Corrosion

Definition of corrosion: The gradual destruction of metals or alloys due to

chemical/electrochemical attack of the components of the environment (i.e. H 2O & O2) is called
corrosion.
Example: When iron is exposed to air in the presence of moisture, hydrated ferric oxide (rust) is
formed.

Electrochemical theory of corrosion: According to electrochemical theory, when a metal such

as iron is exposed to corrosive environment, due to heterogeneity at the metallic surface a large
number of tiny galvanic cells with anodic and cathodic regions are formed.

Reactions:

O2
Fe2+Fe2 OH- H2O
+ OH-ff
Cathodic region
Anodic region
Electrons

Iron Metal

R
eaction at anodic region: Oxidation of metal takes place at anode. The Fe is oxidized to Fe 2+
ions, so corrosion takes place at the anodic region

Reactions at cathodic region: The metal is unaffected at the cathodic region.

At cathode, the reaction is either a) liberation of hydrogen or b) absorption of oxygen.

a) Liberation of hydrogen (in the absence of oxygen)

In acidic medium the reaction is

In neutral or alkaline medium, the reaction is

b) Absorption of oxygen (in the presence of oxygen)

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In neutral or alkaline medium, the reaction is

Overall reaction:

Ferrous ion Hydroxide ion ferrous hydroxide

In the presence of moisture and atmospheric oxygen ferrous hydroxide is oxidized to ferric
hydroxide or hydrated ferric oxide.

Electrochemical Series:
The arrangement of elements in their increasing order of their standard reduction electrode
potential.
The electro chemical series does not take passivity of the metals into account. In order to
overcome the limitations of electrochemical series, galvanic series was introduced.

Galvanic Series: It is a series which lists various metals and alloys in increasing of their
resistance to corrosion determined in sea water at 250C.

Different types of corrosion:

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1. Differential metal corrosion or Metallic Corrosion or Galvanic Corrosion:

Zn metal Fe metal Fe metal Cu metal

Anode Cathode Anode Cathode

This type of corrosion occurs when two dissimilar metals are in contact with each other and are
exposed to a corrosive environment. The two metals differ in their electrode potentials. The
metal with lower electrode potential acts as anode and the other metal with higher electrode
potential acts as cathode.
The anodic metal undergoes oxidation and gets corroded.
A reduction reaction occurs at the cathodic metal.
The cathodic metal does-not undergo corrosion.
The reactions may be represented as follows:

Higher the potential difference between the anodic and cathodic metals, higher is the rate of
corrosion.
Example: When iron is in contact with zinc, iron acts as cathode and zinc acts as anode because
zinc has lower electrode potential compared to iron. Therefore, zinc undergoes corrosion. On
the other hand, when iron is in contact with tin, iron acts as anode because iron has lower
electrode potential compared to tin. Hence iron undergoes corrosion.

Other examples:
1. Steel screws in copper sheet.
2. Steel screws with copper washer.
3. Bolt and nut made of different metals.

2. Differential aeration corrosion:

Iron

More O2,
(Cathode)

Less O2,
(Anode)
Water

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This type of corrosion occurs when two different parts of the same metal are exposed to different
oxygen concentrations. (E.g. An iron rod partially dipped in water.) The part of the metal which
is exposed to less oxygen concentration acts as anode. The part which is exposed to more
oxygen concentration acts as cathode. The anodic region undergoes corrosion and the cathodic
region is unaffected.

The reactions may be represented as follows:

Other examples:
1. Part of the nail inside the wall undergoes corrosion.
2. When a dirt particle sits on a metal bar, the part under the dirt undergoes corrosion.
3. Partially filled iron tank undergoes corrosion inside water.

3. Pitting corrosion: This is an example of differential aeration corrosion.

When a small dust particle gets deposited on a steel surface, the region below the dust particle
is exposed to less oxygen compared to the remaining part. As a result, the region below the dust
particle acts as anode undergoes corrosion and forms a pit. The remaining region of the metal
acts as cathode and is unaffected.

The reactions may be represented as follows:

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Formation of a small anodic area and a large cathodic area results in intense localized corrosion
below the dust particle.

4. Water line corrosion: This is an example of differential aeration corrosion.

More oxygen,
(Cathode) Rust

Less
Oxygen Water
(Anode)

When a steel tank is partially filled with water for a long time, the inner portion of the tank
below the water line is exposed only to dissolve oxygen, where as, the portion above the water
line is exposed to more oxygen. Thus the portion below the water line acts as anode and
undergoes corrosion. The upper portion acts as cathode and is unaffected.
A distinct brown line is formed just below the water line due to the deposition of rust.

The reactions may be represented as follows:

Other example: Ships which remain partially immersed in sea water for a long time undergo
water line corrosion.

5. Stress corrosion:

Stress corrosion occurs when stressed region of metals are exposed to corrosive environments.
The stressed region acts as anode and undergoes corrosion (refer diagram below). The
unstressed region acts as cathode and is unaffected.
Stress corrosion is highly specific in nature. For corrosion to occur, the presence of both stress
and a specific corrosive environment are necessary.
For example, stressed regions of mild steel undergo stress corrosion in the presence of NaOH
solution. (caustic embrittlement of boilers)

Anode undergoes
corrosion
Factors affecting the rate of corrosion:

1. Nature of the metal: Metals with lower electrode potentials are more reactive and are more
susceptible to corrosion. For example, elements such as Mg and Zn, which have low electrode
potentials, are highly susceptible to corrosion. Noble metal such as, which have higher electrode
potentials, are less susceptible to corrosion.

Exceptions: Metals and alloys which show passivity are exceptions for this general trend. Such
metals form a protective coating on the surface which prevents corrosion.

2. Nature of corrosion product: If the corrosion product is insoluble, stable and non-porous,
then it acts as a protective film which prevents further corrosion. The film acts as a barrier
between the fresh metal surface and the corrosive environment. On the other hand, if the
corrosion product is soluble, unstable and porous, then the corrosion process continues even after
the formation of corrosion product.

Example: Aluminium, titanium and chromium form a protective film of metal oxide on the
surface. But in the case of Zn and Fe, the corrosion products formed don’t have protective value.

3. Difference in potential between anodic and cathodic regions: Larger the potential
difference between the anodic and cathodic regions, higher is the rate of corrosion. For example,
the potential difference between iron and copper is 0.78 V, and between iron and tin is 0.3 V.
Therefore, corrosion is faster when iron is in contact with copper.
The use of dissimilar metals should be avoided wherever possible. Otherwise, the anodic metal
gets corroded.

4. Anodic and cathodic areas: Smaller the anodic area and larger the cathodic area, more
intense and faster is the corrosion. For example, a broken coating of tin on iron surface results in
intense corrosion at the broken region. Iron is anodic to tin. Exposed region of iron acts as
anode with small area. Tin acts as cathode which has large area.

Rate of Corrosion of Anodic area = Cathodic Area

Anodic Area
5. Anodic and cathodic polarizations: The polarization at anode and cathode decreases the rate
of corrosion.
Anodic polarization occurs due to the accumulation of metal ions in the vicinity of anodic
region. This retards the formation of new metal ions by oxidation. Thus the corrosion process is
retarded.
Cathodic polarization occurs due to the accumulation of hydroxyl ions in the vicinity of
cathodic region. This accumulation retards the movement of cathodic reactant (oxygen) towards
the cathodic surface. Hence, the rate of cathodic reaction decreases. A retarded cathodic reaction
results in retardation of anodic reaction also. Thus the corrosion process is slowed down.

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6. Hydrogen over voltage: When the cathodic reaction is liberation of hydrogen, a high
hydrogen over voltage retards the cathodic reaction. A retarded cathodic reaction retards the
anodic reaction also. Thus the corrosion process is slowed down.

(Note: Theoretically, certain potential difference between anode and cathode is required for the
hydrogen evolution to occur at cathode. But in practice, the potential difference required is more
than the theoretical value. This excess potential difference is called hydrogen over voltage.)

7. pH of the medium: In general, lower the pH of corrosion medium, higher is the corrosion
rate. (Exception: Metals like Al, Zn etc. undergo fast corrosion in media with high pH.)

Iron does not undergo corrosion at pH greater than 10. This is due to the formation of protective
coating of hydrous oxides of iron. Between p H 10 and 3, the presence of oxygen is essential for
corrosion. If the pH is less than 3, corrosion occurs even in the absence of oxygen.

8. Temperature: Higher the temperature, higher is the rate of corrosion.

In general, the rate of a chemical reaction increases with increase in temperature. Corrosion is
one such chemical reaction. Therefore, the rate of corrosion increases as the temperature
increases.
Increase in temperature increases the ionic conductivity of the corrosive medium. This also
contributes to the increase in corrosion rate.
9. Humidity: The rate of corrosion generally increases with increase in Humidity of the
medium.
10. Conductivity of the medium: Higher the conductivity of the medium higher is the rate of
corrosion. p. Eg: 1) Corrosion occurs faster in sea water than river water 2) Corrosion occurs
faster in clayed soil than dryed soil.

Corrosion control

Metal Coating: The process of covering the base metal with a layer of another metal is known
as metal coating.

Galvanizing: The process in which a thin layer of zinc is coated on the surface of iron by hot
dipping process is known as Galvanization.

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Galvanization is carried out by hot dipping method. It involves the following steps.

1. The metal surface is washed with organic solvents to remove organic matter on the surface.
2. Rust is removed by washing with dilute sulphuric acid.
3. Finally, the article is washed with water and air-dried.
4. The article is then passed through zinc chloride and ammonium chloride flux.
5. Then the article is dipped in a bath of molten zinc at 430- 4700C.
6. The excess zinc is removed by rolling, wiping or by means of air blast.

Applications: Galvanization of iron is carried out to produce roofing sheets, fencing wire,
buckets, bolts, nuts, pipes etc.

(Note: Even if the Zn coating falls off at some places, the base metal (Fe) does not get corroded
at those places. This is because the base metal acts as cathode. In corrosion process, the
cathodic metal always remains unaffected.)

(Note: Galvanized articles are not used for preparing and storing food because zinc dissolves in
dilute acids producing toxic zinc compounds)

Tinning: Tinning is the process of coating the surface of a base metal (such as iron) with tin by
hot dipping process is known as Tinning. Tinning of iron metal is an example of cathodic metal
coating on an anodic base metal.

Tinning of iron is carried out by hot dipping method. It involves the following steps.

1. The metal surface is washed with organic solvents to remove organic matter on the
surface.
2. Rust is removed by washing with dilute sulphuric acid.
3. Finally, the article is washed with water and air-dried.
4. The article is then passed through zinc chloride and ammonium chloride flux.
5. Then the article is dipped in a bath of molten tin.
6. Finally it is dipped in palm oil to avoid oxidation of tin. The oil prevents the oxidation of
tin coating.

Applications: Tin-coated steel is used for manufacturing containers.

(Note: In tin-coated iron, if a small crack is formed in the tin coating, an intense and accelerated
corrosion of iron occurs at the exposed region because of the formation of a small anodic area
(iron) and large cathodic area (tin)).

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(Note: Copper utensils are coated with tin to prevent contamination of food with poisonous
copper salts.)
Cathodic protection:- sacrificial anode method:

Electrical conductor

Steel pipe Cathode

Mg/Zn
Block
(Anode)

In cathodic protection, the metal to be protected is completely converted into a cathode. Since
cathodes do not undergo corrosion, the metal is protected against corrosion.

In sacrificial anode method, the metal to be protected is electrically connected to a more active
metal. For example, when steel is to be protected, it may be connected to a block of Mg or Zn.
In such a situation, steel acts as cathode and is unaffected. Mg and Zn act as anode and
undergo sacrificial corrosion. When the sacrificial anode gets exhausted, it is replaced with new
ones.
Other examples: Mg bars are fixed to the sides of ships to act as sacrificial anode.
Mg blocks are connected to buried pipe lines.
Cathodic protection:- impressed current method (impressed voltage method)

In cathodic protection, the metal to be protected is completely converted into a cathode. Since cathodes
do not undergo corrosion, the metal is protected against corrosion.
In impressed current method, the metal to be protected is connected to the negative terminal of an
external d.c. power supply. The positive terminal is connected to an inert electrode such as graphite.

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Under these conditions, the metal acts as cathode and hence does not undergo corrosion. The inert
electrode acts as anode; but it does not undergo corrosion because it is inert.

Battery

Protected metal
Inert anode

Questions
1. Define corrosion.
2. Explain electrochemical theory of corrosion with iron as example.
3. What are the reactions taking place at anodic area and cathodic area in the corrosion of
iron? What is overall reaction?

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4. Describe I) differential metal corrosion or galvanic corrosion. II) differential aeration
corrosion, III) pitting corrosion,IV) waterline corrosion.
5. Explain Stress corrosion.
6. Explain the factors affecting the rate of corrosion.
7. What is anodizing? Describe anodizing of alumiinium.
8. What is phosphating? Explain.
9. What is galvanizing? Describe galvanizing of iron.
10. What is Tinning? Describe tinning.
11. Pin holes in tin-coated iron are more prone to corrosion than pin holes in zinc coated iron.
Why?
12. Explain cathodic protection by sacrificial anode method OR Explain corrosion control by
sacrificial anode method.
13. Describe cathodic protection by impressed current method OR Explain corrosion control by
impressed voltage method.

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