above Ag in the electrochemical series
and is expected to corrode faster than Ag. But it is just the reverse , Ti is corroded to a lesser extent than
Ag. This can only be explained by passivity of the metals. A metal is said to be passive in a certain
environment if it exhibits much lower corrosion rate than what is expected from its position in the
electrochemical series (however on changing the environment it can be rendered active).
Passivity of the metal is due to the formation of a highly protective , but very thin film on the surface of
the metal under particular env. Conditions. The film formed at the surface makes the metal or alloy act like
a noble metal.

Galvanic series :
It is the series in which metals are arranged in the order of their corrosion behavior in different
environments. This is a more reliable series for predicting the corrosion behavior than the electrochemical
series.
4.5 Factors influencing corrosion

There are two factors that influence the rate of corrosion. Hence a knowledge of these factors and the
mechanism with which they affect the corrosion rate is essential because the rate of corrosion is different
in different atmospheres.

1. Nature of the metal 2. Nature of the corroding environment

4.5.1 Nature of the metal:

(a) Physical state:

The rate of corrosion is influenced by the physical state of the metal (such as grain size, orientation of
crystals, stress, etc). The smaller the grain size of the metal or alloy, the greater will be its solubility and
hence greater will be its corrosion. Moreover, areas under stress, even in a pure metal, tend to be anodic
and corrosion takes place at these areas.

(b) Purity of metal:

Impurities in a metal cause heterogeneity and form minute/ tiny electrochemical cells (at the exposed
parts), and the anodic parts get corroded. The cent percent pure metal will not undergo any type of
corrosion. For example, the rate of corrosion of aluminium in hydrochloric acid with increase in the
percentage impurity is noted.

% purity of aluminium 99.99 99.97 99.2

Relative rate of corrosion1 1000 30000

(d) Nature of surface film:

In an aerated atmosphere, practically all metals get covered with a thin surface film (thickness=a few
angstroms) of metal oxide. The ratio of the volumes of the metal oxide to the metal is known as a specific
volume ratio. Greater the specific volume ratio, lesser is the oxidation corrosion rate. The specific volume
ratios of Ni, Cr and W are 1.6, 2.0 and 3.6 respectively. Consequently the rate of oxidation of tungsten is
least, even at elevated temperatures.

(e) Relative areas of the anodic and cathodic parts:

When two dissimilar metals or alloys are in contact, the corrosion of the anodic part is directly
proportional to the ratio of areas of the cathodic part and the anodic part.
Corrosion is more rapid and severe, and highly localized, if the anodic area is small ( eg., a small steel
pipe fitted in a large copper tank) , because the current density at a smaller anodic area is much greater and
the demand for electrons can be met by smaller anodic areas only by undergoing corrosion more briskly.

(f) Position in galvanic series:

Extent of corrosion depends upon the position of the metal in galvanic series. When two metals are in
electrical contact in presence of an electrolyte, the metal higher in the galvanic series becomes anodic and
suffers corrosion. The greater the difference in their position in the series, the faster is the corrosion in the
anodic metal.
(g )Solubility of corrosion products:
If the corrosion product is soluble in the corroding medium then corrosion proceeds at a faster rate.
Whereas insoluble products suppress the corrosion as it functions as a physical barrier. For example the
corrosion of Pb in H2SO4 proceeds at much lower rate because of the formation of insoluble PbSO4.

(h) Volatility of corrosion :

The volatility of the corrosion product has a marked significance on corrosion rate. If the corrosion
product is volatile, it leaves the surface of the metal as soon as it is formed leaving the underlying fresh
metal surface exposed to further attack. This results into excessive corrosion.

4.5.2 Nature of the Corroding Environment:

(a)Temperature:
The rate of corrosion is directly proportional to temperature ie., rise in temperature increases the rate of
corrosion. This is because the rate of diffusion of ions increases with rise in temperature.

(b) Humidity of air:

The rate of corrosion will be more when the relative humidity of the environment is high. The moisture
acts as a solvent for oxygen, carbon dioxide, sulphur dioxide etc. in the air to produce the electrolyte
which is required for setting up a corrosion cell.

(c) Presence of impurities in atmosphere:

Atmosphere in industrial areas contains corrosive gases like CO2, H2S, SO2 and fumes of HCl, H2SO4
etc. In presence of these gases, the acidity of the liquid adjacent to the metal surfaces increases and its
electrical conductivity also increases, thereby the rate of corrosion increases.

(d) Presence of suspended particles in atmosphere:

In case of atmospheric corrosion: (i) if the suspended particles are chemically active in nature (like NaCl,
Ammonium sulphate), they absorb moisture and act as strong electrolytes, thereby causing enhanced
corrosion; (ii) if the suspended particles are chemically inactive in nature (eg., charcoal), they absorb both
sulphur gases and moisture and slowly enhance corrosion rate.

(e) Influence of pH:

Generally acidic media (ie., pH<7) are more corrosive than alkaline and neutral media. However,
amphoteric metals (like Al, Zd, Pb, etc.) dissolve in alkaline solutions as complex ions. The corrosion rate
of iron in oxygen- free water is slow, until the pH is below 5. The corresponding corrosion rate in the
presence of oxygen is much higher. Consequently corrosion of metals, readily attacked by acid, can be
reduced by increasing the pH of the attacking environment, e.g. Zn ( which is rapidly corroded, even in
weakly acidic solutions such as carbonic acid suffers minimum corrosion at pH=11.

a) Nature of ions present:

b) Conductance of the corroding medium:
c) Formation of oxygen concentration cell:
d) Flow velocity of process stream:
e) Polarization of electrodes:
 4.6 Corrosion Control ( Protection Against Corrosion )

As the corrosion process is very harmful and losses incurred are tremendous, it becomes necessary to
minimize or control corrosion of metals. Corrosion can be stopped completely only under ideal conditions.
But the attainment of ideal conditions is not possible. However, it is possible only to minimize corrosion
considerably. Since the types of corrosion are so numerous and the conditions under which corrosion
occurs are so different, diverse methods are used to control corrosion. As the corrosion is a reaction
between the metal or alloy and the environment, any method of corrosion control must be aimed at either
modifying the metal or the environment.

1. Choice of metals and alloys:

(i) The first choice is to use noble metals such as gold and platinum. They are most resistant to corrosion.
As they are precious, they cannot be used for general purposes.

(ii) The next choice is to use the purest possible metal. But in many cases, it is not possible to produce a
metal of high chemical purity. Hence, even a trace amount of impurity leads to corrosion.

(iii) Thus, the next choice is the use of corrosion resistant alloys. Several corrosion resistant alloys have
been developed for specific purposes and environment. For example, a) Stainless steel containing
chromium produces an exceptionally coherent oxide film which protects the steel from further attack. (b)
Cupro-nickel (70% Cu + 30%Ni) alloys are now used for condenser tubes and for bubble trays used in
fractionating columns in oil refineries. (c) Highly stressed Nimonic alloys (Ni-Cr- M o alloys) used in gas
turbines are very resistant to hot gases.

2. Proper Designing:

Proper geometrical design plays a vital role in the control of corrosion of equipment and structures. The
general guidelines of the design of materials and components to control corrosion are the following:
Use always simple design and structure
The design must avoid more complicated shapes having more angles, edges, corners etc.
Avoid the contact of dissimilar metals as they may lead to galvanic type corrosion. To overcome this,
insulation can be used.
When two dissimilar metals are to be in contact, the anodic area must be as large as possible and the
cathodic area should be as small as possible.
As far as possible, crevices (gap or crack) should be avoided between adjacent parts of a structure.
Bolts and rivets should be replaced by proper welding
Metal washers should be replaced by rubber or plastic washers as they do not adsorb water. They also act
as insulation.
Corrosion in pipelines can be prevented by using smooth bends.
Heat treatment like annealing minimizes the stress corrosion.
A good design of water storage container is the one from which water can be drained and cleaned easily.
Such a design avoids accumulation of dirt etc
.
3. Cathodic protection:

The reduction or prevention of corrosion by making metallic structure as cathode in the electrolytic cell is
called cathodic protection. Since there will not be any anodic area on the metal, corrosion does not occur.
There are two methods of applying cathodic protection to metallic structures.
a) Sacrificial anodic protection (galvanic protection)
b) Impressed current cathodic protection

(a) Sacrificial anodic protection method

In this method, the metallic structure to be protected is made cathode by connecting it with more active
metal (anodic metal). Hence, all the corrosion will concentrate only on the active metal. The parent
structure is thus protected. The more active metal so employed is called sacrificial anode. The corroded
sacrificial anode block is replaced by a fresh one. Metals commonly employed as sacrificial anodes are
magnesium, zinc, aluminium and their alloys. Magnesium has the most negative potential and can provide
highest current output and hence is widely used in high resistivity electrolytes like soil.

Applications:
Protection of buried pipelines, underground cables from soil corrosion.
Protection from marine corrosion of cables, ship hulls, piers etc.
Insertion of magnesium sheets into the domestic water boilers to prevent the formation of rust.
Calcium metal is employed to minimize engine corrosion.

Advantages:
a. Low installation and operating cost.
b. Capacity to protect complex structures.
c. Applied to a wide range of severe corrodents.

Limitations:
a. High starting current is required.
b. Uncoated parts cannot be protected.
c. Limited driving potential, hence, not applicable for large objects.
 Sacrificial anodic protection method

(b) Impressed current cathodic protection method:

In this method, an impressed current is applied in the opposite direction to nullify the corrosion current
and convert the corroding metal from anode to cathode.
Usually the impressed current is derived from a direct current source (like battery or rectifier on AC line)
with an insoluble, inert anode (like graphite, scrap iron, stainless steel, platinum or high silica iron).
A sufficient DC current is applied to an inert anode, buried in the soil (or immersed in the corroding
medium) and connected to the metallic structure to be protected. The anode is, usually, a backfill,
composed of coke breeze or gypsum, so as to increase the electrical contact with the surrounding soil.
Impressed current cathodic protection has been applied to open water box coolers, water tanks, buried
oil or water pipes, condensers, transmission line towers, marine piers, laid up ships etc.
This kind of protection technique is particularly useful for large structures for long term operations.
 Impressed current cathodic protection method

Comparison of Sacrificial anode method with Impressed current cathodic

method:

Sl. Sacrificial Anode method Impressed Current method

No.
1 External power supply is not required. External power supply is required.

2 The cost of investment is low. The cost of investment is high.

3 This requires periodic replacement Replacement is not required as anodes

sacrificial anode. are stable.

4 Soil and microbiological corrosion effects are not Soil and microbiological corrosion
considered. effects are taken into account.

5 This is the most economical method especially when This is well suited for large structures
short term protection is required. and long term operations.

6 This is a suitable method when the current This is a suitable method even when the
requirement and the resistivity of the electrolytes are current requirement and the resistivity of
relatively low. the electrolytes are high.
4. Modifying the environment-corrosion control:

Environment plays a major role in the corrosion of metals. Hence, we can prevent corrosion to a great
extent by modifying the environment. Some of the methods are

i) Deaeration:
Fresh water contains dissolved oxygen. The presence of an increased amount of oxygen is harmful and
increases the corrosion rate. Deaeration involves the removal of dissolved oxygen by increase of
temperature together with mechanical agitation. It also removes dissolved carbon dioxide in water.

ii)By using inhibitors:

Inhibitors are organic or inorganic substances which decrease the rate of corrosion. Usually the inhibitors
are added in small quantities to the corrosive medium. Inhibitors are classified into
1) Anodic inhibitors (chemical passivators)
2) Cathodic inhibitors (adsorption inhibitors)
3) Vapour phase inhibitors (volatile corrosion inhibitors)

Anodic Inhibitors:

Inhibitors which retard the corrosion of metals by forming a sparingly soluble compound with newly
produced metal cations. This compound will then adsorb on the corroding metal surface forming a
passive film or barrier. Anodic inhibitors are used to repair
1.the crack of the oxide film over the metal surface
2.the pitting corrosion
3.the porous oxide film formed on the metal surface. Examples: Chromate, phosphate, tungstate, nitrate,
molybdate etc.

Cathodic Inhibitors:

Depending on the nature of the cathodic reaction in an electrochemical corrosion, cathodic inhibitors are
classified into

(a) In an acidic solution:

The main cathodic reaction is the liberation of hydrogen gas, the corrosion can be controlled by slowing
down the diffusion of H+ ions through the cathode. Eg., Amines, Mercaptans, Thiourea etc.
2H+ + 2 e- -------- H2

(b) In a neutral solution:

In a neutral solution, the cathodic reaction is the adsorption of oxygen or formation of hydroxyl ions.
The corrosion is therefore controlled either by eliminating oxygen from the corroding medium or by
retarding its diffusion to the cathodic area.
The dissolved oxygen can be eliminated by adding reducing agents like Na2SO3.
The diffusion of oxygen can be controlled by adding inhibitors like M g, Zn or Ni salts. Eg., Na2SO3,
N2H4, Salts of Mg, Zn or Ni.

½ O2 + H2O + 2e- ---------- 2OH-

Vapour phase inhibitors:

These are organic inhibitors which are readily vapourised and form a protective layer on the metal
surface.
These are conveniently used to prevent corrosion in closed spaces, storage containers, packing materials,
sophisticated equipment etc.
Examples are Dicyclohexylammonium nitrate, dicyclohexyl ammonium chromate, benzotriazole,
phenylthiourea etc.

5. Anodic protection

This is an electrochemical method of corrosion control in which an external potential control system,
called potentiostat, is used to produce and maintain a thin non corroding, passive film on a metal or an
alloy. The use of potentiostat is to shift corrosion potential into passive potential so that the corrosion of
the metal is stopped.
The potential of the object (say acid storage tank) to be protected is controlled by a potential controller
(potentiostat) so that under a certain potential range, the object becomes passive and prevents further
corrosion. This potential range depends upon the relationship between the metal and the environment.
Applications:
1.Used in acid coolers in dilute sulphuric acid plants
2.used in storage tanks for sulphuric acid
 3.used in chromium in contact with hydrofluoric acid
Limitations:
This method cannot be applied in the case of corrosive medium containing aggressive chloride.This
cannot be applied if protection breaks down at any point, it is difficult to reestablish.

Protective Coatings:
In order to protect metals from corrosion, it is necessary to cover the surface by means of protective
coatings. These coatings act as a physical barrier between the coated metal surface and the environment.
They afford decorative appeal and impart special properties like hardness, oxidation resistance and
thermal insulation.

Classification:
Protective coatings can be broadly classified into two types. They are
1. Metallic coatings
2. Non-Metallic coatings

Metallic Coating:
Corrosion of metals can be prevented or controlled by using methods like galvanization, tinning, metal
cladding, electroplating, cementation, anodizing, phosphate coating, enamelling, electroless plating. Some
of the methods are

1) Hot dipping: It is used for producing a coating of low-melting metals such as Zn (m.p.=419 deg C), Sn
(m.p.=232 deg C), Pb, Al etc., on iron, steel and copper which have relatively higher melting points. The
process involves immersing the base metal in a bath of the molten coating-metal, covered by a molten flux
layer (usually zinc chloride).

2) Galvanizing: It is the process of coating iron or steel sheets with a thin coat of zinc to prevent them
from rusting. The process of iron or steel articles is first cleaned with dil. Sulphuric acid and washed with
distilled water and dried. The dried metal is dipped in a bath of molten zinc, now the thin layer of zinc is
coated on the iron or steel article.

3) Metal cladding: It is the process by which a dense, homogeneous layer of coating metal is bonded
firmly and permanently to the base metal on one or both sides. Corrosion resistant metals like nickel,
copper, 24 lead, silver, platinum and alloys like SS, nickel alloys, copper alloys, lead alloys can be used as
cladding materials.
 4) Tinning: It is a method of coating tin over iron or steel articles. The process is first treating steel sheet
in dilute sulphuric acid and it is passed through a flux (ZnCl2), next steel passes through a tank of molten
tin and finally through a series of rollers from underneath (bottom of) the surface of a layer of palm oil.

5) Electroplating: Electroplating is a coating technique. It is the most important and most frequently
applied industrial method of producing metallic coating. Electroplating is the process by which the coating
metal is deposited on the base metal by passing a direct current through an electrolytic solution containing
the soluble salt of the coating metal. The base metal to be plated is made cathode whereas the anode is
either made of the coating metal itself or an inert material of good electrical conductivity (like graphite).
Objectives: Electroplating is carried out for 1) Decoration or better appearance 2) Increasing the resistance
to corrosion of the coated metal. 3) Improving the hardness of the metal 4) Increasing the resistance to
chemical attack.

Non-Metallic coatings:
The non metallic coatings may be organic or inorganic .

Organic coatings consist of Paints, Varnishes, Lacquers and Enamels.

Inorganic coatings consist of Surface coating or chemical conversion coating – Chromate coating,
Phosphate coating and Oxide coating. 2. Anodising 3. Enamel coating or Vitreous or Porcelain coating.

Applications:

Corrosion of metal products can seriously affect their reliability, enormous wastage of machinery ,
unpredictable machinery failure which may lead to accidents causing loss of life, the efficiency of the
machine decreases and frequent replacement of corroded equipment increases cost.
To minimize corrosion of a multi-metal assembly, it is necessary to understand the basic corrosion
processes that produce certain types of corrosion and recognize the factors that affect the rate of corrosion.
Once this is understood, the information is applied to all phases of design engineering, including materials
selection, methods of prevention, and physical design. The final step must be the determination of the
feasibility of the proposed engineering design.

Questions:

1. Define the term corrosion and differentiate it from erosion

2. Discuss the different types of corrosion that we commonly come across.

3. Discuss the importance of design and material selection in controlling corrosion.

4. State two conditions for electrochemical corrosion to occur

5. Indicate the principles of cathodic protection?

6. Discuss about the use of inhibitors in corrosion control?

7. What is cathodic protection? Under what conditions is this protection more useful?

8. Explain with suitable examples the corrosion due to differential aeration and dry corrosion?

9. Explain the two important factors that influence the corrosion of metals?
10. Mention the theories of corrosion and explain any one of them in detail.

13. What is meant by electrochemical corrosion?

14. What is the mechanism by which rusting occurs?

15. Distinguish between wet and dry corrosion

16. A copper equipment should not possess a steel component. Why?

17. State Pilling Bedworth rule.

18. How does the corrosion product influence further corrosion?

19. What are the factors that influence corrosion?

20. How is the metallic surface prepared by various methods before electroplating?

21. Explain why magnesium corrodes faster when it is in contact with copper than when it is in contact
with iron.

22. How is Ni plating done by electroless plating?

23. Explain the control of corrosion by the use of sacrificial anodes and by impressed current cathodic
protection.

25. Bolt and nut made of the same metal is preferred in practice. Why?

26. What is cathodic protection?

27. Discuss the mechanism of electrochemical corrosion.

29. Write in detail about the galvanic corrosion and differential aeration corrosion.

30. What are corrosion inhibitors? Classify different types of inhibitors with examples.

31. Describe briefly the important parameters involved in electroplating.

32. What is rust?

33. Why does corrosion of water filled steel tanks occur below the water line?

34. Discuss the mechanism of chemical and electrochemical corrosion.

35. Write a note on pitting corrosion, stress corrosion and cathodic protection.

36. What is a sacrificial anode? Give two examples.

38. What is pitting corrosion?

39. Write a brief note on Differential aeration corrosion.

40. Discuss how the nature of the metal influences the rate of corrosion.
41. Explain the process of electroplating with a suitable example. Mention the uses of electroplating.

42. Explain electrochemical theory of corrosion with suitable examples.

43. What are the conditions for electrochemical corrosion?

44. How is pilling Bedworth ratio related to the protective capacity of an oxide layer?

45. What is a sacrificial anode? How does it protect a submerged pipeline?

47. What are the factors affecting electroplating?

52. What are the limitations of electroplating?

55. How can galvanic corrosion be prevented?

57. What type of corrosion will occur on a piece of iron covered with dust?

58. Why does any impure metal corrode faster than pure metal under identical conditions?

59. What are the drawbacks of cathodic protection?

60. Will a piece of steel covered with ice undergo corrosion? Why?

Assignment:

Q1.Explain electrochemical and chemical corrosion and their mechanisms.

Q2. What is passivity ? Explain its significance.

Q3. What is Pilling-Bedworths rule?

Q4. What are the various environmental factors which affect corrosion?

Q5. Explain any three methods of protection against corrosion.

Q6. Compare the Sacrificial anode protection method with Impressed current cathodic protection
method.

Q7. What are the consequences of corrosion in our everyday life ?

Q8. Explain why

a) Copper equipment should not possess a small steel bolt.

b) Bolt and nut made of the same material is preferred.

c) Impurities increase the rate of corrosion.

d) Galvanic series is a preferred way of predicting the corrosion of metals as compared to

electrochemical series.

e) Rusting of iron is quicker in saline water than fresh water.

Q9. Write short notes on the following

a) hydrogen embrittlement

b) decarburation

c) liquid metal corrosion

d) water line corrosion

e) crevice corrosion

References:

1) Jain and Jain, Engineering Chemistry, 15th Edition, Dhanpat Rai Publishing Co., New Delhi.

2) S.S. Dara, Engineering Chemistry, 1st Edition, S. Chand & Co, New Delhi

3) https://onlinelibrary.wiley.com/doi/book/10.1002/9780470277270
4) https://www.researchgate.net/publication/275028997_Corrosion_and_Corrosion_Control