2. Dr. SURENDRAN PARAMBADATH
(M.Sc, M.Phil, M.Tech)
Formerly: Post Doctoral Research Associate,
Nano-Information Materials Research Laboratory,
Pusan National University, Busan-South Korea
Currently: Assistant Professor
Govt. Polytechnic College, Perinthalmanna
3. Nature keep all metals in their ore form.
Metals are usually extracted from their ores.
Nature tries to convert the metals again into their ore form.
The surface of metals are attacked when exposed to
environment.
Chemicals are chemically turned to new substances such as
oxides, hydated
oxides, carbonates, chlorides, sulphides, sulphates etc having
entirely new properties.
Once the metals begin to decay or get corroded, they suffer a
loss in electrical conductivity, tensile strength, color and luster.
4. Definition of Corrosion
Corrosion is the slow process of decay of the
metal, due to the attack of the atmosphere gases
on the surface of the metal, resulting in the
formation of metallic compounds such oxides,
hydroxides, carbonates, sulphides etc.
OR
Corrosion is the process of destruction or
deterioration of the metal and alloys by unwanted
or unintentional chemical or electrochemical
attack by its environments staring its surface.
5. Type of Corrosion
1. Dry or Chemical Corrosion:
It is due to the direct chemical action of
environmental atmosphere gases such as O2.
H2S, SO2, N2, halogens or anhydrous
inorganic liquids with metal surfaces.
2. Wet or Electrochemical Corrosion:
It is due to the existence of separate
“anodic” and “cathodic” areas in the system
between which current flows through the
conducting liquid and the anode gets oxidized
and wasted.
6. Electrochemical Theory of Corrosion
(Mechanism Of Corrosion)
Corrosion is an electrochemical process.
Galvanic cells are setup between dissimilar metals in contact with each
other or between dissimilar parts of he same metal when surrounded by
moist air or liquid.
Anodic area oxidation takes place
Cathodic area reduction takes place.
Metallic ions formed at the anodic part and ions formed at the
cathodic part diffuses towards each other through conducting
medium and the corrosion product is formed somewhere between
the anodic and cathodic areas.
7. Anodic iron gets oxidised to Fe2+
M Mn+ + ne-
The oxygen at the cathode changes to OH- by reduction.
½ O2 + H2O +2e- 2 OH-
9. 1. The position of metals in the electrochemical Series
Metal SRP, Eo
Lithium----------------- -3.05 V
More oxidation Less Reduction
Potassium
Calcium
Sodium
Decreasing Magnesium Increasing
tendency Aluminum order of std
to loose Zinc reduction
electrons Nickel potential
Tin
Hydrogen--------------- 0.00
Copper
Silver
Platinum
Gold---------------------- +1.15 V
10. When two metals are in contact with each other, in
presence of an electrolyte, the more active metal
becomes the anode and easily undergoes oxidation
ie, corrosion.
More active
Less active
Greater the difference between the reduction
potentials of the metals, more severe will be the rate of
corrosion.
11. 2. Relative areas of anodic and cathodic parts
When two dissimilar metals or alloys are in
contact and the anodic area is smaller than the
cathodic area, the corrosion of the anode become
rapid and severe.
Cathode
Anode
12. 3. Purity of the metal
A pure metal is more corrosion resistant than an
impure metal. The rate and extent of corrosion
increases with the amount of impurities present.
13. 4.Physical state of the metal
Smaller the size of the metal, more the area under
stress and greater is the corrosion.
14. 5. Solubility and volatility of corrosion product
If the product of corrosion is soluble in the
corroding medium and also is volatile, corrosion
occurs faster.
15. 6. Nature of the corroding environment
A.Temperature: Corrosion generally increases
with rise in temperature of environment.
B.Humidity of the air: Rate of corrosion
increases with presence of moisture in the
atmosphere.
C.Impurities: Presence of impurities like
CO2, H2S, SO2, acid fumes etc increases
corrosion rate.
D.Influence of pH: In acid medium corrosion is
more and in alkaline medium it is less.
18. Rust: Fe2O3.XH2O or Fe(OH)3
A galvanic cell is set up between two dissimilar parts of the same
metal iron.
a) The portion of iron which is in contact with water acts as
anode and the other portion in contact with air acts as cathode.
b) Anodic iron gets oxidised to Fe2+
Anode
Fe Fe2+ + 2e-
Cathode
The oxygen at the cathode changes to OH- by reduction.
½ O2 + H2O +2e- 2 OH-
The electron released at the anode move through the metal to the
cathodic site.
19. c) Fe2+ and OH- ions combine to form Fe(OH)2 which gets
oxidized to Fe(OH)3.
Fe2+ + OH- Fe(OH)2
4 Fe(OH)2 + O2 + 2H2O 4Fe(OH)3
Overall reaction
2 Fe + O2 + 4H+ 2Fe2+ + 2H2O
20. Conditions for rusting
1.Impurities in iron
2.Presence of oxygen
3.Presence of moisture
4.Presence of electrolyte
5.Presence of Cl2 or SO2 in the atm.
28. i. Metallic Coating
a) Using less active metal
b) Using a more active metal
Galvanization is the process of
coating iron or steel sheets with a
thin layer of zinc by dipping in
molten zinc.
29.
30. The iron or steel article is first cleaned with
dil. H2SO4 at 60-90oC (Pickling).
Then the article is treated with 5% HF to remove
grains of sand, washed with water and dried.
It is dipped in molten zinc (425-450oC).
A layer of zinc gets coated on the article, which
is then pressed through a roller to remove excess
zinc.
31. ii. Non-Metallic Coating
a) Phosphate coating by alkaline solution of phosphate.
b) Chromate coating using chromate solutions.
c) Anodizing on non-ferrous metals.
32. iii. Organic Coating
Plastics, polythene, rubber etc, are used for coating to
prevent corrosion.
Mainly on articles like, ship, submarines. Etc.
