FOR B TECH 2023-24

APPLIED CHEMISTRY

UNIT II: CONCEPTS OF CORROSION SCIENCE

INTRODUCTION

The term corrosion is used to denote a change. A metal changes from its elementary state to the
combined state, more or less rapidly, when it comes into contact with the gaseous/liquid
medium. This is actually owing to the chemical interaction between the metal and the
environment.

DEFINITION

Corrosion “The spontaneous destruction and consequent loss of a metal/alloy due to

unavoidable chemical/electrochemical attack by the environment”
Many metals exist in nature in combined form as their oxides, carbonates, sulphides, chlorides
and silicates (except noble metals) such as Au (gold), Pt (Platinum) etc. During extraction
process these are reduced to their metallic state from their ores and during extraction of ores
considerable amount of energy is required. Compounds are in lower energy state than the metals.
Hence when metals are put into use in various forms, they get exposed to environment such as
dry gases, moisture, liquids etc. and slowly the exposed metal surface begin to decay by
conversion into a compound.
Example:
1. When Cu is exposed to the industrial environment it forms an adherent protective green deposit
which isolates the metal from the environment, hence the further action is very slow.
2. When iron metal is exposed to the industrial environment, the metal forms a loosely adherent
product of hydrated ferric oxide called rust, which is relatively non-protective.

Hence, the fundamental approach to the phenomena of corrosion, the structural features of the
metal, reactions which occur at the interface and nature of the environment are to be considered.
 Disadvantages of corrosion:

The process of corrosion is slow and occurs only at surface of metals but the losses incurred are
enormous. Destruction of machines, equipment, building materials and different types of metallic
products, structures etc. Thus the losses incurred are very huge and it is estimated that the losses
due to corrosion are approximately 2 to 2.5 billion dollars per annum all over the world.

THEORIES OF CORROSION:

Corrosion can be explained by the following two theories.

1. Dry or chemical corrosion.

2. Wet or electrochemical corrosion.

DRY OR CHEMICAL CORROSION:

This type of corrosion occurs mainly by the direct chemical action of the environment i.e., by the
direct attack of atmospheric gases such as O2, halogens, H2S, SO2, N2 or anhydrous inorganic
liquids on the metal surface with which they are in contact. There are 3 main types of chemical
corrosion. 1) Corrosion by oxygen (or) oxidation corrosion. 2) Corrosion by other gases like
SO2, CO2, H2S and F2 etc and 3) Liquid metal corrosion.

WET (OR) ELECTROCHEMICAL CORROSION: Most of the corrosion cases are

electrochemical in nature, taking place by an electrochemical attack on the metal in the presence
of moisture/conducting medium- called wet corrosion.

This type of corrosion is observed when, a conducting liquid is in contact with a metal (or)

 When two dissimilar metals (or) alloys are either immersed (or) dipped partially in a solution.

 The corrosion occurs due to the existence of separate anodic and cathodic areas or parts between
which current flows through the conduction solution.

 In the anodic area oxidation reaction takes place so anodic metal is destroyed by dissolving (or)
forming a compound such as an oxide. Hence corrosion always occurs at anodic area.

 At Anode: M M n+ + ne –

 At cathode: M n+ + ne – M


ELECTROCHEMICAL THEORY OF CORROSION

According to the theory, when a metal is in contact with the conducting medium or when
dissimilar metals/alloys are either immersed partially/completely in a solution, a large number of
galvanic cells with the existence of anodic and cathodic area on the metal, are formed.
In this corrosion, oxidation of the metal and reduction of species present in solution takes place
at anodic and cathodic parts, respectively.
 The electrons are transferred through the metal from anode to cathode.
 The anodic part of the metal suffers from corrosion and cathode is protected from corrosion.
Corrosion reactions are:
a) At anode (oxidation reaction): Fe Fe 2+ + 2e-
The reaction at cathode (reduction reaction) depends on the nature of the environment
If the medium is acidic,
b) In the presence of dissolved oxygen: 2H+ + ½O2 +, 2e- H 2O
c) In the absence of dissolved oxygen: 2H+ + 2e- H 2
d) If the medium is alkaline/neutral,
In the presence of dissolved oxygen: H2O+½ O2 + 2e- 2OH-
In the absence of dissolved oxygen: 2H2O+2e- 2OH- + H2
Example: Rusting of an Iron in the presence of moist air
At anode: Fe Fe+2 + 2e-
At cathode: H2O+ ½O2 + 2e- 2OH-
Net reaction: Fe+2 + 2OH- Fe (OH) 2
Net cell reaction: Fe2+ + 2OH- Fe (OH) 2
In the presence of excess of oxygen: 2Fe (OH) 2+ ½ O2 Fe2O3.2H2O
(Red rust) Hematite

In the limited supply of oxygen: 3Fe (OH) 2+½ O2 Fe3O4.3H2O

(Black rust) Magnetite

GALVANIC SERIES

Galvanic series is an arrangement of metals and semi-metals according to their nobility. Nobility is
the resistance to corrosion and oxidation in the presence of moist air. This list includes metals,
semi-metals and alloys because only these materials can undergo corrosion. The list is arranged
according to the relative potentials of these materials in a specified environment. Generally, the
environment that we consider is sea water.
The galvanic series is a chart showing the relationships and a guide for selecting metals that
can be joined, with the aim of helping in the decision-making process. The use of the galvanic
series has to be done with caution and basic knowledge of the environments that are a
necessary part of this serious form.

ELECTROCHEMICAL SERIES
Electrochemical series is a list of chemical elements which shows the order of standard electrode
potentials of them. It provides enough information about the relative reactivity of metals in
aqueous solutions under standard conditions. Another common name for this series is “activity
series”. Furthermore, this series lists metals in order of decreasing the reactivity.
At the top of the series, it has alkali metals and alkaline earth metals. These are more reactive
and easily undergo oxidation than the metals at the bottom. Moreover, they easily react to form
compounds. These metals are thus called “active metals”.
 DIFFERENCES BETWEEN ELECTROCHEMICAL SERIES AND GALVANIC
SERIES:
Electrochemical series:
 Electrode potentials are measured by dipping pure metals in their salt solution of 1 M
concentration, without any oxide films on them.
 The position of a metal in the electrochemical series is fixed.
 It gives no information regarding the positions of alloys.
 The position of metal is permanently fixed in this series.
 This series comprises metals and non-metals.
 It predicts the relative displacement tendencies.
 It is absolute.
 It is quantitative.
 It is a series only for pure metals.
 It is used for theoretical calculations.
Galvanic series:
 This series was developed by studying corrosion of metals and alloys in unpolluted
seawater, without their oxide films, if any removed.
 In galvanic series, the position of a given metal may shift.
 Their corrosion can be studied from this series since alloys are included in galvanic series
 The position of metal, when present in the form of an alloy is different from that of pure
metal.
 This series comprises metals and alloys.
 It predicts the relative corrosion tendencies.
 It is relative.
 It is qualitative.
 It is a series for pure metals and alloys also.
 It is used for practical applications.
TYPES OF CORROSION
1) DIFFERENTIAL METAL CORROSION or GALVANIC CORROSION
When two dissimilar metals are in direct contact with one another and exposed to a corrosive
conducting medium, the metal higher up in the electrochemical series behaves as anode and
suffers from corrosion, whereas the metal lower in the electrochemical series becomes cathode
and protected from corrosion. This type of corrosion is also known as Galvanic corrosion. If the
potential difference between the electrodes is high, greater the extent of corrosion.
When Zn and Cu metals are electrically connected and exposed to an electrolyte, Zn (higher in
electrochemical series) forms anode and suffers from corrosion whereas Cu (lower in
electrochemical series) forms cathode and protected from corrosion.
 Examples: Steel screws in a brass marine hardware.
Steel pipe connected to copper plumbing.

Reactions:

At Anode: Fe Fe2+ + 2e-

At Anode: ½ O2 + H2O+ 2e- 2OH-

Net cell reaction: Fe2+ + 2OH- Fe (OH) 2

2) DIFFERENTIAL AERATION CORROSION

This type of corrosion is due to the formation of differential aeration cell or oxygen
concentration cell. When a metal surface is exposed to differential air or oxygen concentrations-
forms differential aeration cell. The more oxygenated part of the metal behaves as cathode and
less oxygenated part becomes cathode. Differential aeration of metal causes a flow of current
called the differential current and the corrosion is called differential aeration corrosion.
Example: Consider a piece of Fe metal is partially immersed in water and agitated properly. The
part of the metal above and closely adjacent to the water-line are more oxygenated, because of
easy access of oxygen and hence become cathodic. The part of the Fe metal immersed to greater
depth, which have less access of oxygen and becomes anode. Hence a difference in potential
between the electrodes is created, which causes a flow of current between the two differentially
aerated areas of the same metal and causes corrosion at anode.

Reactions:

At Anode: Fe Fe2+ + 2e-

At Anode: ½ O2+ H2O+ 2e- 2OH-

Net cell reaction: Fe2+ + 2OH- Fe (OH) 2

a) PITTING CORROSION
In pitting corrosion, a pit is formed when the protective coating on the metal surface break, a
micro pit (anode) formed on the metal surface. Once the pit is form the process of corrosion
become very fast due to different amount of oxygen in contact with metal surface. The portion
(pit) with lower concentration of oxygen become anode and portion with higher concentration of
oxygen becomes cathode. Pitting is very destructive and frequently ruins the tubes, pipes etc.
This attack becomes more intensified with time.

b) CREVICE CORROSION

Crevice corrosion occurs between two joining surface which may be metal-metal or metal-
nonmetal crevice. It is localized corrosion of metal or alloy surface. In crevice corrosion, the
joining surface area has low concentration of oxygen as compare to other area. The portion with
lower concentration of oxygen become anode and portion with higher concentration of oxygen
becomes cathode. The area of the crevice become anodic and undergoes corrosion. Corrosion is
further accelerated by the deposition of dirt's, scales, debris etc.

Crevice corrosion more commonly observed at, sharp re-entrant corners, incomplete weld
penetration, overlapping surfaces, under gaskets etc.
3) INTER GRANULAR CORROSION OR GRAIN BOUNDARY CORROSION

This type of corrosion occurs along grain boundaries and only where the material especially sensitive
to corrosive attack exists and corrosive media attacks only at the grain boundaries. This may be due
to the precipitation of certain compounds at the grain boundaries.

Inter granular corrosion occurs when certain alloys reach temperatures between 425°C and 870°C.
This type of corrosion is generally takes place in welding, annealing process.

Chromium present in the stainless steel alloy reacts with carbon to create chromium carbide near the
grain boundaries. .
For example when stainless steel containing (18% Cr, 8% Ni and 0.1% C) carbon more than
0.1% heated to high temperature and cooled slowly or held at 650 ℃ for short interval of time, a
rapid reaction takes place between carbon and chromium forming Chromium carbide (CrC). The
CrC precipitates at the grain boundaries and results in the formation of galvanic cell. Thereby the
region just adjacent to grain boundaries become depleted in chromium composition and is more
anodic with respect to the grain center (cathodic). Corrosion takes place at the grain boundary.
This is known as grain boundary corrosion.

It can be reduced when different grades of carbon are used, preferably low carbon grades. It can
also be stabilized by the addition of niobium or titanium, which has greater affinity to the
formation of carbides than chromium. This is useful as the precipitation of carbides from the
material will not affect the chromium content. Grain boundary corrosion is not observed in pure
metals.
FACTORS AFFECTING THE RATE OF CORROSION

1. The physical state of metal or nature of metal or electrode potential: The tendency of a
metal to undergo corrosion is dependent on the nature of the metal. Metals with lower reduction
potential undergo corrosion easily whereas metals with higher reduction potential do not undergo
corrosion easily.
The extent of galvanic corrosion depends on the position of the metals in galvanic series. The
metal which is placed at higher position in the series are more reactive and undergoes corrosion.
The rate and severity of corrosion, depend upon the difference in their position and greater is the
difference, the faster is the corrosion of the anodic metal.

The reactive metals like Na, K, Mg, and Zn are more susceptible for corrosion. The noble metals
like Ag, Au, Pt, Pd are less susceptible for corrosion.

2. Nature of the oxide film or nature of corrosion product: If the corrosion product is formed
on the metal surface is soluble, porous, less ionic, non uniform, and non stoichiometric in ratio,
then corrosion proceeds at a faster rate. If the corrosion product is insoluble, nonporous, uniform,
and stoichiometric in ratio, then the corrosion will be suppressed. Passive metals (Al, Cr, Ni etc)
get protected by corrosion product but not the active metals (Fe, Zn, Mg etc).

3. Area of anode and cathode: The rate of the corrosion is greatly influenced by the relative
sizes of cathodic and anodic areas.
• If the metal has smaller the anodic area and larger the cathodic area exposed to corrosive
atmosphere, more intense and faster is the corrosion occurring at anodic area because at anode
oxidation takes place and electrons are liberated. At the cathode these electrons are consumed.
When anode is smaller and cathode region is larger all the liberated electrons at anode are rapidly
consumed. This process makes the anodic reaction to takes place at its maximum rate thus
increasing the corrosion rate. If the cathode is smaller and reverse process takes place decrease
rate of corrosion.
For e.g. If tin (Sn) coated on iron (Fe) and in that some area are not covered or some pin holes
are left, there forms smaller anodic area and larger cathodic area because tin is cathodic with
respect to iron so intense localized corrosion takes place. On the other hand if Zn coated to Fe
then if there are some pin holes are there creates larger anodic area and smaller cathodic area
because Fe is cathodic with respect to zinc so that rate of corrosion is very less.

4. pH of the medium: In general, lower the pH higher is the rate of corrosion.

• If the pH is greater than 10, corrosion of iron is very less due to the formation of protective
coating of hydrous oxides of iron.
• If pH is between 10 and 3, then presence of oxygen is essential for corrosion of iron.
• If the pH is 3 or lower than 3 severe corrosion occurs in the absence of air due to the
continuous evolution of H2 at cathode. However metals like Al, Zn etc. undergo fast corrosion in
highly alkaline medium.

5. Temperature and conductance of the electrolytic medium: - Raise in temperature increases

the rate of corrosion due to increase in the rate of diffusion of ions. Higher the electrical
conductivity of the liquid in contact with the metal surface, greater is the rate of corrosion. Saline
environment is more corrosive than normal environment due to the presence of higher
concentration of dissolved salts in the liquid due to which metals nearby sea area is more
corroded.

6. Humidity in 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, CO 2, SO2 etc in the air to
produce the electrolyte which is required for setting up of an electrochemical cell.

7. Polarization of electrode: The shift in electrode potential which results from the effects of
current flow w.r.t. the zero current flow potential. All corrosion reactions involve current flow
and will alter the potential of the metal surfaces involved.

CORROSION MEASUREMENT BY TAFEL EXTRAPOLATION METHOD

Tafel extrapolation is a mathematical technique used to estimate the corrosion current (Icorr) or
the corrosion potential (Ecorr) in an electrochemical cell and thereby, the corrosion rate. So, the
corrosion rate of a metal can be measured using an electrochemical cell.

The electrochemical cell consists of working electrode (metal), auxiliary electrode and reference
electrode connected to a potentiostat.
A Tafel plot represents the relationship between the measured potential and the logarithmic
current density. Extrapolation involves extending a known sequence of linear values within a
Tafel plot to determine the parameters. Tafel extrapolations can be performed either manually or
by specialized computer software.

Prior to the application of external current, the potential of the sample remains same and that will
be the corrosion potential (Ecorr) of the metal in the respective environment. At this potential,
the anodic and cathodic reactions are usually minimum. The potential of the working electrode is
varied at a predetermined rate in the negative direction and the corresponding change in current
is measured. Similar measurements are also performed in the anodic direction. The measured
potential vs log current values are plotted in the semi-logarithmic plot as shown in Fig.
The curve consists of two diverging logarithmic plot lines representing the anodic and cathodic
currents. Extrapolation is performed by extending the linear portions of the anodic and cathodic
plots back to their intersection. These two lines eventually meet at a point where the corrosion
current, Icorr is obtained. The value of Icorr can be used in the mathematical equation to
calculate the corrosion rate.

Where K is a constant,
Icorr is the corrosion current density,
ρ is the density of the metal and
Eq.wt: is the equivalent weight of the respective metal.

POURBIAX DIAGRAM

A Pourbaix diagram provides information about the stability of a metal as a function of pH

and potential. These diagrams are available for over 70 different metals. Pourbaix diagrams
have several uses, including in corrosion studies.
A Pourbaix diagram is also known as a potential/pH diagram, equilibrium diagram describing
various chemical phase formation at different pH and potential. Pourbaix diagrams are plotted
by using the Nernst equation (an equation used to calculate electrode potential). As the Nernst
equation is derived entirely from thermodynamics, the Pourbaix diagram can be used to
determine which species (metals or alloys) is thermodynamically stable at a given electrode
potential (E) and pH.

Characteristics of Pourbaix diagram:

A Pourbaix diagram is divided in regions of “immunity”, “corrosion” and “passivity”. These
regions provide information about the stability of a particular metal or alloy in a specific
aqueous electrochemical environment under certain pH, E vs SHE, pressure and temperature
conditions.
 The immunity region is the region in which there is no metal dissolution.
 The corrosion region is the region in which there is active metal dissolution.
 The passivation region is the region in which a protective metal-oxide film that prevents metal
dissolution is formed.
 Horizontal lines represent electron transfer reactions. They are pH-independent, but potential-
dependent.
 Vertical lines are potential-independent but pH-dependent and not accompanied by any
electron transfer.
 Sloping, straight lines give the redox potentials of a solution in equilibrium. This equilibrium
indicates electron transfer as well as pH.
 The diagram is computed for the equilibrium conditions at 25°C.
The main objectives of the Pourbaix diagrams are:
1. To show the direction of the various reactions at a given pH and potential.
2. To make a basis for estimation of the corrosion product compositions at various pH and
potential combinations.

3. To show which environmental pH and potential changes will reduce or preventcorrosion.

Pourbaix diagram of water

As water is the medium in the electrochemical or wet corrosion we must understand the
stability region of H2O in the Pourbaix diagram. From the diagram below we see the stability
region of water. When the E is increased or decreased the water undergoes oxidation or
reduction. This appears along with all metal system. Water is thermodynamically stable at the
range of zero to 1.23 V of potential.
POURBAIX DIAGRAM OF ALUMINIUM

In the Pourbaix diagram of Aluminum, the water is shown as dotted lines. The region below and
above the line water is not stable. In the above Pourbaix diagram, in the Passivation region Al is
covered by the Al2O3 layer. The Al2O3 layers protect the base Al from further corrosion in the
particulate pH and E conditions. However, when the pH increases or decreases beyond the
passivation zone the passive layers undergo dissolution and Al 3+ appears in the solution as
corroded products. From the Pourbaix diagram it is obvious that Al (immune zone) is stable
only in more cathodic region, even at neutral conditions.

The reactions corresponding to lines in the diagram are as follows.

1) Al → Al3+ + 3e (Redox reaction it is Potential dependent and pH independent)

2) 2Al + 3 H2O → Al2O3 + 6H+ (Acid base reaction it is pH dependent and Potential
independent)

3) ½ Al2O3 + 3H++ 3e- →Al + 3/2 H2O (both redox and acid base reaction, it depends both pH
and Potential)

In both acidic and alkali medium Aluminum undergo corrosion at higher potentials (pH between 1-4
and 8-14). In neutral condition Aluminum forms thin layer on the surface called passive layer.
POURBAIX DIAGRAM OF IRON

Iron is the most common and hard engineering material. But this undergoes severe corrosion
and forms various corrosion products at different potential and pH conditions. The below
Pourbaix illustrates the immune zone in more cathodic and acidic pH. It is obvious that pure
Iron is relatively stable than Al in E-pH diagram. In the oxidative or positive potential and
acidic (pH <6) condition iron undergoes dissolution and corroded products ( Fe +2 and Fe3+)
remains inthe solution.

In neutral and slight basic conditions formation Fe(OH)2 is observed in Poubaix diagram. As
the potential increases Fe(OH)3 and oxide formation is observed in the E-pH diagram.

Rust formation is observed more oxidative condition with higher pH conditions. In highly
basic and cathodic potential conditions (in boiler corrosion – caustic embrittlement) formation
of FeO2- is observed.

From the Pourbaix diagram it is obvious that Iron undergoes various phase formation with
change in pH and potentials.

The reactions corresponding to lines in the diagram are as follows.

1) Fe+2 + 2e- → Fe (pure redox reaction, it is pH independent and Potential dependent)

2) 2Fe3+ + 3H2O → Fe2O3 + 6H+ (Pure acid base reaction, it is pH dependent and
Potentialindependent)

3) 2Fe+2+ 3H2O → Fe2O3 + 6H+ + 2e - (Both redox and acid base reaction, it depends on both
pH and Potential)

The knowledge and understanding the stability of these engineering metals is very important
in device, construction and engineering material applications.

WHAT IS PASSIVITY?

Formation of an oxide/compound on a metal or alloy surface that is stable in the electrolyte, so

that the metal is rendered “passive” in the environment (i.e., the material is “passivated”)
Passivation limits corrosion.

The film formed during passivation is insoluble, non-porous and of such a “self healing nature”,
that when broken, it will repair itself on re-exposure to oxidizing conditions.

The common examples of passive metals and alloys are Ti, Al, Cr and a wide variety of stainless
steel alloys containing Cr. These exhibit outstanding corrosion resistance in oxidizing
environments, but in reducing environments, they become chemically active.

Fe-Cr alloys exhibit an increasing tendency to passivate as the Cr content increases!

Cl- ions can break down an existing passive film.

It is generally difficult/impossible to passivate an alloy or metal in the presence of aggressive

ions such as Cl-.

Aggressive ions affect the critical current density for passivity and may become incorporated
into the oxide (and cause defects) Local breakdown of a passive film can lead to pitting
corrosion.

Many metals like Cr, Ti, Al, Ni exhibit a reduction in their corrosion rate above certain critical
potential because of the formation of a protective, thin oxide film.

Passivation is the reason for the excellent corrosion resistance of Al and S.S.

IMMUNITY VERSUS PASSIVITY

Immunity means that the metal or alloy will not corrode – Period! It is thermodynamically
impossible.

Passivity means that the alloy will not corrode under specific conditions, but it is not permanent.

It is still thermodynamically possible for the alloy to corrode.

 STANDARD ELECTRODE POTENTIAL & POLARIZATION

Standard electrode potential is the potential developed when all the reactants have unit activity
and the temperature is 25 ℃. Oxidation and reduction reactions are of equal rate.

When these reactions are not equal, there will be a net oxidation or reduction and the potential
shifts away from its equilibrium value. This process is polarization.

CORROSION CONTROL

Corrosion can be completely avoided only under ideal conditions. Since it is impossible to attain
such conditions, it can be minimized by using various corrosion control methods.

They are:
1.Protective coatings
2.Cathodic protection
3.Anodic protection
4.Corrosion Inhibitors
5.Spray Technique

1. PROTECTIVE COATINGS

Corrosion is prevented by the application of protective coating on the surface of metal, thereby
the metal surface is isolated from the corrosive environment. The coatings being chemically inert
to the environment under specific conditions of temperature and pressure, forms a physical
barrier between the coated surface and its environment. Coatings are not only preventing
corrosion but also decorate the surface of the metal.
Important types of protective coatings are:
(i) Inorganic coatings Ex: Metal coating
(ii) Organic coatings Ex: Paints, Enamels, Varnishes, Grease etc

METAL COATINGS

Corrosion of base metal can be prevented by coating a layer of another metal over it. Metal
coatings are of two types

1. Anodic metal coating

2. Cathodic metal coating
ANODIC METAL COATING

Coating a layer of metal which is anodic to base metal is known as anodic metal coating.
Ex: Coating of Zn, Al, Mg etc over iron/steel metals.
Galvanization or Galvanizing:
Galvanization is a process used for the protection of steel or iron from rusting. In this process, a
protective zinc coating is applied on the iron surface. The most common method of galvanizing
is to hot dip the metal in a bath of molten zinc. Zinc is more reactive metal than iron; hence it
reacts with oxygen to form a protective oxide layer, which prevents inner iron from getting in
contact with oxygen.
It involves the following steps.
1. Degreasing: Oil, grease on the metal surface is removed by washing with organic solvents
(CCl4, toluene)
2. Descaling: The metal surface is washed dilute sulphuric acid (Pickling process) to remove
dirt, rust on the surface.
3. Finally, the article is washed with water and air-dried.
4. Then passed through flux of ammonium chloride and zinc chloride to provide adhesive
property to the iron surface.
5. The article is then dipped in a bath of molten zinc. The excess zinc on the surface is
removed by passing through a pair of hot rollers.

Advantages
Galvanization provides excellent corrosion resistance to the iron metal surface.
Disadvantages: Galvanised articles cannot be used for storing or preparing food stuff since Zn
dissolves in acidic medium and forms toxic compounds.
Application: Galvanization of iron is carried out to produce roofing sheets, fencing wire,
buckets, bolts, nuts, pipes, building construction purpose etc.

CATHODIC METAL COATING

It is a process of coating of base metals with cathodic metals such as Sn, Ni, Cr, Cu etc.
Example: Tinning
TINNING:
Tinning is the process of coating a layer of tin metal on iron by hot dipping method.
1. Degreasing: Oil, grease on the metal surface is removed by washing with organic solvents
(CCl4, toluene)
2. Descaling: The metal surface is washed dilute sulphuric acid (Pickling process) to remove
dirt, rust on the surface.
3. Finally, the article is washed with water and air-dried.
4. Then passed through flux of ammonium chloride and zinc chloride to provide adhesive
property to the iron surface.
5. The article is then dipped in a bath of molten tin. The excess tin on the surface is
removed by passing through a pair of hot rollers.
6. Finally treated with palm oil to prevent the oxidation of tin.

Advantages
Tinning provides excellent corrosion resistance to the iron metal surface.
Tin is non toxic in nature.
Application: Tin coated articles can be used for storing or preparing food stuff. Copper utensils
coated with tin to prevent the contamination of food stuff with poisonous copper salts.
 DIFFERENCE BETWEEN GALVANIZING AND TINNING
Galvanizing Tinning
Coating of zinc on iron or steel, by hot dipping Coating of tin on iron or steel, by hot dipping
process is called galvanizing. (M.P of Zn = 419 oC) process is called tinning. (M.P of Sn =
The article is washed with organic solvents to 232oC).
remove oil/grease, with sulphuric acid to remove The metal surface is washed with organic
scale/rust, then with water and dried, before coating. solvents to remove oil/grease, with sulphuric
Coating metal is anodic to iron/steel, called acid to remove scale/rust then with water and
anodic coating. dried, before coating.
The molten metal bath is covered with a flux of Coating metal is cathodic to iron/steel, called
Ammonium chloride, which prevents the oxidation of cathodic coating.
the coated metal. The molten metal bath is covered by a
The article is dipped in a molten bath of Zn. The flux of Zinc chloride.
excess of coated metal is removed by passing through The clean and dry sheet is passed through
a pair of hot rollers and cooled gradually. flux layer, molten tin, finally removed out
Galvanizing is applied to nails, bolts, pipes, through palm oil, which prevents the
roofing sheets etc. oxidation of the coated tin.
Galvanized sheets cannot be used for It possesses more resistance against
preparing/storing food stuffs, since Zn dissolves in atmosphere.
acidic medium and forms toxic compounds. It is non-toxic in nature and more noble
If any crack is produced on the galvanized sheets, than the base metal.
causes severe corrosion on the coated Zn metal and Tinning is widely used for coating the
the base metal is protected. steel sheets, Cu and brass sheets used for
Zn is chosen as a protective coating for iron/steel manufacturing containers for storing/packing
because of its natural resistance against corrosion in food materials, cooking utensils, refrigeration
most atmospheric conditions, and Zn is equipments, etc.
electronegative to iron and can protect it sacrificially. If any crack is produced on the tinned
sheets, causes severe corrosion of the base
metal.
Tin coatings form a useful preparation for
protective painting in general applications.
 2. CATHODIC PROTECTION:

The principle is to force the metal to be protected, to behave as cathode. There are two types of
cathodic protections namely,
1) Sacrificial anode method
2) Impressed current method

SACRIFICIAL ANODE METHOD

The metallic structure to be protected is connected to a more anodic metal using a metallic wire.
The more active metal gets corroded, while the parent structure is protected from corrosion. The
more active metal so employed is called sacrificial anode. The sacrificial anodes to be replaced
by fresh ones as and when it is required.
Commonly used sacrificial anodes are: Mg, Zn, Al etc.
This method is generally used for the protection of buried pipelines, ship hulls, water tanks, etc.
Example: steel pipe is protected by connecting it to a block of Mg or Zn. In such cases steel acts
as a cathode and is unaffected Mg or Zn acts anode and undergoes sacrificial corrosion

IMPRESSED CURRENT METHOD

The metallic structure to be protected is made as cathode by impressing the current. The current
is applied in the opposite direction to nullify the corrosion current. The impressed current is
obtained from a source like battery. An insoluble anode (ex: graphite, high silicon content iron,
etc.) is buried in the soil, and connected to the structure to be protected. This method is suitable
for large structures and for long term operations. The anode is usually placed in a backfill, to
provide a better electrical contact with the surroundings.

3. ANODIC PROTECTION

“Protection of a metal from corrosion by developing a protective oxide film in oxidizing solution
by applying anodic current continuously from an external source is called anodic protection”.

Such a protection is possible on metals like Al, Cr, W, Ti, Ni and their alloys that get passivated
in a given environment.

Electrochemical measurements helps in finding out the optimum potential required for anodic
protection. Potentials are applied to the surface that is corroding and it is increased. The current
at various applied potential is measured. When logarithm of current (log I) is plotted against
applied potential (V) a typical current potential curve is obtained.

A plot of potential versus current curve shows the corrosion tendency and optimum
potential for anodic protection
When potential is increased, current also increases in the beginning. This is due to dissolution of
the metal at anode (corrosion). This continues until the potential reaches a critical value after
which the current drops to a smaller value. At the potential (Ep) the formation of a protective
layer occurs. If the potential is further increased, the metal remains un attacked up to a particular
potential. This is known as passive region. The optimum potential required for the protection can
be chosen in the mid way of passive region. For best results a metal should have a broad passive
region.

4. CORROSION INHIBITORS

Corrosion inhibitor is a chemical substance, when they added in small concentrations (e.g 0.1%)
to the corrosive environment, minimizes or prevents corrosion.

The inhibitor is chemically adsorbed on the surface of the metal and forms a protective thin film
with inhibitor effect or by combination between inhibitor ions and metallic surface.

The inhibitor leads a formation of a film by oxide protection of the base metal.

The inhibitor reacts with a potential corrosive component present in aqueous media and the
product is a complex.

The anodic type of inorganic inhibitor includes chromates, nitrites, molybdates and phosphates
and cathodic type includes zinc and polyphosphate inhibitors.

Mixed type is the major class of organic inhibitor as it includes amines, amine salt and
imidazoilnes – sodium benzoate mercaptans, esters, amines and ammonia derivatives.

Inorganic or anodic type of inhibitors

The addition of inorganic inhibitors causes suppression of electrochemical reaction at anodic–
cathodic areas. Most of the times, inhibitors are used in a blended form. These inhibitors only
react at an adequate level of concentration. Chromate Inhibitors

They are most effective inhibitors, but they are toxic and, hence, their application is restricted
and is not advised. In industrial water, the threshold concentration is 120 mg/L.

Cathodic inhibitors are chemical compounds which inhibit the cathodic reaction of the
corrosion cell (either liberation of H2 gas or formation of OH-).

Examples of cathodic inhibitors:-

Compounds such as ZnSO4, MgSO4 and Ca(HCO3)2, these compounds inhibit the cathodic
reaction by forming insoluble Zn(OH)2 film or Mg(OH)2 film or CaCO3 film with the
cathodically formed 𝑂𝐻− ions (in neutral solutions) the insoluble film isolates the cathodic
regions of the corrosion cells from the corrosive medium and stops corrosion. Addition of
oxygen scavengers or reducing agents such as Na2SO3 and hydrazine (N2H4) removes oxygen
from the media to minimize the formation of OH- ions.

Mixed Inhibitors

These types of corrosion inhibitors also form a film on the surface of the metal. They work to
reduce cationic reactions as well as anionic reactions. This is done via the formation of a
precipitate on the surface of the metal.
Examples for Mixed inhibitors include silicates and phosphates which are used as water
softeners to stop the rusting of water.

5. SPRAY TECHNIQUE
The coating metal in molten state is sprayed on base metal by means of spraying gun.

Surface engineering techniques generally consist of surface treatments, where the composition or
physical, chemical, and mechanical properties of existing surface are altered. They are achieved
by adding different material deposited as a thin or thick layers (coatings) to create a new surface.

Thermal Spraying is not a new process. It has proved itself to be extremely effective in the 90
years of its existence in all manner of applications ranging from coatings in gas turbines to
corrosion protection on park benches.
Thermal Spraying consist of
 Heating of the powder
 Melting of the powder
 Acceleration of the particles on to high velocity
 Impacting the particles on to the substrate and subsequent bonding
These include finishing coatings, such as anti-corrosion or decorative coatings, and engineering
coatings such as wear resistant and thermal barrier coatings.

There are four commonly used processes in thermal spraying;

1. Laser beam treatment.
2. Electron beam treatment.
3. Flame treatment.
4. Induction treatment.
It should be noted that surface engineering may be used effectively to improve the surface
properties, without changing the chemical composition of the bulk materials. It is frequently used
to satisfy engineering requirements. It provides an important means of increasing the life of
engineering component by increasing oxidation, corrosion, and wear and erosion resistance.
 Question bank

1. Define corrosion. Discuss the electrochemical theory of corrosion by taking rusting of

iron as an example.
2. Differentiate between electrochemical series and galvanic series.
3. Describe differential metal and differential aeration corrosion with an example.
4. Identify the type of corrosion the following.
a. Bolt and nut made from different metals are in contact with each other.
b. Deposition of small particles of dust on iron surface.
5. Pin holes on tin coated iron are more prone to corrosion than those on zinc coated iron.
Why?
6. Write a note on grain boundary corrosion or intergranular corrosion.
7. Describe crevice corrosion.
8. Give Nernst equation. How Nernst equation is useful in corrosion studies.
9. What Pourbiax diagram. List the characteristics. Describe Pourbiax diagram of Al metal
and Fe metal.
10. Define the following terms.
a) Polarization b) Passivity c) Immunity
11. Explain corrosion rate measurement by weight loss using Tafel extrapolation method.
12. What is metal coating? Distinguish between anodic metal coating (galvanization) and
cathodic metal coating (tinning).
13. What is cathodic protection? Explain sacrificial and impressed current method.
14. How metals can be protected by anodic protection method.
15. Write a note on corrosion inhibitors.
16. Write a short note on spray technique.
17. List and explain the factors affecting rate of corrosion.