UNIT – IV:

Corrosion and its control:

Introduction causes and effects of corrosion, mechanism of chemical and electrochemical corrosion.
Types-differential aeration corrosion (Pitting and waterline corrosion), differential metal corrosion
(Galvanic corrosion).
Factors affecting corrosion-nature of metal (position, passivity, purity, areas of anode and cathode) &
nature of environment (temperature, pH, humidity).
Corrosion control methods - cathodic protection-Sacrificial anodic and impressed current cathodic
protection, comparison of galvanizing and tinning.

Corrosion
Causes
Most of the metals (except noble metals) exist in nature in combined form (ores).
Extraction (reduction)
Ores Metals
Natural tendency (oxidation
Combined state/Lower energy state Excited state/Higher energy state

When metals are exposed to environment-Moisture / Dry Gases / liquids- Exposed metal surfaces began
to decay more or less rapidly
Corrosion process is exactly reverse of extraction of metals

“Corrosion is the deterioration/destruction and consequent loss of solid metallic material, through
unwanted chemical or electrochemical attack by its environment, starting at its surface”.

Ex: i) Rusting of iron -formation of reddish scale and powder of hydrated ferric oxide on the surface.
ii) Formation of green film of basic copper carbonate on the surface of Cu, when exposed to moist-air
containing CO2.

Effects of Corrosion
 The annual world wide cost -metallic corrosion -$2-2.5 billion.
 Enormous wastage of metal (25 %) in the form of its compounds.
 The valuable metallic properties like conductivity, malleability, ductility etc., are lost due to
corrosion.
 The failure of the machinery takes place due to lose of useful properties of metals.
 Life span of the metallic parts of the machineries is reduced.
 Leakage in the process - Health & fire hazard
 Causes contamination

Theories of Corrosion
Broadly the corrosion process proceeds in two types by direct chemical action of environment on the
surface of metal in absence of moisture and in presence of a conducting liquid by the formation of
electrochemical cells.
1. DRY CORROSION
2. WET CORROSION
 1. Dry or Chemical corrosion
Occurs
 Due to direct chemical reaction of atmospheric gases
 Due to molten metal in contact with metal surface
Types
i) Oxidation corrosion
ii) Corrosion by gases
iii) Liquid Metal Corrosion

i) Oxidation corrosion
 Occurs due to direct chemical reaction of atm. O 2 at lower or higher temperatures with metal
surface forming metal oxide
 Absence of moisture
 Increases with increase in temp.
 Alkali metals –(Na, Li, K etc) and alkaline metals (Mg, Ca, Sn etc.) are rapidly oxidised at low
temperature.
 At high temperature, almost all metals (except Ag, Au and Pt) are oxidized.
Mechanism
i) On exposure to atm. O2, metal gets oxidized to form metal ions
M → 2M+ + 2e- (Oxidation)

(ii) Electrons lost by metal are taken up by oxygen to forms oxide ions

1/2 O2 +2e- → O-2 (Reduction)

Net reaction : 2M +1/2 O2 → M2O

Metal Oxide

• Oxidation occurs at metal surface and resulting metal oxide layer forms a barrier and restricts
further oxidation.
• For further oxidation occurs only when either the
i) Metal must diffuse outwards through the scale or
ii) Oxygen must diffuse inwards through the scale to the underlying metal
Generally Metal outward diffusion is much rapid than Oxygen inward diffusion because of their smaller
size, metal ions have higher mobility than the oxygen ions.
Nature of Oxide formed plays an important part in oxidation corrosion process
 Metal + Oxygen → Metal oxide (corrosion product)
When oxidation starts, a thin layer of oxide is formed on the metal surface and the natureof this film
decides the further action.
Types of Layers
1) Stable Layer:
The oxide films of Al, Pb, Cu, Sn etc. are stable, tightly adhering and can be impervious in nature.
Behaves as protecting coating and prevents further oxidation

2) Unstable Layer
The oxide layer ( Ag, Au & Pt) formed, decomposes back into the metal and oxygen. Consequently,
oxidation corrosion is not possible.
Metal Oxide Metal + Oxygen

3) Volatile Layer
The oxide layer (MoO3) volatilizes as soon as it is formed and leaves the fresh surface. This causes rapid
and continuous corrosion.

4) Porous Layer
Oxide films (Alkali metals & alkaline earth metals) have pores and cracks, so easy access for oxygen
and oxidation corrosion continues till the entire metal is completely converted into its oxide

Pilling - Bed Worth Rule

This rule is postulated to express the extent of protection given by the corrosion layer.
According to this rule, smaller the specific volume ratio, greater is the oxidation corrosion.
If the volume metal oxide layer is equal/greater than the volume of metal from which it is formed is non-
porous and becomes protective by tightly adhering to the metal from which it is formed.
Ag + Cl2 → 2AgCl (protective film)

Ex: Specific volumes of W (3.6), Cr (2.0) and Ni (1.6) – Corrosion is least in W.

If the volume of the metal oxide layer is lesser than the volume of metal, the metal oxide layer is non-
protective, porous and non-continuous and faces stress and strains and develop pores and cracks and
permits oxygen to reach the underlying meal, so the corrosion is continuous and rapid.
Fe (steel) + H2S → FeS (porous)

ii) Corrosion by gases

• Corrosion by gases like SO2, CO2, Cl2, H2S, F2 etc.
• Extent of corrosive effect depends on the chemical affinity of metal and the gas involved
• The degree of attack depends on the formation of protective and non-protective films on the metal
surface.
a) Protective or non-porous film protects the metal from further attack so intensity/ extent of attack
decreases
Ex: AgCl layer formed by the attack of Cl2 on Ag Film

b) Non protective or porous- surface of the whole metal is gradually destroyed

Ex: In petroleum industry, H2S at high Temp attacks steel forming porous FeS scale

iii) Liquid Metal Corrosion

 Occurs due to the action of flowing liquid metal at high temp on solid metals or alloys.
 This type of corrosion occurs in devices used for nuclear power. Either by
a) Dissolution of solid metal by a liquid metal
b) Internal penetration of liquid metal into the solid metal
Both these modes of corrosion cause weakening of the solid metal.

Ex: i) In nuclear reaction where Na metal used as a coolant leads to corrosion of Cd.
ii) When Liquid Hg diffuse along with the small defects and grain boundaries of an alloy (brass), it
reduces its ductility and strength.

2. Wet or Electrochemical corrosion

It is a common type of corrosion of metal in aqueous corrosive environment.
Occurs
i) when the metal comes in contact with a conducting liquid
ii) when two dissimilar metals/alloys are immersed or dipped partly in a solution.
We have to consider the following
 The formation of a galvanic cell on the surface of metals, i.e., formation of anodic and cathodic
areas/parts, between which current flows through the conducting medium.
 The chemical in the environment and humidity acts as an electrolyte/conducting medium.
 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.

Mechanism of Wet or Electrochemical corrosion

This corrosion involves flow of electron-current between anodic and cathodic areas
At anode- Dissolution of metals as corresponding metallic ions with liberation of electrons
At cathode- Consumes electrons by a) Evolution of H2 or b) Absorption of O2, depending on the nature
of the corroding environment
a) Evolution of H2-
Occurs usually in acid medium
At anode Fe dissolves as ferrous ions with liberation of electrons
The liberated electrons flows from anodic to cathodic areas, through iron metal, where H + are eliminated
as hydrogen gas.

At Anode Fe → Fe2+ + 2e- Oxidation

At Cathode 2H+ + 2e- → H2 ↑ Reduction
Overall Reaction Fe + 2H+ → Fe2+ + H2 ↑

 In this type of corrosion, anodic areas are very large and cathodic areas are small.
 The metals above Hydrogen in the electrochemical series have a tendency to get dissolved in the
acid solution with simultaneous evolution of H2.
 This type of corrosion causes displacement of Hydrogen ions from the acidic solution by the
metal ions
In Neutral solution Evolution of hydrogen:
At Anode : Fe → Fe2+ + 2e‾ (Oxidation)
At cathode : 2H2O +2e‾ → H2 ↑ + 2 OH‾ (Reduction)
Overall Reaction Fe2+ + 2OH‾→ Fe (OH)2
If oxygen is in excess, ferrous hydroxide is easily oxidized to ferric hydroxide.
4Fe (OH) 2 + O2 + 2H2O → 4Fe (OH) 3 → Fe2O3. 3H2O.
This type of corrosion causes displacement of hydrogen ions from the solution by metal ions. All metals
above hydrogen in electrochemical series have a tendency to get dissolved in acidic solution with
simultaneous evolution of H2 gas. The anodes are large areas, whereas cathodes are small areas.
b) Absorption of O2
Rusting of iron in neutral, aqueous solution of electrolytes (NaCl), in the presence of Atm. O 2. The
surface of Fe coated with thin iron oxide film.

 Iron oxide film develops some cracks, anodic areas are created on the surface, while well metal parts
act as cathodes.
 Anodic areas are small and cathodic areas are large
 At anode Fe dissolves as ferrous ions with liberation of electrons
 The liberated electrons flows from anodic to cathodic areas, through iron metal, where e - are
intercepted by dissolve oxygen.
 The ferrous ions (A) and hydroxyl ions (C) diffuse and when they meet, ferrous hydroxide is
precipitated.
 The combination occurs, more commonly near the cathode, because the smaller Fe +2 ions diffuse
more rapidly than the larger OH- ions.
 So corrosion occurs at the anode, but rust is deposited at or near the cathode.

At Anode Fe → Fe+2 + 2 e- Oxidation

At Cathode ½ O2 + H2O + 2 e- → 2OH- Reduction
Overall Reaction Fe+2 + 2OH- → Fe(OH)2 ↓
a) If oxygen is in excess, ferrous hydroxide is easily oxidized to ferric hydroxide.

4Fe (OH) 2 + O2 + 2H2O → 4Fe (OH) 3

The product called brown/yellow rust corresponds to Fe2O3. 3H2O.

b) If the supply of oxygen is limited, the corrosion product may be black rust (black anhydrous
magnetite), Fe3O4 or [FeO. Fe2O3].

 With increase in oxygen content, it forces the cathodic reaction to produce more OH - ions thus
consuming more electrons by accelerating the corrosion at anode.
 The presence of oxygen accelerates the corrosion at anodic area and rust formation at cathodic area.

Types of corrosion

Concentration cell corrosion

Galvanic Corrosion or bimetallic Corrosion

Concentration cell corrosion

This type of corrosion occurs due to electrochemical attack of the metal surface exposed to electrolyte
of varying concentrations or varying aeration.

Differential aeration corrosion

 Most common type of Concentration cell corrosion.

 Corrosion occurs, when one part of the metal is exposed to a different air concentration from the other
part.
 Poorly oxygenated part of the metal becomes anodic and well oxygenated part of the metal becomes
cathodic.
 Difference in potential is created, which causes the flow of electrons between the two differently
aerated parts of the same metal.

Differential aeration corrosion accounts for corrosion of metals, partially immersed in a

solution, just below the waterline.

If Zn rod is partially immersed in dilute solution of NaCl and the solution is not agitated
properly, then, the parts above and closely adjacent to the waterline are more strongly aerated because
of easy access of oxygen and hence become cathodic and the parts immersed to greater depth have less
access of oxygen show a smaller oxygen concentration and thus become anodic.

The circuit is completed by migration of ions, through the electrolyte, and flow of electrons,
through the metal, from anode to cathode.
+2 -
At Anode Zn → Zn + 2 e Oxidation
At Cathode ½ O2 + H2O + 2 e- → 2OH- Reduction

Waterline corrosion
 Corrosion in storage tanks, H2O tanks, marine structures, ship hulls etc., is called water line corrosion.

At Anode 2Fe → 2Fe+2 + 4 e- Oxidation

At Cathode 4H2O → 4H+ + 4OH- Reduction

4H+ + O2 + 4 e- → 2H2O

Overall Reaction 2Fe+2 + 4OH- → 2Fe(OH)2

 When water is stored in an iron tank, it is found that maximum corrosion occurs along a line just
below the water level/ meniscus.
 The area above the H2O line is highly oxygenated and acts as cathode and is completely unaffected
by corrosion, while the area just below the H2O line is less oxygenated and acts as anode and
undergoes corrosion.
 However, if the water is relatively free from acidity, little corrosion takes place.
 In the case of ships, this can be accelerated by marine plants attaching themselves to the side of the
ship. This can be restricted to some extent by the using special antifouling paints.
Pitting corrosion

At Anode Fe → Fe+2 + 2 e- Oxidation

At Cathode ½ O2 + H2O + 2 e-→ 2OH- Reduction
Overall Reaction Fe+2 + 2OH- → Fe(OH)2 ↓
Pitting corrosion is a localized accelerated attack resulting in the formation of pin holes, pits and cavities
around which the metal is relatively unattacked.

It is due to the breakdown or cracking of the protective film on the metal at specified points.
 This gives rise to the formation of small anodic and large cathodic areas.

 Pitting corrosion is explained by considering a drop of water or brine solution (aqueoussolution

of NaCl) on a metal surface, (especially iron).
 In this case, the small part below the impurity acts as the anodic area while the rest of the metal
acts as the cathode area. Due to corrosion a small pit is formed at the anodic area which grows
gradually.

 Once a small pit is formed the rate of corrosion will be increased because the demand for
electrons by the large cathodic area could be satisfied only when the anode undergoes drastic
corrosion.

 Presence of external impurities like sand, dust, water drops etc. on the surface of the metal can
also be a cause for this type of corrosion.
Reasons/ causes for breakdown of the protective film

Surface roughness or non-uniform finish, Scratches or cut edges, Local strain of the metal,
Alternating stresses, Sliding under load, Impingent attack (turbulent flow of a solution over a metal
surface), Chemical attack.
Ex: i) Carry over in boiler causing corrosion to turbine plates- Boiler water concentrated with dissolved
salts

ii) Corrosion due to caustic embrittlement - Boiler water has Na2CO3 (converts to NaOH)

Differential metal corrosion

Galvanic / Bimetallic corrosion

When two dissimilar metals or alloys are electrically connected and exposed to an electrolyte, the metal
higher in electrochemical series under goes corrosion. This type of corrosion is called galvanic or
bimetallic corrosion.

Thus when Zn and Cu are connected,

Zn being higher in the series forms the anode and is attacked and dissolved and undergoes
corrosion, whereas Cu which is lower in the series act as cathode and gets protected.
Mechanism
In acidic solution- Hydrogen evolution process
In neutral/ slightly alkaline solution- Oxygen absorption process
The electron –current flows from the anodic metal, Zn to cathodic metal Cu.

Zn → Zn+2 + 2 e- Oxidation
Corrosion occurs at Zn rod, and Cu rod is protected.
Examples
1) Steel pipe connected to copper plumbing
2) Lead – antimony solder around copper wire
3) Steel screws in a brass marine hardware
4) A steel propeller shaft in bronze bearing

Factors affecting corrosion-

1) Nature of metal
a) Position of the metal in galvanic series
When two metals or alloys are in electrical contact, in presence of an electrolyte, the more active
metal/ alloy (which is higher up in the series) will undergo corrosion easily.
• Rate and severity of corrosion, depends upon the difference in the position of metals.
• The greater is the difference, the faster and higher is the corrosion.
Ex: Au Cu Pb Fe Zn Al Mg Na K- Reactivity increases from left to right.

b) Purity of the metal

Lesser is the percentage purity of a metal, faster is the rate of corrosion.
The impurities present in a metal cause heterogeneity and thus tiny electrochemical cells are set up at the
exposed part of the impurity and corrosion of metal around the impurity takes place due to local action.
Ex: Zn metal containing impurity such as Pb or Fe undergoes corrosion due to the formation of local
electrochemical cells.
The rate or extent of corrosion increases with the increasing exposure and the extent of the
impurities. Consequently, corrosion resistance of a metal may be improved by increasing purity
c) Passivity of the metal
Tl, Al, Cr, Mg, Ni & Co are passive and they exhibit much higher corrosion resistance
than expected from their position in galvanic series, due to formation of highly protective, but very thin
film (of oxide) on the metal/alloy surface.
The film is self-healing in nature, if broken, repairs itself on re-exposure to oxidizing conditions.
Passivity means the lack of activity under conditions where a metal would be expected to react readily.
 There are certain metals which are passive to certain corroding agents.
Ex: Iron is passive to conc. HNO3.
Al has no action with conc. HNO3 in absence of chlorides etc.
 Metals which are susceptible to corrosion are made passive by alloying with one or more of metals
which are passive or resist to corrosion.
Ex: Iron is rendered passive by alloying it with any of the transition metals such as chromium, nickel
and molybdenum • (e.g. stainless steel)
Ni based alloys, Al alloys are used in many engineering applications
 The quantity of passive metal to be added depends upon the environments.
Ex: lron-chromium alloy, when the quantity of chromium added is from 12 to 20%, the alloy is
passive to neutral salt solution but no concentration of chromium in iron makes the alloy passive to
hydrochloric acid.

d) Areas of anode and cathode

 The rate of the corrosion is greatly influenced by the relative sizes of cathodic and anodic areas.
 Rate of corrosion is more if the size of cathodic area is larger, the demand for electrons will be more
and this results in the increase of rate of oxidation of metal at anode i.e., high rate of corrosion at
anode.
 When two dissimilar metals or alloys are in contact, the corrosion at anodic part is directly
proportional to the ratio of areas of the cathodic part and the anodic part.

Ex: 1) Steel pipe (corroded) fitted in a large copper tank

2) Tin (Sn) coated on iron (Fe) and in that some area is 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.

3) 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.

2) Nature of environment
a) Temperature
With increase of temperature of environment, the reaction as well as diffusion rate increase, thereby
corrosion rate is generally enhanced.
b) pH
Corrosion of the metals which are readily attacked by acids can be reduced by increasing the pH of the
attacking environment.
Ex: Zn corrodes rapidly even in weak acidic solutions, H 2CO3, but suffers minimum at pH=11.
Generally acid media, at pH < 7 are more corrosive than alkaline and neutral media.

Ex: Corrosion rate of iron in oxygen-free water is slow, until the pH < 5.

c) Humidity of air
 Humidity is the deciding factor in atmospheric corrosion.
 The greater is humidity, the greater is the rate and extent of corrosion. This is due to the fact that the
gases O2, H2S, SO2, CO2 and vapour present in atmosphere furnish water to the electrolyte, essential
for setting up an electrochemical corrosion cell.
 The nature of water source may also play an important role.
Rain water, not only supplying the necessary moisture, but also wash away the good part of oxide film
from the metal surface which leads to enhanced corrosion (Ag, Cr form adherent oxide film).

“Critical humidity is the relative humidity above which the atmospheric corrosion rate of metal
increases sharply.”
This value depends on the physical characteristics of the metal as well as nature of the corrosion products.

Corrosion control methods

Cathodic protection

The method of protecting the base metal by making it to behave like a cathode is called as cathodic
protection. There are two types of cathodic protection
(a) Sacrificial anode method
(b) Impressed current method.
a. Sacrificial anode method
□ In this protection method, the metallic structure to be protected (base metal) is connected by a wire
to a more anodic metal so that all the corrosion is concentrated at this more anodic metal.
□ The more anodic metal itself gets corroded slowly, while the parent structure (cathodic) is
protected. The more active metal so employed is called sacrificial anode. The corroded sacrificial
anode is replaced by a fresh one, when consumed completely.
□ Metals commonly employed as sacrificial anode are Mg, Zn, Al and their alloys which possess
low reduction potential and occupies higher end in electrochemical series.
Eg. A ship-hull which is made up of steel is connected to sacrificial anode (Zn-blocks) which undergoes
corrosion leaving the base metal protected.
The underground water pipelines and water tanks are also protected by sacrificial anode method.
By referring to the electrochemical series, the metal with low reduction potential is connected to the
base metal which acts as anode.

Fig. Sacrificial anode method: ship hull and underground water pipeline

Applications:
Suitable for applications / structures that are buried / placed in the soil with low resistivity
1) Protection of buried pipelines
2) Underground cables
3) Water tanks
4) Marine corrosion of cables, ship-hulls and piers
5) Insertion of Mg sheets into the domestic water boilers to prevent the formation of rust.
6) Calcium metal is employed to minimize engine corrosion.

Advantages:
1) Low installation and operating cost.
2) Capacity to protect complex structures.
3) Applied to wide range of severe corrodents.
Limitations:
1) High starting current is required.
2) Uncoated parts cannot be protected.
3) Frequent monitoring and replacement of sacrificial anode block should be done
4) Limited driving potential, hence, not applicable for large objects.

b. Impressed current method

For larger structures, galvanic anodes cannot economically deliver enough current to provide
complete protection.
□ In this method, an impressed current is applied in opposite direction to nullify the corrosion
current, and convert the corroding metal from anode to cathode

□ The impressed current is slightly higher than the corrosion current. Thus the anodic corroding
metal becomes cathodic and protected from corrosion.
□ The impressed current is taken from a battery or rectified on A.C. line.
□ The anode is usually insoluble anode like graphite, high silica iron, scrap iron, stainless steel, or
platinum.
□ A sufficient DC current is applied to an inert anode, buried in the soil (or immersed in the corroding
medium) and connected to the metallic structure to be protected.
□ The anode is, usually, a back fill, composed of coke breeze or gypsum, so as to increase the
electrical contact with the surrounding soil.
The impressed current protection method is used for water tanks, water & oil pipe lines,
transmission line towers etc.

• This kind of protection technique is particularly useful for large structures for long term
operations
Advantages:
 Unlimited driving potential.
 Capable of protecting large steel structures when designed properly.
 Requires less anodes then a galvanic system.
 Output can be controlled using a permanent reference electrode, desirable when the
 electrolyte resistivity is known to change.

Limitations:
 Initial costs can be more expensive.
 Requires an external DC power source along with an AC supply.
 System requires routine maintenance and monitoring.
 Anode wires can be susceptible to damage.

Metallic coatings: The surface of the base metal is coated with another metal (coating metal).
Metallic coatings are broadly classified into anodic and cathodic coatings.

Anodic coating:

The metal used for the surface coating is more anodic than the base metal which is to be protected.
For example, coating of Al, Cd and Zn on steel surface are anodic because their electrode potentials
are lower than that of the base metal iron. Therefore, anodic coatings protect the underlying base
metal sacrificially.
The formation of pores and cracks over the metallic coating exposes the base metal and a
galvanic cell is formed between the base metal and coating metal. The coating metal dissolves
anodically and the base metal is protected.
Cathodic coating:
 Cathodic coatings are obtained by coating a more noble metal (i.e. metals having higher electrode
potential like Sn, Au, Ag, Pt etc.) than the base metal. They protect the base metal as they have
higher corrosion resistance than the base metal due to cathodic nature.
 Cathodic coating protects the base metal only when the coating is uniform and free from pores.
 The formation of pores over the cathodic coating exposes the base metal (anode) to environment and
a galvanic cell is set up. This causes more damage to the base metal.

S.No. Galvanizing Tinning

1 Anodic coating Cathodic coating

2 Process of covering iron or steel with a thinProcess of coating steel with a thin coat of TIN to
coat of ZINC to prevent it from rusting. prevent it from corrosion
3 Zinc protects iron sacrificially, Since it is moreTin protects the base metal iron, from corrosion due
electro-positive than iron and does not permitto its noble nature and higher corrosion resistance.
iron to pass into the solution.

4 In galvanized articles, zinc continues toTin protects underlying iron till the coat is intact.
protect the underlying iron by galvanic cellAny break in coating causes rapid corrosion of iron.
action, even if the coating of zinc is broken at
any place

5 Galvanized containers cannot be used forTin coated containers and utensils can be used for
storing acidic foodstuffs as zinc reacts withstoring any food stuff as tin is non-toxic and
food acid forming poisonous compounds protects metal from corrosion

6 Ideal temperature is around 450 OC Ideal temperature is around 250 OC

7 Ammonium chloride is used as flux Zinc chloride is used as flux

8 Galvanized articles are good engineeringTinned articles are used majorly in food processing
materials industries