A brief introductory write-up on corrosion and its prevention.

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CORROSION & CORROSION PREVENTION
Basic Principles
Dr. T. K. G. Namboodhiri
Consultant-Metallurgy & Corrosion, Tiruvalla, Kerala
(Ex-Professor of Metallurgical Engineering, Banaras Hindu University)
1. INTRODUCTION
1.1 What is corrosion?
Corrosion may be defined as the destruction or deterioration in properties of materials by
interaction with their environments. It is a natural phenomenon. Engineers generally
consider corrosion when dealing with metallic materials. However, the process affects all
sorts of materials, for example, ceramics, plastics, rubber etc. Rusting of iron and steel is the
most common example of corrosion. Swelling in plastics, hardening of rubber, deterioration
of paint, and fluxing of the ceramic lining of a furnace are all incidences of corrosion in non
metallic materials. Metallurgists may think of corrosion as reverse extractive metallurgy.
Metals are extracted from their compounds occurring in nature through extractive metallurgy
processes involving considerable expenditure of energy, natural resources, time, and man
power. Corrosion works to convert the metal I back into the same compounds.
1.2. Why is corrosion important?
Corrosion is a destructive natural process. It causes huge losses of money, material, and
natural resources. Estimated annual loss due to corrosion ranges from 3.5 to 5 % of the GNP
of a nation. Any effort to reduce this huge loss will be of much advantage to the economy.
Hence, it is advisable that all, particularly those involved with engineering materials, are
aware of the basic principles of corrosion and corrosion prevention.
1.3. Cost of corrosion
Loss due to corrosion may be direct or indirect.
Direct costs: cost of material lost, cost of repair and replacement of corroded parts, cost of
painting & other protective measures, over-design to allow for corrosion, and inability to use
otherwise suitable & cheaper materials.
Indirect costs: May be economical or social in nature. These include, contamination of
products like food items or drugs, loss of valuable products from leaking tanks or pipes, loss
of production due to shut downs, loss in appearance, as in automobiles or homes, and loss in
safety reliability of structures, machines, pipelines, storage tanks etc.
As per 2004 estimates, the annual direct loss due to corrosion in India was Rs. 36,000 crores,
while in the USA the loss was $364 billion. If the indirect costs are also taken into account,
this figure will be several times more. The higher the state of industrialization of a country,

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the higher will be the loss due to corrosion. India with its hot and humid tropical climate will
experience a larger proportion of corrosion loss than a temperate country like the U.K. or
France. A developing country like India, where a considerable percentage of population lives
under the poverty line, can ill afford the loss of such huge amounts due to corrosion. Hence,
corrosion should be dealt with all the seriousness it deserves.
2. PRINCIPLES OF CORROSION
Why do metals corrode?
Every system in the universe tries to reduce its energy content so as to become stable. This is
true for all types of reactions we see in nature. A spontaneous chemical reaction will occur
only if it leads to a lowering of the total energy content of the system. In thermodynamics we
say that all spontaneous reactions are accompanied by a lowering of the free energy. Most
metals and alloys, except the few noble metals, have higher free energies than those of their
chemical compounds. This is the reason they are not seen in nature as native metals.
Metallurgists spend lot of energy to convert metal compounds into metals, which thus
become unstable. As soon as they are put to service, they tend to get converted into their
more stable compound form. This is the basis of metallic corrosion.
Corrosion is an interdisciplinary phenomenon. It involves principles of thermodynamics,
electrochemistry, metallurgy, and physics & chemistry.
2.1 Thermodynamics of corrosion
Thermodynamics deals with the energy content of metals. The free energy of formation of a
compound is a measure of its energy content and stability. A compound with a high negative
free energy is very stable and requires a high energy input to convert it to the metal. This
metal will have a high tendency to be converted back to the compound, so as to reduce its
energy content, and so a high tendency for corrosion. The free energy change associated with
the compound formation is thus indicative of the tendency for corrosion of the metal. As is
discussed in the next section, the free energy change is used to calculate the electrode
potentials of metals and their corrosivity. Potential – pH diagrams or Pourbaix diagrams are
used to predict the corrosion tendency of a metal electrolyte system at various conditions.
2.2 Electrochemical nature of corrosion
Metallic corrosion is essentially electrochemical in nature. An electrochemical reaction is a
chemical reaction where, in addition to mass transfer, electron transfer also takes place. In
order to understand electrochemistry, we must understand what an electrode means. A
metallic conductor in contact with an ionic conductor (electrolyte) is called an electrode. Any
electrode will have a stable electrode potential, which develops due to the electrochemical
reactions taking place at the interface, and is the difference in potential of the metal and
electrolyte surrounding it. The electrode potential developed under standard conditions, the
standard electrode potential, is a characteristic property of the electrode.

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Electrochemical reactions take place in electrolytic cells, which consist of two different
electrodes immersed in an electrolyte
and connected electronically outside
the cell, as in Fig. 1.
Let us consider the corrosion of Zn
in hydrochloric acid. This process
can be represented by the electrolytic
cell shown in Fig. 1. We have Zn
metal as one electrode, and the inert
pt as the second electrode. When Zn
comes in contact with HCl, Zn atoms
dissolve, as per the following
equations.
DISSOLUTION OF ZN METAL IN HYDROCHLORIC ACID,
Zn + 2 HCl = ZnCl 2 + H 2 -------------------- -(1)
Written in ionic form as,
Zn + 2 H + + 2Cl − = Zn 2 + + 2Cl − + H 2 ----------------------(2)
The net reaction being,
Zn + 2 H + = Zn 2 + + H 2 ------------------------- (3)
Equation (3) is the summation of two partial reactions,
Zn → Zn 2* + 2e -----------------------------------------(4) and
2 H + + 2e → H 2 ------------------------------------------(5)
Equation (4) is the oxidation / anodic reaction and
Equation (5) is the reduction / cathodic reaction
The Zn atoms get converted to Zn ions by the oxidation reaction (4) at the Zn electrode and
get into the electrolyte, leaving behind two electrons/atom, which travel through the external
circuit to the Pt electrode and reduce two hydrogen ions to a hydrogen molecule by the
reaction (5). Thus one Zn atom dissolves while one molecule of hydrogen is liberated at the

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Pt electrode. The metal continues to dissolve liberating hydrogen molecules continuously.
Such an electrochemical reaction leads not only to a mass transfer from the metal to the
electrolyte, but also electron transfer from one electrode to the other. He electrode on which
oxidation takes place is called the anode, while that on which reduction occurs is a cathode.
Every electrochemical reaction thus has two components, one oxidation, and the second
reduction. These two reactions occur simultaneously at the same rate so that there is a charge
balance. The corrosion of a piece of Zn metal in HCl is shown schematically in Fig.2.
Here both the
anodic and
cathodic
reactions take
place on the
same piece, at
different
locations. The
metal dissolves
as ions and the
electrons left
behind move to
another point on
the metal
surface where
they reduce
hydrogen ions
from the
electrolyte to
hydrogen
atoms..
Fig. 2 Corrosion of Zn in HCl
As mentioned before, each electrode has a characteristic electrode potential, called the single
electrode potential. This potential is related to the nature of the electrode reaction taking
place, and thus to the free energy change involved in the reaction. The relationship may be
given as,
∆G = -nFE --------------------------------------- (6)
Where,
∆G is the free energy change in joules
n is the number of electrons involved in the reaction
E is the electrode potential in volts
F is the Faraday constant, 96500 Coulombs/ g.equivalent.

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The single electrode potential of an electrode when all the reactants are at unit activity and at
250 C is called the standard electrode potential (E0). These potentials are also referred to as
redox potentials because they represent the equilibrium between an oxidation and a reduction
reactions taking place at the electrode interface.
EMF series & Galvanic series
Standard electrode potentials of many common elements are tabulated as an EMF series
which is used by electrochemists to determine the possible direction of reactions. Corrosion
engineers use another series, the galvanic series where metals are listed in the order of their
electrode potentials measured under actual service conditions. Galvanic series predicts
corrosion reactions more accurately.
Polarization
When corrosion takes place on a metal, its electrode potential shifts away from the standard
electrode potential, according to the equation,
E = E0 +2,303 RT/nF (product of activities of reactants/ product of activities of products)
This shifting of the potential from the standard value is called polarization, which forms the
basis of the kinetics of corrosion reactions.
Kinetics of corrosion
While thermodynamics predict the possible direction of a reaction, it cannot predict the rate
at which the reaction will occur. For this we require kinetics. The single electrode potential
of an electrode gets polarized when an electric current is flowing, ie, when corrosion takes
place. The rate at which corrosion occurs is determined by the Mixed Potential Theory of
corrosion.
Mixed Potential Theory
The mixed potential theory of Wagner
and Traud helps us to determine the
kinetic parameters of electrochemical
corrosion. It consists of two simple
hypotheses, 1) any electrochemical
reaction can be split into two or more
partial oxidation and reduction reactions,
and 2) there can be no net accumulation
of electrical charge during an
electrochemical reaction. Accordingly, a
corroding metal cannot spontaneously
accumulate electrical charge. All the
Fig. 3 Mixed potential theory

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electrons released by the anodic oxidation reaction must be consumed simultaneously by one
or more cathodic reduction reactions. The potential of a corroding metal will be determined
by the partial oxidation and reduction reactions involved in the process. This is schematically
shown in Fig.3.
2.3 Metallurgy of corrosion
Metals and alloys are widely used in engineering applications and are the ones which
undergo corrosion in service. The nature and extent of corrosion strongly depend upon the
metallurgical characteristics of the material. The tendency for corrosion of a metal depends
primarily on its electrode potential and polarization behavior. The chemical composition of
the material determines the electrode potential. Crystal structure of the metal may influence
corrosion. Metals are made up of tiny crystals called grains and many properties of the
material depend upon the grain size. The grain boundary, the region between adjacent grains,
is a defective region where impurities accumulate. Besides grains, engineering materials
generally contain different phases in their structure. These phases will have different
chemical compositions and hence different corrosion tendencies. Metals may also contain
defects like dislocations, internal surfaces, inclusions and voids. All of these may affect the
corrosion behavior of the material. Mechanical properties like strength and ductility have
much influence on certain forms of corrosion.
2.4 Physics & Chemistry in corrosion
Materials undergo corrosion during service because they are exposed to corrosive
environments. Different environments have different corrosive properties. Hence, chemistry
of the environment is a very important factor in corrosion. Physical properties like density,
viscosity, melting point and boiling point, surface tension of the environment as well as the
corroding material may affect the corrosion behavior. Velocity of liquid media will have a
great influence on the extent of corrosion. Temperature and pressure are two other important
variables in corrosion.
3. FORMS OF CORROSION
The process of corrosion may be classified in different ways. Some of these classifications
are given below.
A) Wet corrosion and Dry corrosion: Wet corrosion is corrosion in presence of water or a
liquid corrodent. Dry corrosion occurs in presence of gaseous atmospheres or in
contact with solids.
B) Room Temperature corrosion and High Temperature corrosion: Generally all wet
corrosion takes place at or around room temperature. High temperature corrosion
occurs generally much above the boiling point of water or other aqueous corrodents,
and is a dry corrosion processes.
C) Electrochemical corrosion and Chemical corrosion: Metallic corrosion is always
electrochemical in nature. Dissolution of a metal in an acid was previously thought to
be a simple chemical reaction and was called chemical corrosion. But now it is also
seen as an electrochemical corrosion.

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the original Fontana classification of corrosion based on the appearance of the corroded
metal. . Here we have 8 forms of room temperature or aqueous corrosion, and oxidation
and corrosion under complex gaseous environments at high temperatures.
Room Temperature or Aqueous Corrosion
Based on the appearance of the corroded metal, wet corrosion may be classified as
• Uniform or General
• Galvanic or Two-metal
• Crevice
• Pitting
Dealloying
• For the purpose of this lecture, let us consider
• Intergranular
• Velocity-assisted
• Environment-assisted cracking
UNIFORM CORROSION
• Corrosion over the entire exposed surface at a uniform rate. e.g.. Atmospheric
corrosion.
• Maximum metal loss by this form.
• Not dangerous, rate can be measured in the laboratory
GALVANIC CORROSION
• When two dissimilar metals are joined together and exposed, the more active of
the two metals corrode faster and the nobler metal is protected. This excess
corrosion is due to the galvanic current generated at the junction
CREVICE CORROSION
• Intensive localized corrosion within crevices & shielded areas on metal surfaces
• Small volumes of stagnant corrosive caused by holes, gaskets, surface deposits,
lap joints
PITTING

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• A form of extremely localized attack causing holes in the metal
• Most destructive form
• Autocatalytic nature
• Difficult to detect and measure
• Mechanism
Fig.4 shows the mechanism of pitting.
Fig. 4 Mechanism of Pitting.
DEALLOYING
• Alloys exposed to corrosives experience selective leaching out of the more active
constituent. e.g. Dezincification of brass.
• Loss of structural stability and mechanical strength
INTERGRANULAR CORROSION
• The grain boundaries in metals are more active than the grains because of
segregation of impurities and depletion of protective elements. So preferential
attack along grain boundaries occurs. e.g. weld decay in stainless steels
VELOCITY ASSISTED CORROSION
• Fast moving corrosives cause
• a) Erosion-Corrosion,
• b) Impingement attack , and
• c) Cavitation damage in metals
CAVITATION DAMAGE
• Cavitation is a special case of Erosion-corrosion.
• In high velocity systems, local pressure reductions create water vapour bubbles
which get attached to the metal surface and burst at increased pressure, causing
metal damage
ENVIRONMENT ASSISTED CRACKING
• When a metal is subjected to a tensile stress and a corrosive medium, it may
experience Environment Assisted Cracking. Four types:
• Stress Corrosion Cracking
• Hydrogen Embrittlement
• Liquid Metal Embrittlement

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• Corrosion Fatigue
STRESS CORROSION CRACKING
• Static tensile stress and specific environments produce cracking
• Examples:
• 1) Stainless steels in hot chloride
• 2) Ti alloys in nitrogen tetroxide
• 3) Brass in ammonia
HYDROGEN EMBRITTLEMENT
• High strength materials stressed in presence of hydrogen crack at reduced stress
levels.
• Hydrogen may be dissolved in the metal or present as a gas outside.
• Only ppm levels of H needed
LIQUID METAL EMBRITTLEMENT
• Certain metals like
Al and stainless
steels undergo
brittle failure when
stressed in contact
with liquid metals
like Hg, Zn, Sn, Pb
Cd etc.
• Molten metal
atoms penetrate
the grain
boundaries and
fracture the metal
•
•
Fig 5 a). Tensile behavior under LME
•
Fig. 5 b). Brittle IG
fracture in Al alloy
by Pb
CORROSION
FATIGUE:
S-N DIAGRAM

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Fig. 6a) gives schematic S-N curves for fatigue and corrosion-fatigue.
 Synergistic action of corrosion
& cyclic stress. Both crack
Stress Amplitude
nucleation and propagation
are accelerated by corrodent
and the S-N diagram is Air
shifted to the left
Corrosion
log (cycles to failure, Nf)
Fig. 6a) S-N curves for fatigue and corrosion fatigue
CRACK PROPAGATION
Fig. 6b) shows schematic crack
log (Crack Growth Rate, da/dN)
propagation curves under fatigue as well
as corrosion-fatigue conditions.
 Crack propagation rate is
increased by the corrosive action
Log (Stress Intensity Factor Range, K

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Fig. 6b) Crack propagation rates for fatigue and corrosion-fatigue
High Temperature or Dry Corrosion
Oxidation under dry conditions, high temperature corrosion reactions in gaseous
atmospheres, and hot corrosion come under this classification.
OXIDATION
Oxidation refers to the reaction between a metal and air or oxygen in the absence of water or
an aqueous phase. Scaling, tarnishing and dry corrosion are other names for this process.
Nearly all metallic materials react with oxygen at high temperatures. As the temperature
increases, the oxidation resistance of materials decreases. As there are many applications of
metals at high temperatures, like gas turbines, rocket engines, refineries, and furnaces, the
importance of high temperature oxidation is considerable.
The oxidation resistance of a material may be related to the relative volumes of the metal and
its oxide, through the Pilling-Bedworth ratio, R = Md / nmD, where, M is the molecular
weight of the scale, D is the density of the scale, m is the atomic weight of the metal, d is the
density of the metal, and n is the number of metal atoms in a molecular formula of the scale.
R gives the volume of oxide formed from a unit volume of the metal. At a ratio of less than 1,
the scale does not cover the metal completely, and the metal continues to get oxidized, while
a ratio much greater than one tends to produce too much oxide which introduces high
compressive stress and tendency for spalling of the scale. The ideal R value is close to one.
Oxidation, like aqueous corrosion is an electrochemical process, and consists of two partial
processes,
M → M +2 + 2 e- ----------- Metal oxidation at metal-scale interface
½ O2 + 2 e- → O2 --------- Oxygen reduction at scale-gas interface.
----------------------

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M + ½ O2 → MO --------------------Overall reaction
The oxide scale acts as the electrolyte through which ions and electrons move to make the
above reactions possible. The electronic and ionic conductivities of the scale thus determine
the rate of oxidation of the metal.
Oxidation kinetics
When a fresh metal is
exposed to oxygen, a
thin surface layer of the
oxide forms on the
metal. As oxidation
continues, the scale
thickness increases and
the reaction rate
decreases depending
upon the scale
characteristics. Many
empirical rate equations
have been developed to
fit experimental
oxidation data. Some of
these are linear,
parabolic, logarithmic
and cubic. These are
shown schematically in Fig. 7. Fig.7 Oxidation Rate Laws
Oxidation-resistant alloys
The oxide characteristics determine the oxidation resistance of an alloy. Most oxides are non-
stoichiometric compounds with structural defects. They may be n-type or p-type
semiconductors whose conductivities could be altered by alloy additions. This principle is
used in developing high temperature oxidation resistant alloys like Fe-Cr, Fe-Cr-Al, and Ni-
base alloys.
Catastrophic oxidation
Metals that follow linear oxidation kinetics at low temperatures may experience oxidation at
continuously increasing rates at high temperatures. Metals like Mo, W, Os, Rh, and V which
have volatile oxides may oxidize catastrophically. Alloys containing Mo and V may oxidize
catastrophically by the formation of low melting eutectic oxide mixtures. Combustion of fuel
oils with high V compounds produces vanadium oxides in the gas phase, and can lead to
catastrophic oxidation.
Internal Oxidation
In some alloys, one or more dilute components may form more stable oxides than the base
metal which get distributed below the metal-oxide interface. This is called internal oxidation

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because the oxide precipitate forms within the metal matrix. Dilute copper and silver based
alloys containing Al, Zn, Cd, Be show such a behavior.
CORROSION IN OTHER GASEOUS ENVIRONMENTS
Sulfur compounds
High temperature degradation of metals occurs when exposed to sulfur compounds like H2S,
SO2 and vaporized sulfur. This process is referred to as sulfidation. In reducing gases
containing hydrogen, such as gasified coal, H2S is a major gaseous constituent. In oxidizing
gases such as fossil fuel combustion products, considerable SO2 may exist. These sulfur
bearing gaseous compounds can lead to rapid scaling and to internal precipitation of stable
sulfides. Mechanical properties of high temperature alloys are seriously affected by these
precipitates.
Decarburization and hydrogen attack.
When steels are exposed o hydrogen at high temperatures, the carbon present either in
dissolved form or s carbides, reacts with hydrogen to produce methane gas, as per the
following reaction.
C (Fe) + 4 H→CH4
This phenomenon is called hydrogen attack. Decarburization leads to a decrease in the
strength of the steel. The methane formed inside the steel may lead to cracking. Cr and Mo
additions to steels improve their resistance to decarburization and cracking. A Nelson
diagram is used to predict safe working conditions of hydrogen partial pressure and
temperature for various steels.
Hot Corrosion
Hot corrosion refers to the accelerated high temperature corrosion of materials under sulfur
gaseous atmospheres and the presence fused sulphate compounds on the metal surface.
4) CORROSION TESTING.
Corrosion tests are of four types;
1. Laboratory tests
2. Pilot-plant tests
3. Plant or actual service tests
4. Field tests
Laboratory tests use small specimens and small volumes of corrodents and actual conditions
are simulated as far as possible. These are most useful as screening tests to determine which
material warrants further studies.
Pilot plant or semi-works tests are made in a small-scale plant that essentially duplicates the
intended large-scale operation. This type of tests generally gives the best results.

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Actual plant tests are done when an operating plant is available. The purpose is in evaluating
better or more economical materials or in studying corrosion behavior of existing materials
as process conditions are changed.
Field tests are designed to obtain general corrosion information. Examples are atmospheric
exposure of a large number of specimens in racks at one or more geographical locations and
similar tests in soil or sea water.
Corrosion tests are essential for the following purposes;
1. Evaluation and selection of materials for a specific environment or a given
application.
2. Evaluation of materials as regards to their compatibility to various environments. The
information generated helps in the selection of materials for a specific application.
3. Control of corrosion resistance of the material or corrosiveness of the environment.
These are routine quality control tests.
4. Study of the mechanisms of corrosion or other research and development purposes.
These are specialized tests involving precise measurements and close control.
Corrosion test standards have been developed by many organizations like ASTM, NACE in
USA and similar ones in other countries. Corrosion data exists for a large proportion of
materials used in several environments which could be used by material selectors.
Corrosion Rate
The most basic corrosion property of a material is the rate at which material is lost due to
exposure to an environment. This is expressed as the corrosion rate, which is expressed in
two ways; 1) weight of material lost per unit surface area per unit time, and 2) the rate of
penetration or thinning down of a material. The common corrosion rate units are mdd
(mg/dm2/day) in the first category, and mpy (mils/year) in the second category.
Weight loss after immersion in the corrodent for a specific time is measured on a specimen of
known surface area and then the corrosion rate can be calculated as
R = KW/ATD
Where,
K is a constant for a specific unit of R
W is weight lost in gm
A is the surface area in sq.cm
T is time of exposure in hours
D is density in g/cu.cm
The constant K varies from unit to unit. For mdd, the value of K is 2.4 x 106D, and for mpy
the K value is 3.45x 106.
5) PROTECTION AGAINST CORROSION
Need for corrosion prevention
• The huge annual loss due to corrosion is a national waste and should be minimized
• Materials already exist which, if properly used, can eliminate 80 % of corrosion loss

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• Proper understanding of the basics of corrosion and incorporation in the initial design
of metallic structures is essential
Methods
• Material selection
• Improvements in material
• Design of structures
• Alteration of environment
• Cathodic & Anodic protection
• Coatings
Material Selection
• Most important method – select the appropriate metal or alloy.
• “Natural” metal-corrosive combinations like
• S. S.- Nitric acid, Ni & Ni alloys- Caustic
• Monel- HF, Hastelloys- Hot HCl
• Pb- Dil. Sulphuric acid, Sn- Distilled water
• Al- Atmosphere, Ti- hot oxidizers
• Ta- Ultimate resistance
Improvement of materials
1) Purification of metals- Al , Zr
2) Alloying with metals for:
• Making more noble, e.g. Pt in Ti
• Passivating, e.g. Cr in steel
• Inhibiting, e.g. As & Sb in brass
• Scavenging, e.g. Ti & Nb in S.S
• Improving other properties
Design of Structures
• Avoid sharp corners
• Complete draining of vessels
• No water retention
• Avoid sudden changes in section
• Avoid contact between dissimilar metals
• Weld rather than rivet
• Easy replacement of vulnerable parts
• Avoid excessive mechanical stress
Alteration of Environment

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• Lower temperature and velocity
• Remove oxygen/oxidizers
• Change concentration
• Add Inhibitors
– Adsorption type, e.g. Organic amines, azoles
– H evolution poisons, e.g. As & Sb
– Scavengers, e.g. Sodium sulfite & hydrazine
– Oxidizers, e.g. Chromates, nitrates, ferric salts
Cathodic & Anodic Protection
• Cathodic protection: Make the structure more cathodic by
– Use of sacrificial anodes
– Impressed currents
Used extensively to protect marine structures, underground pipelines, water heaters
and reinforcement bars in concrete
• Anodic protection: Make Passivating metal structures more anodic by impressed
potential. e.g. 316 s.s. pipe in sulfuric acid plants
Coatings
• Most popular method of corrosion protection
• Coatings are of various types:
– Metallic
– Inorganic like glass, porcelain and concrete
– Organic, paints, varnishes and lacquers
• Many methods of coating:
– Electro deposition
– Flame spraying
– Cladding
– Hot dipping
– Diffusion
– Vapour deposition
– Ion implantation
– Laser glazing
Surface Engineering
The process of altering the surface characteristics of materials is known as surface
engineering. Corrosion is a surface property and all the coating processes mentioned above
come under surface engineering. Besides corrosion, wear, fretting, fatigue etc are also
dependent on the surface characteristics of materials.
CONCLUSION

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• Corrosion is a natural degenerative process affecting metals, nonmetals and even
biological systems like the human body
• Corrosion of engineering materials lead to significant losses
• An understanding of the basic principles of corrosion and their application in the
design and maintenance of engineering systems result in reducing losses considerably
REFERENCES
1. Corrosion Engineering, Mars G. Fontana, 3rd Ed. McGraw-Hill International,
Singapore, 1987
2. Corrosion and Corrosion Control, Herbert H. Uhlig, 3rd Ed. John Wiley & Sons,
New York, 1985
3. Metals Handbook, 9th ed. Volume 13, Corrosion, ASM International, Metals Park,
Ohio, 1988