3. INTRODUCTION
• Corrosion is the deterioration of a metal as a result of
chemical reactions between it and the surrounding
environment.
• It is an oxidation process. It causes loss of metal.
• The responsible factors for the corrosion of a metal
are the composition of the metal, the environmental
chemicals , temperature and the design.
4. • Example: Formation of
rust on the surface of
iron.
• Formation of green film
on the surface of copper.
5. Causes of corrosion
• Most metals exist in nature in combined forms that is in the
form of ore.
• Therefore metals in their finally refined state are highly
unstable and have a tendency to revert back in their original
state that is to stable state by the process of
corrosion(oxidation).
Ore of metal
(thermodynamically
stable)
(M n+)
Metal (M)
(thermodyna
mically
unstable)
Corroded metal
(Mn+ )
(more stable
than pure metal)
Extraction
Gain of e’s
(+ne)
corrosion
loss of e’s
(-ne)
6. So…why study corrosion?
• Materials are precious resources.
• Engineering design is incomplete without
knowledge of corrosion.
• Applying knowledge of corrosion can minimize
disasters.
• Corrosion products are thereat to environment.
7. Types Of Corrosion
Dry or Chemical Corrosion
Wet or Electrochemical Corrosion
It involves direct chemical attack of
atmospheric gases like CO2, O2, H2S,
SO2, halogen, moisture on metal. e g-
tarnishing of silver ware in H2S laden
air.
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.
8. Dry or Chemical Corrosion
Occurs
• Due to direct chemical reaction of atmospheric gases.
• Due to molten metal in contact with metal surface.
Types
• Oxidation corrosion
• Corrosion by gases
• Liquid Metal Corrosion
9. Types of dry corrosion
Oxidation corrosion
Occurs
• due to direct chemical reaction of atm. O2 with metal surface
forming metal oxide
• Absence of moisture
• Increases with increase in temp.
Mechanism
– on exposure to atm., metal gets oxidized to form metal ions
(i) M (S) Mn+ + ne- (if n=1)
– Electrons lost by metal are taken up by oxygen to forms oxide ions
(ii) 1/2O2 (g) + 2e- O2-
M2+ + 1/2O2 MO(metal oxide)
10. Types of dry corrosion
Corrosion by Other Gases
• 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.
• 2Ag + Cl2 AgCl
(Non-Porous layer)
• Fe + H2S FeS + H2
(Porous Layer)
• Sn + 2Cl2 SnCl4
(Volatile Layer)
11. Types of dry corrosion
LiquidMetal 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.
12. Wet or Electrochemical corrosion
Wet corrosion may be classified as:-
Uniform corrosion
Galvanic corrosion
Pitting corrosion
Stress corrosion
Conc. Cell corrosion
Dealloying
13. Uniform corrosion
• It is defined as the type of corrosion attack that
is more or less uniformly distributed over the
entire exposed surface of a metal.
• Cast irons and steel corrode uniformly when
exposed to open atmospheres.
14. Mechanism of uniform corrosion
• Anode :
• M (S) M+ + e- (Oxidation)
• cathode :
• 2H+ + 2e- H2 (g) (Reduction)
• 2M(S) + 2H+ 2M+ + H2
• In alkaline or neutral environment i.e. pH=7 or pH>7,
reduction of dissolved O2 is predominant cathodic process
that cause uniform corrosion.
• O2 + 2 H2O + 4e = 4OH-
•
15. Rusting of iron
• Rusting of iron in neutral aqueous solution of electrolytes in
presence of atmospheric oxygen.
• Anode :
• Fe Fe2+ + 2e- (Oxidation)
• Cathode :
• ½ O2 + H2O + 2e- 2OH- (reduction)
• Fe2+ + 2OH- Fe(OH)2
• In excess of O2 , ferrous hydroxide is easily oxidised to ferric hydroxide.
• 4Fe2+(OH)2 + O2 + 2H2O 4Fe(OH)3
• The product called rust corresponds to Fe2O3.XH2O
16.
17. Galvanic corrosion
• When two dissimilar metals are electrically connected and
exposed to an electrolyte, the metal higher in electrochemical
series undergoes corrosion.
• e.g., In the Zn-Cu galvanic cell , Zn act as anode where
oxidation & corrosion occurs & Cu act as a cathode & is
protected.
19. Pitting Corrosion
• Formed as a result of pit and cavities.
• Localized attack and formed by cracking
protective coating.
20. Pitting Corrosion
• Mechanism :
• Anode :
Fe = Fe2+ + 2e- (dissolution of iron)
• Cathode :
1/2O2 + H2O + 2e- = 2(OH-)
The positively charged pit attracts negative ions of
chlorine Cl- increasing acidity of the electrolyte
according to the reaction:
FeCl2 + 2H2O = Fe(OH)2 + 2HCl
21.
22. Stress corrosion
• Occurs in the presence of tensile stress and
corrosive environment
• E.g. brass get corrode in traces of ammonia.
23. Mechanism :
For stress corrosion to occur:
1. Presence of tensile stress.
2. Presence of a specific corrosive environment.
3. A susceptible material.
The corrosive agents are highly specific and
selective like:
(i) Caustic alkalis and strong nitrate solution for
mild steel.
(ii) Traces of ammonia for brass.
(iii) Acid chloride solution for stainless steel.
24.
25. Concentration cell corrosion
• It occurs when metallic surface is partially immersed
in an electrolyte and partially exposed to air.
• Mechanism:
• Anode :
• Zn Zn2+ + 2e-
• Cathode:
• 1/2O2+H2O+2e- 2OH-
• The Zn2+ and OH- ions interact to give Zn (OH)2
Therefore, corrosion occurs at anode.
26.
27. Dealloying
• Dealloying is a rare form of corrosion found in
copper alloys, gray cast iron & some other
alloys.
• Also called selective leaching.
• It is the removal of an element from an alloy
by corrosion.
• Common e.g. Dezincification
• Dezincification is the selective leaching of zinc
in brasses.
28. Dealloying
Dezincification
• If zinc content in a brass is less than 15% then
dezincification doesn’t occur significantly.
• If the zinc content is higher and acidic
environment is present dezincification will
occur.
• Prevented by keeping copper content above
85% , alloy substitution e.g. 1 % tin to brass.
30. Effects of corrosion :
(i) Loss of useful properties of metal and thus loss
of efficiency.
(ii) Decrease in production rate, because efficiency
is less and replacement of corroded equipment
or machinery is time consuming.
(iii) Increase in maintenance and production cost.
(iv) Contamination of product.
(v) Economic losses
31. Dry corrosion Wet corrosion
Corrosion occurs in the absence of
moisture.
Corrosion occurs in presence of
conducting medium.
It involves direct attack of chemicals
on the metal surface.
It involves formation of
electrochemical
cells.
The process is slow. It is a rapid process.
Corrosion products are produced at
the site of corrosion.
Corrosion occurs at anode but rust is
deposited at cathode.
The process of corrosion
is uniform.
It depends on the size of the anodic
part of metal.
Corrosion is the deterioration of a metal as a result of chemical reactions between it and the surrounding environment.
Corrosion is the reverse process of metallurgy.
Green film of basic carbonate CuCO3 + Cu(OH)3 =Cu(OH)2.CuCO3 On the surface of copper when exposed to moist air containing CO2.
Metals exist in nature in the combined form of carbonates, sulphides and sulphates.
These chemically combined states of metal "known as ore" has a low energy and is thus thermodynamically stable state of metal.
2. During the process of extraction a number of steps like concentration, Roasting ,and Smelting are involved and finally the ore is reduced into metal.
3. Corroded metal is thermodynamically more stable than pure metal but due to corrosion useful properties of a metal such a malleability, ductility and electrical conductivity are lost.
1. Chemical (or dry) corrosion: It involves direct chemical attack of atmospheric gases like CO2, O2, H2S, SO2, halogen, moisture and inorganic acid vapours on metal.Example, turnishing of silver ware in H2S laden air.
2. Electrochemical (or wet)corrosion: 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.
Generally more active metal act as anode and original metal act as cathode so anode is the area where corrosion occurs. Example, rusting of iron in moist atmosphere.
Atmospheric gases like o2,h2s,so2
1.Size of cation is smaller than anion.
consider a metal surface comes in contact in contact with atm. Gas O2 & metal oxide is formed on the metal surface.
These metal oxide act as the electrolyte & allow charge flow of metal ions & O2- ions
Here every metal part act as a conductor . But here we donot have moisture.
This formation of metal oxide indicates the formation of corrosion.
1.The extent of corrosion depends on the chemical affinity between metal and the gas involved.
2. The degree of attack depends upon the formation of protective and non- protective films on the metal surface.
If the film is protective or non porous the extent of attack decreases (e.g. AgCl).
If the film formed is non-protective or porous. The surface of the whole metal is gradually destroyed e.g. SnCl4 is volatile.(tin chloride)
In petroleum industry H2S at high temperature attacks- steel forming scale which is porous.
Amalgam- alloy of mercury with other metal
1.Uniform corrosion or general corrosion
1.This type of corrosion occurs in acidic medium. Anodic reaction is dissolution of metal into metal ions with liberation of electrons.
2.(cathode part): the electrons released flow through the metal from anode to cathode, whereas H+ of acidic solution are eliminated as H2 gas.
3.This type of corrosion causes displacement of H+ ions from sol. By metal ions
4.Metal have a tendency to get dissolved in acidic soln. with evolution of H2 Gas .
5.The anodes are large areas & cathodes are small areas.
1.Usually the surface of iron is coated with a thin film of iron oxide. If
the film develops cracks, anodic areas are created on the surface. While the metal parts act as cathodes.
2. Fe2+ ion originates at anode and OH- ions originate from cathode.
Smaller Fe2+ ions (ferrous ions)diffuse more rapidly than the larger OH- ions, so corrosion occurs at the anode, but corrosion product rust deposited near cathode.
3. Rust Fe2O3.XH2O , where x= is the no. of water molecules is variable
Zinc (higher in electrochemical series) forms the anode; whereas copper (lower in electrochemical series) acts as cathode.
Electrode potential of Zn= - 0.763
Electrode potential of Cu=+0.340
Thus the corrosion is a localized accelerated attack resulting in the formation of pits, holes
or cavities.
These Zn2+ will react with OH- ions to give Zn(OH)2
O2+2H2O+4e- = 4OH-(CATHODIC PART)
Pitting corrosion is a localized accelerated attack resulting in the formation of cavities around which the metal is relatively unattacked, therefore, pitting corrosion results in the formation of pinholes, pits and cavities in the metal. Pitting is, usually, the result of the breakdown or cracking of the protective film on the metal at specific points. This gives rise to the formation of small anodic and large cathodic areas.
2.If the presence impurities like sand, dust and scale is checked, pitting can be reduced.
3.If the concentration of oxygen is uniform, then also pitting is reduced.
Pitting corrosion of a stainless steel is illustrated in the figure.
Anodic reactions inside the pit: Fe = Fe2+ + 2e- (dissolution of iron) The electrons given up by the anode flow to the cathode (passivated surface) where they are discharged in the cathodic reaction: 1/2O2 + H2O + 2e- = 2(OH-) As a result of these reactions the electrolyte enclosed in the pit gains positive electrical charge in contrast to the electrolyte surrounding the pit, which becomes negatively charged. The positively charged pit attracts negative ions of chlorine Cl- increasing acidity of the electrolyte according to the reaction: Ferrous chloride + 2H2O = Ferrous hydroxide + 2HCl PH of the electrolyte inside the pit decreases (acidity increases) from 6 to 2-3, which causes further acceleration of corrosion process. Large ratio between the anode and cathode areas favors increase of the corrosion rate.
Stress corrosion is the combined effect of static tensile stresses and the corrosive environment on the metal stresses that cause cracking result from heavy working like rollling,drawing or insufficient annealing, welding, etc.
In such cases, the metal under stress becomes more anodic and tend to increase the rate of corrosion.
1. Stress corrosion involves in a localized electrochemical corrosion, occurring along narrow paths,
forming anodic areas with respect to the more cathodic areas at the metal surface.
2. Presence of stress produces strains, which result in localized zones of higher electrode potential.
3.These become so chemically-active that they are attacked, even by a mild corrosive environment, resulting in the formation
of attack, which grows and propagates in a plant, until failure occurs or it may stop, after progressing a
finite distance.
It occurs when metallic surface is partially immersed in an electrolyte and partially exposed to air.
anodic part: Poor oxygenated metallic part becomes anodic and undergoes oxidation.
cathodic part: Oxygen rich metallic part becomes cathodic. At the cathode, O2 takes up electrons to form OH- ions.
Zn is partially immersed in a dilute solution of a neutral salt e.g., NaCl.
2.Zn rod above and adjacent to the waterline are strongly aerated and hence
become cathodic.
3. Whereas parts immersed show a smaller oxygen concentration and become anodic.
4. So there is a difference of potential which causes flow of current between two differentially aerated areas of
same metal. Zinc will dissolve at anodic areas and oxygen will take up electrons at the cathodic areas forming hydroxyl ions.
Generally less noble metal is removed from alloy
Elements undergoing selective removal are zinc, aluminium, iron ,etc.
BRASS contains generally 70% Cu & 30% Zn ----dipped in dil.HCl
beginning BRASS colour changes from yellow to red.
Cu content increase from 70% to 92-95% in the surface. Cu surface is porous in nature .
So ??? Comes where is Zn??
After corrosion Zn reduces to 1-2% i.e. Zn has gone out onto the solution and copper has been left out
& this form a particular type of corrosion called dealloying. Since Zn is going out we call “dezincification” .
(if required) the rate at which dezincification takes place also affected by environmental temperature shows that increasing zinc % from 15 to 40 at 80°c increases the rate by 5 times.
..4.this process can be observed with naked eyes as colour changes from yellow(brass) to red(copper).
Prevented by keeping copper content above 85% , alloy substitution e.g. 1 % tin to brass.
