LECTURE NOTES

Unit- 3

Topic- Corrosion& its theories Subject Teacher: Dr. PrashantKriplani

Syllabus

Corrosion and its control: Definition and significance of corrosion, Mechanism of chemical (dry) and
electrochemical (wet) corrosion, galvanic corrosion, concentration corrosion and pitting corrosion.
Protection from corrosion; protective coatings-galvanization and tinning, cathodic protection,
sacrificial anode and modifications in design.

3.1 Introduction

Corrosion may be defined as a destructive phenomenon, chemical or electrochemical, which can

attack any metal or alloy through reaction by the surrounding environment and in extreme cases may
cause structural failure. In other words Corrosion is the spontaneous decay of metal due to action of
moisture, chemical and other environmental factors on it.

3.2 Importance of Corrosion Studies

The importance of corrosion studies is two folds. The first is economic, including the reduction of
material losses resulting from the wasting away or sudden failure of piping, tanks, metal components
of machines, ships, hulls, marine, structures…etc. The second is conservation, applied primarily to
metal resources, the world’s supply of which is limited, and the wastage of which includes
corresponding losses of energy and water resources accompanying the production and fabrication of
metal structures.

3.3- Basic Causes of Corrosion

For the purpose of electrochemical corrosion which is the most important classification of
corrosion,Four conditions must exist before electrochemical corrosion can proceed:

3.3.1- There must be something that corrodes (the metal anode).

3.3.2- There must be a cathode.

3.3.3- There must be continuous conductive liquid path (electrolyte, usually condensate and salt or
other contaminations).

3.3.4- There must be a conductor to carry the flow of electrons from the anode to the cathode. This
conductor is usually in the form of metal-to-metal contact such as in bolted or riveted joints. The
elimination of any one of the four conditions will stop corrosion.

3.3.5- Environmental factors: At normal atmospheric temperatures the moisture in the air is enough
to start corrosive action. Oxygen is essential for corrosion to occur in water at ambient temperatures.
Other factors that affect the tendency of a metal to corrode are: Acidity or alkalinity of the
conductive medium (pH factor), Stability of the corrosion products, Biological organisms (particularly
anaerobic bacteria), Variation in composition of the corrosive medium, Temperature etc.

3.4Types of Corrosion

There are Threetheories/types of corrosion:

(i) Acid/ Carbonate Corrosion theory

(ii) Dry/ Chemical Corrosion theory

(iii) Wet/ electrochemical corrosion theory

3.4.1Acid/ Carbonate Corrosion theory

CO2 is present in water as a dissolved gas under the high pressures common in underground oil and
gas reservoirs. In the dissolved state it forms carbonic acid. The primary material of construction for
pipelines in the oil and gas industry is mild steel, because of its price, strength and availability. Initial
action of corrosion occurs as a result of combined chemical action of water, oxygen and CO2.

CO2 + H2O ----------------- H2CO3

Fe + H2CO3 ----------------- FeCO3 + H2

4FeCO3 + 10H2O + O2 ---------------- 4Fe(OH)3 + 4H2CO3

Except CO2 other gases like H2S, SO2, SO3, Cl2 etc. also causes acid corrosion.

3.4.2Dry/ Chemical Corrosion theory

According to this theory, corrosion on the surface of a metal is due to direct reaction of atmospheric
gases like oxygen, halogens, oxides of sulphur, oxides of nitrogen, hydrogen sulphide and fumes of
chemicals with metal. The extent of corrosion of a particular metal depends on the chemical affinity
of the metal towards reactive gas.

Chemical Corrosion may be further classified into 3 subtypes:

(i) Oxidation corrosion (Reaction with oxygen): Some of the metals directly react with oxygen in the
absence of moisture. Alkali and alkaline earth metals react with oxygen at room temperature and
form corresponding oxides, while some metals react with oxygen at higher temperature. Metals like
Ag, Au and Pt are not oxidized as they are noble metals. During oxidation of a metal, metal oxide is
formed as a thin film on the metallic surface which protects the metal from further corrosion.
If diffusion of either oxygen or metal is across this layer, further corrosion is possible. Thus, the layer
of metal oxide plays an important role in the process of corrosion.

Oxides of Pb, Al and Sn arestable and hence inhibit further corrosion. They form a stable, tightly
adhering oxide film.

In case of porous oxide film, atmospheric gases pass through the pores and react with the metal and
the process of corrosion continues to occur till the entire metal is converted into oxide. Porous oxide
layer is formed by alkali and alkaline earth metals.

Molybdenum forms a volatile oxide film of MoO3 which accelerates corrosion. Au, Ag, Pt form
unstable oxide layer which decomposes soon after the formation, thereby preventing further
corrosion.
Pilling Bedworth Rule: If volume of metal oxide on the surface of a metal is more than or equal to the
volume of metal, the oxide layer will be protective.For example, for Al2O3, Fe, Ni, ZnW, Cr.

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

It will be non-protective if volume of oxide is less than volume of metal. For example, Na, Mg, Ca etc.

(ii) Corrosion by other gases such as Cl2, SO2, H2S, NOx: In dry atmosphere, these gases react with
metal and form corrosion products which may be protective or non-protective. Dry Cl2 reacts with
Ag and forms AgCl which is a protective layer, while SnCl4 is volatile. In petroleum industries at high
temperatures,H2S attacks steel forming FeS scale which is porous and interferes with normal
operations.

(iii) Liquid metal corrosion: In several industries, molten metal passes through metallic pipes and
causes corrosion due to dissolution or due to internal penetration. For example, liquid metal mercury
dissolves most metals by forming amalgams, thereby corroding them.

3.4.3 Wet or electrochemical theory of corrosion by taking rusting of iron as example

It is a common type of corrosion of metal in aqueous corrosive environment. This type of corrosion
occurs when the metal comes in contact with a conducting liquid or when two dissimilar metals are
immersed or dipped partly in a solution. According to this theory, there is the formation of a galvanic
cell on the surface of metals. Some parts of the metal surface act as anode and rest act as cathode.
The chemical in the environment and humidity acts as an electrolyte. Oxidation of anodic part takes
place and it results in corrosion at anode, while reduction takes place at cathode. The corrosion
product is formed on the surface of the metal between anode and cathode.
To understand the wet theory, let us take the example of corrosion of iron. Oxidation of metal takes
place at anode while the reduction process takes place at cathode. By taking rusting of iron as an
example, the reaction can be explained as that it may occur in two ways: (i) evolution of hydrogen
and (ii) absorption of oxygen.
At anode: oxidation occurs.
At cathode:
Case I: Evolution of H2
The hydrogen ions (H+ ) are formed due to the acidic environment and the following reaction occurs
in the absence of oxygen
2H+ + 2e ---------- H2 ↑ (reduction)
The overall reaction is Fe + 2H+ ---------- Fe+2 + H2
In this case, metals react in the acidic environment and are dissolved (undergo corrosion) to release
H2 gas. Allmetals above hydrogen in electrochemical series can show this type of corrosion. In
hydrogen evolution type ofcorrosion, anodic area is large as compared to its cathodic area

(Mechanism of wet corrosion by (a) hydrogen evolution and (b) oxygen absorption)

Case II: Absorption of O2

This type of corrosion takes place in neutral or basic medium in the presence of oxygen. The oxide of
iron covers the surfaceof the iron. The small scratch on the surface creates small anodic area and rest
of the surface acts as cathodic area. TheFollowing chemical reactions occur at anode and cathode.
At anode Fe -------------- Fe+++ 2e¯¯ (oxidation)

Ferric hydroxide is actually hydrated ferric oxide, Fe2O3.H2O, which is a yellowish rust. Anhydrous
magnetite, Fe3O4 [a mixture of (FeO + Fe2O3)], is also formed, which is brown-black in colour. It is
markable that the corrosion occurs at anode but the corrosion product is formed near cathode. It is
because of the rapid diffusion of Fe++ as compared to –OH. Hence corrosion occurs at anode, but rust
is deposited at or near cathode.

3.4.4 Differences between dry and wet corrosion

Dry corrosion Wet or electrochemical corrosion

Corrosion occurs in the absence of moisture. Corrosion occurs in presence of
conducting medium.
It involves direct attack of chemicals on the metal It involves formation of electrochemical
surface. cells.
The process is slow. It is a rapid process.
Corrosion products are produced at the Corrosion occurs at anode but rust is
site of corrosion. deposited at cathode.
The process of corrosion is uniform. It depends on the size of the anodic part
of metal.

3.5 Galvanic corrosion

This type of electrochemical corrosion is also called bimetallic corrosion. When two dissimilar
metals are connected and exposed to an electrolyte, they will form a galvanic cell. The anodic metal
will be oxidized and it will undergo corrosion. Zinc and copper metals connected with each other in
an electrolyte medium form a galvanic cell. Zinc acts as anode and undergoes corrosion while
cathode will be unaffected.
At anode: Zn-------------- Zn++ + 2e– [Oxidation] (Corrosion)
At cathode: Cu++ + 2e– ------------- Cu [Reduction] (unaffected)
Galvanic corrosion can be avoided by coupling metals close to the electrochemical series. Fixing
insulating material between two metals. By using larger anodic metal and smaller cathodic metal.
Example of galvanic corrosion: i) Steel screws in brass marine hardware, ii) steel pipe connected to
copper plumbing, iii) steel propeller shaft in bronze bearing, iv)zinc coating on mild steel, v) lead–tin
solder around copper wires.

3.6Pitting corrosion

Due to crack on the surface of a metal, local straining of metal, sliding under load, chemical attack,
there is formation of a local galvanic cell. The crack portion acts as anode and rest of the metal
surface acts as cathode. It is the anodic area which will be corroded and the formation of a pit is
observed. This type of corrosion is thus called pitting corrosion. Metals owing to their corrosion
resistance to their passive state show pitting and ultimately result in formation of passivity. Presence
of external impurities such as sand, dust, scale embedded on the surface of metals lead to pitting. For
example, stainless steel and aluminum show pitting in chloride solution.
3.7 Stress corrosion
In a metallic structure, if there is a portion under stress, it will act as anode and rest part of the
structure will act as cathode. It is now a galvanic system and hence anodic part which is small in area
will corrode more. Stress corrosions are observed in the following systems:
Caustic embrittlement is a type of stress corrosion occurring in steel tank (Boiler) at high temperature
and in alkaline medium. Boiler water has Na2CO3; it will be hydrolyzed at high temperature to give
NaOH. It flows into hair cracks and crevices. There it reacts with iron and forms Na 2FeO2 (sodium
ferroate) which decomposes to give Fe3O4 (ferroferric oxide) and NaOH.

3Na2FeO2 + 3H2O ------------ Fe3O4 + H2 + 6NaOH

NaOH thus formed further reacts with iron to cause corrosion. It is called caustic embrittlement.
Addition of Na2SO4to boiler water in addition to tannin and lignin to boiler water prevents caustic
cracking. By neutralization of excess of alkali with dilute acid (or) control of pH value caustic
embrittlement can be controlled.

3.8Concentration corrosion: Practically such type of cells/corrosion occur due to differences in the
environment surrounding the metal. The corrosion commonly occurs in localized areas where small
volume of stagnant solution exists.
3.8.1 Differential aeration corrosion: If a metal rod is dipped in an electrolyte, the portion dipped in
water is poor in oxygen concentration and works as anode which gets corroded and the portion above
water acts as cathode which is protected. The system will act as a concentration cell and the chemical
reactions for zinc dipped in water are given as:

Zn(OH)2 appears as corrosion products.

Note: Galvanic series is the series of metals that is made keeping in view the process of corrosion of a
metal in a particular atmosphere, i.e. sea water. In galvanic series, oxidation potential of metals is
arranged in the decreasing order of activity of a series of metals. The series is towards the increasing
noble nature.
More anodic: Mg, Mg alloys, Zn, Al, Cd, Fe, Pb, Sn, Ni–Mo–Fe alloys) Brasses, Cu, Ni, Cr–steel
alloy,Ag, Ti, Au, Pt towards noble nature.

3.9 FACTORS INFLUENCING CORROSION

Since corrosion is a process of destruction of metal surface by its environment, the two factors that
govern the corrosion process are:
(i) Metallic and (ii) Environmental.
3.9.1 Nature of metal: Different properties of a metal are responsible for corrosion. These properties
are given here.
(a) Position of metal in galvanic series: It decides the corrosion rate. A metal having higher position in
galvanic series undergoes corrosion when connected to another metal below it. Also, more difference
in the position of galvanic series will cause faster corrosion at anodic metal.
(b) Purity of metal: Pure metal resists corrosion, while impurities in a metal form a local galvanic cell
(metal as anode and impurity as cathode) and result in the corrosion of metal. Rate of corrosion
increases due to more exposure of impurities. For alloys the system is a homogeneous solid solution,
hence no local action and no corrosion.
(c) Relative areas of anode and cathode: Smaller the area of anode com- pared to cathode will lead to
faster corrosion of anode. It is because the corrosion current at anode and at cathode will be same.
But for small anodic area the current density will be large at anode and larger cathodic area will
demand more electron which will be fulfilled by fast reaction at anode (oxidation), i.e. rapid
corrosion.
(d) Physical state of the metal: Small granular metal will corrode faster than the larger one. Also the
type of structure
formed by a metal will have effect on the corrosion rate. A bent metal (stress) is rapidly corroded due
to stress.
(e) Nature of oxide film: An oxide film is formed by the reaction between metal and oxygen. If this
oxide film is porous
and oxygen can be diffused through it, more corrosion is expected (already shown in dry or chemical
corrosion). Also,if volume of metal oxide is more than the volume of metal (The specific volume
ratios of Ni, Cl, W are 1.6, 2.0 and 3.6) least corrosion or no further corrosion occurs.
(f) Volatility and solubility of corrosion product: In both the cases, the corrosion will be faster. MnO3,
SnCl4 are volatile, so faster is corrosion of Sn in chlorine atmosphere. In case of soluble corrosion
product, it will be enhanced by water and metal surface will be exposed for further corrosion.

3.9.2Effect of environment
The role of environment in the corrosion of a metal is very important. Environmental parameters like
temperature,
humidity, pH, etc. play impor- tant role. The effect is discussed here.
(a) Temperature: The rate of diffusion increases by rise in temperature, hence the rate of corrosion is
also increased. At higher temperature, passive metals also become active and undergo corrosion. But
higher temperature reduces the concentration of oxygen and hence corrosion is reduced (in case of
water where oxygen is dissolved).
(b) Humidity: In humidity, gases like CO2, SO2, NOx are dissolved which form electrolytes. It will
cause galvanic corrosion. Some oxides are water soluble, humidity washes away the corrosion
products and metal surface is further corroded. Other soluble corrosion products can also be washed
away by humidity, causing further corrosion.
(c) Impurity of atmosphere: Pollutants like H2S, CO2, SO2 and acid vaporscause more pollution where
they dissolve. In sea water (salty in nature which acts as an electrolyte) corrosion rate increases. Some
suspended particles are dissolved in humidity and form electrolyte which helps in corrosion.
(d) pH : pH value means concentration of H+ (acidic nature). In acidic medium (pH less than 7),
corrosion is faster. Also, in basic medium pH > 7, some metals such as Pb, Zn, Al, etc. form complexes
and hence they corrode.
Example: Zn corrodes minimum at pH 11, but at higher pH (more than 11) it corrodes faster.
(e) Conductance effect: Due to presence of salts and water in earth, it is of con- ducting nature. More
conductance leads to more stray current and hence fast corrosion. Dry sandy soil is less conducting
and hence less corrosion, while mineralized clay soil is more conducting hence more corrosion
occurs.
(f) Oxygen concentration and oxygen concentration cell: Oxygen is one of the important elements
responsible for corrosion. It forms oxides and hydroxides (in presence of H2O) on the surface of metal
as corrosion product. Oxygen concentration cell is formed on the surface of metal due to difference in
oxygen concentration (iron rod half dipped in water corrodes due to this effect). Dipped portion will
be anode and outer portion will be cathode.
3.10 PROTECTION FROM CORROSION
Due to corrosion, there is a great loss of material and money. Therefore, it is essential to protect
metals from corrosion. Since, there are two components involved in corrosion—the metal and
environment, both are considered in corrosion protection. Following methods have been adopted for
the protection of metal from corrosion.
(i) Proper designing of an object helps in prevention of corrosion.
(ii) Contact of two dissimilar metals must be avoided. If it is unavoidable, an-odic area should be very
large compared to cathodic area.
(iii) Two different metals used in the structure should be such that they are oc- cupying near
positions in galvanic series.
(iv) Putting an insulator between two metals resists corrosion.
(v) As far as possible, metal used in a structure should be extremely pure. Small amount of impurity
causes corrosion.
(vi) While using an alloy, it should be completely homogeneous.
(vii) Design or fabricate equipment or metal parts in such a manner that they have minimised sharp
edges and corners and also avoid, as for as possible, the crevices in joints, etc.
(viii) The modification of environment also helps in protection from corrosion. It includes:
(a) De-aeration removes oxygen by adjusting temperature and mechanical ageing.
(b) Deactivation involves addition of chemicals such as Na2SO3, NH2 – NH2 which absorb oxygen.

(c) Dehumidification of environment is done by adding alumina or silica gel. These chemicals absorb
humidity from metallic surface. In humidity, gases such as CO2, H2S, SO2 and HCl give acidic medium
responsible for corrosion. They are neutralized by NH3 or NaOH or lime.
(ix) There are some chemicals which reduce the rate of corrosion. These chemi- cals are called
corrosion inhibitors. They are of two types, anodic and cathodic. Phosphate, chromate, tungstate
protect anode. They form sparingly soluble products which are adsorbed on the surface of metal and
hence check corrosion. Cathodic protection is done by organic amines, mercap- tans, thiourea and
substituted urea. The above chemicals retard reduction reaction taking place at cathode. Also, by the
use of salts of Mg, Zn or Ni, the insoluble hydroxides of Mg, Zn, Ni are deposited preventing
corrosion.

(x) Methods of application of metal coatings: Metallic dipping causes a coat of metal over the base
metal. It is of two types, anodic coating and cathodic coating.
(a) Anodic coating is done by a metal which is situated prior to the base metal in electrochemical
series. Coating of zinc over iron is anodic coat- ing and this process has its special name galvanization.
Zinc occupies a position before iron. In this process base metal is immersed in molten coating metal.
The base metal should by very clear. Anodic coating of iron is done by Al and Cd also. If any crack or
pit is formed on the coated metal, Zn will act as anode and Fe as cathode and hence Fe will not
corrode .Since Zn is toxic, galvanized utensils are not used for cooking or food storing. Zine coating
protects iron sacrificially. Galvanized iron is used for buckets, tubes, wires, roof sheets etc.

Fig : Functioning of anodic coating. In galvanized steel, Zn serves as anode; while iron of steel serves as the cathode. Therefore, the iron is
protected, even if it is exposed, when a part of the zinc coating is scrapped off
b) Cathodic coating of a base metal (say iron) is done by another metal which occupies a position
below it in electrochemical series but it will be higher corrosion resistant. Coating of Sn over iron
(tinning) is an example of cathodic protection. Cathodic protection of iron is successful only when
there is no pit or crack formation on the surface of metal. Coating of tin is called tinning. Tin-coated
iron is used for cooking as well as storage purposes.
 Fig. Tin-plated steel. Tin protects the iron, when the coating is continuous. When the coating is broken,
the iron of the steel becomes the anode and is subjected to accelerated local corrosion.

Difference between Galvanization and Tinning

Galvanization Tinning
Galvanizing is the process of covering iron or In tinning, steel is covered with a thin coat of tin
steel with thin layer of Zn to prevent corrosion
Zinc protects iron sacrificially Due to noble nature tin protects base metal
Zinc protects iron even when coating of zinc is If the coating is puncture intense corrosion of steel
Punctured occurs.
Since zinc is toxic galvanized utensils are not Since tin is non-toxic tinned contain- ers can be
used for storing food used even for storing food.

(xi) Metal cladding is a process of sandwiching a metal between two corrosion resistant metals. Ni, Pb,
Cu, Al,Ag, Pt are used as cladding metals over a base metal. For example, aluminium acts as cladding
metal for duralumin.

The cladding of duralumin

(xii) Metal spraying: Molten metal is sprayed on the cleaned surface of a metal for its protection from
corrosion. It is helpful to protect towers and bridges.
(xiii) Pack cementation: Heating a base metal with another powdered metal is done. The powdered
metal forms a coat on the base metal, and hence protects it from corrosion. Sherardising and
colourising are two important pack cementation processes in which Zn and Al powders are heated,
respectively. This type of coating is also known as diffusion coating. The third one is chromising in
which the chromium powder and alumina are heated with Fe.
(xiv) Electroplating: It is an electrochemical process in which a base metal is coated by Zn, Ag, Cr,
Au, Sn, etc. to protect it from corrosion and also to make it shining and decorative. The base metal is
made cathode, dipped in a suitable electrolyte, and the metal to be deposited is made the anode.
Anodised coating is done for non-ferrous metals, such as Zn, Mg and Al. (xv) In organic coatings,
paints, varnishes, enamels, lacquers and emulsion paints are included. They not only protect from
corrosion but also give a good look to the metal. Organic coats must have chemical inertness, good
surface adhesiveness and non effectiveness towards inorganic chemicals and water.
(xv) Water-repellent paints are organosilicon compounds, which on hydrolysis give a permanent coat
that can work for 4–6 years as corrosion resistant.
(xvi) Cathodic protection or electrochemical method: The principle involved in this method is that
the metal is forced to behave like cathode, thereby preventing corrosion. For protecting corrosion–
electrochemically, there are two methods for the protection from corrosion.
Sacrificial anodic protection: Underground steel pipes are protected from corrosion by this method.
A magnesium rod is fixed near the metal under protection and both are connected with a conducting
wire. Magnesium is more positive than iron and, hence, in electrochemical cell it acts as anode and
the iron acts as cathode. According to the principle of galvanic cell, it is anode that undergoes
oxidation, and, hence, corrosion occurs at anode saving cathode (iron) from corrosion. Thus,
magnesium sacrifices itself for saving the iron.

Sacrificial anodic protection

Impressed current cathodic protection: The object to be protected is made cathode and it is connected
to the negative terminal of a DC (direct current) source. The positive terminal of the source is
connected to the other electrode made of graphite or platinum, lead or nickel. The impressed current
opposes the galvanic current (corrosion current) and, hence, protection from corrosion takes place.

Impressed current cathodic protection

Passivity : The process in which a metal exhibits higher corrosion resistance is called passivity. When
a very thin, invisible and highly protective film is formed on the surface of a metal or an alloy, it is
called passivity. This film is insoluble and non-passive. A metal is passive in a certain environment if
its corrosion rate is very low. By the change of the environment, the passivity of a metal may change
and may become active towards corrosion.
Low carbon steel does not corrode in conc. HNO 3 due to protection effect of passive film. However,
in dil. HNO3 does not form a stable passive film and therefore dissolves steel. Passive film is formed
on the surface of aluminium, Cr, Si, Ti in air, water and dilute acids. There is a good corrosion and
oxidation resistance of stainless oxide passive layer. A damage of passive film may cause intensive
localised corrosion (pitting corrosion). Passive oxide layers are dissolved in electrolytes containing
SO4- and chlorides. Phosphate and chromate ions stabilise passive films, promoting the regain of its
defects.
Quick Review
1. What is corrosion? Why do metals corrode?
2. Why most of the metals are found in the ore form and not in the pure form? Explain.
3. Describe the electrochemical theory of corrosion.
4. Show the reactions involved in the hydrogen evolution and oxygen absorption types of corrosion.
5. Describe the factors on which corrosion depends.
6. Discuss various methods of protection corrosion.
10. Define passivity with an example.
12. Why rusting of iron is fast in saline water than in ordinary water?
13. What is differential aeration corrosion?
15. What is cementation?
18. Explain sacrificial anodic and impressed current cathodic protection?
20. Differentiate anodic and cathodic protection methods?
21. Differentiate galvanizing and tinning?