Corrosion ppt part 1

Dr. Swastika Das
Professor, Chemistry
BLDEA’s Engineering College
Vijayapur.

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.
•Corrosion comes in many different forms and can be
classified by the cause of the chemical deterioration of a
metal.

Corrosion has been defined as
“a destructive chemical and electrochemical
reaction of a metal with its environment (like O2
, moisture, CO2 etc.) which disfigures metallic
products leading to reduction in their thickness
and also causes loss of useful properties such as
malleability, ductility, electrical conductivity and
optical refractivity.” Except few metals such as
gold, platinum (called noble metal), other
metals are prone to corrosion.

Rusting of iron
Rusting of Aluminium

 Metals are electropositive in nature. Except few metals
like gold, platinum (noble metal) other metals are found
in nature as their compounds (such as oxides,
hydroxides, carbonates, chlorides, sulphides,
phosphates, silicates etc.) which are called their ore.
Metals are thus obtained by extraction from their ores
by reduction process. In nature, when metals exist as
their compounds (or ore) they are stable and they are in
the low energy states. However, during extraction of
metals from their ores, free metals become less stable
and are in the higher energy state than in the ionic state.
So, metals have a tendency to go back to the ionic state
and hence metal atoms are prone to get attacked by
environment . This is the main reason for corrosion of
metals.

 Uniform or General attack
 Galvanic or Two-metal
 Crevice
 Pitting
 Intergranular
 Selective leaching or Parting
 Erosion corrosion
 Stress corrosion
 Hydrogen damage

 It is also known as uniform attack corrosion,
general attack corrosion is the most common type
of corrosion and is caused by a chemical or
electrochemical reaction that results in the
deterioration of the entire exposed surface of a
metal. Ultimately, the metal deteriorates to the
point of failure.
 Factors influencing corrosion
Corrosion is enhanced by :
the presence of impurities , air and moisture,
electrolytes , strains in metals like dent, scratches
etc.

 Dry or Chemical Corrosion
 Wet or Electrochemical Corrosion
Electrochemical theory :Two essential requirements
are
i) Formation of anodic and cathodic areas
ii) ii) Electrical contact between the cathodic and
anodic parts to enable the conduction of e-

It involves direct chemical attack of
atmospheric gases like CO2, O2 , H2S, SO2 ,
halogen, moisture and inorganic acid vapors
on metal.
Example, tarnishing of silver ware in H2S
laden air.

It occurs due to setting up of a large number of tiny galvanic
cells in metals in presence of an impurity as well as in
presence of moisture. Generally impurity (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.

 Oxidation corrosion
 Occurs due to direct chemical reaction of atm.
02 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------—> Mn+
+ ne_
Electrons lost by metal are taken up by oxygen to forms oxide ions (ii)
1/2O2 (g) + 2e ------—> O2-
2M + n/2O2 ------—> 2Mn+ + nO2- ------—> M2On
Metal Oxide
(corrosion product)

 Stable Layer: An impervious layer is formed, which checks
further oxidation corrosion. e.g. oxide films on Al, Sn, Pb,
Cu etc.
 Unstable Layer: Metal oxide decomposes e.g. in Au and Ag ,
Pt
 Volatile Layer: oxide layer volatilizes leaving theunderlying
metal surface for further attack. E.g. molybdenum oxide
(MoO3) is volatile.
 Porous Layer:- atmospheric O2 have access to the
underlying surface of metal. Alkali metals & alkaline earth
metals
Pilling - Bedworth Rule:
 A protective and Non-Porous metal oxide layer has volume equal to or
greater than the volume of metal from which it is formed.
 A Non-Protective and Porous metal oxide layer has volume lesser than
the volume of metal from which it is formed.
 Specific Volume Ratio = Volume of oxide formed /Volume of metal

 Occurs
o When a metal is contact with moist air or any liquid
medium
o When two diff. metals are partially immersed in a soln.
o Chemically non- uniform surfaces of metals behave like
electrochemical cells in the presence of water containing
dissolved O2 & CO2
o Always occurs at anodic areas
Mechanism Involves
 oxidation- reduction process depending on the nature of
corroding environment,
 electrons released at anode are consumed at the
cathodic area by two ways .
(i)Evolution of H2
(ii)Absorption of O2

 According to electrochemical theory of
corrosion, when a metal is exposed to an acidic
environment the process of corrosion sets in by
the formation of separate anodic and the
cathodic areas within the metal surface. A
driving force is necessary for electrons to flow
between the anodes and cathodes.
 Corrosion always occurs at the anodic area of
the metal due to the oxidation process and thus
electrons are liberated. The metal ions Mn+
formed during the destruction process of metal ,
either dissolves in the medium or forms a thin
film of oxide on the metal surface.

 The electrons liberated at the anodic area
flow through and are consumed at the
cathodic area by the following process:
(i) liberation of H2
(ii) Absorption of O2
 Corrosion current flows from anodic to
cathodic area.

i)Anodic reactions
M(s) ------—> Mn+ (aq) + ne- (oxidation)
Fe(s) ------—> Fe2+ (aq) + 2e- (oxidation)
ii) Cathodic reactions
a)In acidic solution in the absence of O2
2H+ + 2e- ------—> H2
b) In acidic solution in the presence of O2
O2 + 4H+ + 4e- ------—> 2H2O
c) In neutral or alkaline medium in the absence of O2
2H2O + 2e- ------—> H2 + 2OH-
d) In neutral or alkaline medium in the presence of O2
O2 + 2H2O + 4e- ------—> 4OH-

e.g. Rusting of iron occurs by O2 in the presence of aqueous
solution
At anode
Fe ------—> Fe2+ + 2e-
At cathode
½ O2 + H2O + 2e- ------—> 2OH-
Overall reaction
Fe + ½ O2 + H2O ------—> Fe2+ + 2OH- ---- Fe(OH)2
(i) In excess supply of oxygen: In excess supply of oxygen, ferrous
hydroxide is easily oxidized to ferric hydroxide.
2Fe(OH)2 + H2O + 1/2O2 ------—> 2Fe(OH)3 ------—>Fe2O3.xH2O
Yellow rust
(ii) In limited supply of oxygen:
Black magnetite Fe3O4 or ferroferric oxide is formed.
Fe(OH)2 ------—> Fe2O3.FeO.6H2O
Black Rust

 Hydrogen evolution:
Anode: Oxidation
M------—> Mn+
+ ne_
Cathode: Reduction
2H+ (g) + 2e- ------—> H2
Î
overall reaction:
2M + 2nH+ ------—> 2Mn+ + nH2

 Anode:
 M------—> Mn+
+ ne_
 Cathode:
O2 (g) + 4e- + 2H2O ------—> 4OH-
Overall reaction:
4 M +2nH2O ------—> 4Mn+ + 4OH-

 Galvanic corrosion (also called bimetallic
corrosion) is an electrochemical process in
which one metal corrodes preferentially to
another when both metals are in electrical
contact, in the presence of an electrolyte.
 This same galvanic reaction is exploited in
primary batteries to generate an electrical
voltage.
 A galvanic couple forms between the two
metals, where one metal becomes the anode
and the other the cathode.

Three conditions must exist for galvanic corrosion to
occur:
 Electrochemically dissimilar metals must be
present.
 The metals must be in electrical contact
 The metals must be exposed to an electrolyte.
Fe or iron
( -0.44V)
Cu or Copper
( +0.34V)
Zn Fe Zn
Corrosion
Fe
corrosion

 Different metals and alloys have different electrochemical potentials (or
corrosion potentials) in the same electrolyte.
 When the corrosion potentials of various metals and alloys are
measured in a common electrolyte (e.g. natural seawater) and are listed
in an orderly manner in a tabular form, a Galvanic Series is created.
 It should be emphasized that the corrosion potentials must be
measured for all metals and alloys in the same electrolyte under the
same environmental conditions (temperature, pH, flow rate etc.),
otherwise, the potentials are not comparable.
 The potential difference between two dissimilar metals is the driving
force for the destructive attack on the active metal (anode).
 The conductivity of electrolyte will also affect the degree of attack. The
cathode to anode area ratio is directly proportional to the acceleration
factor.

zn
c
u
Ccorrosive
environment
Anodic area
corroded
Cathodic
area
protected
soil
Anode: Zn ------—> Zn+2 + 2e-
Cathode: ½ O2 + H2O +2e- ------—> 2OH-
Overall Reaction: Zn(OH)2 is formed.

 Select metals/alloys as close together as
possible in the galvanic series.
 Avoid unfavorable area effect of a small
anode and large cathode.
 Insulate dissimilar metals wherever
practical.
 Apply coatings with caution. Paint the
cathode (or both) and keep the coatings in
good repair on the anode.
 Avoid threaded joints for materials far apart
in the galvanic series.

 Difference between the electrode potentials of the two metals.
The greater the difference the higher the driving force electric
force of corrosion.
Contact resistance at the boundary between the two metals. High
contact resistance limits the electrons transfer through the
boundary and decrease the corrosion rate.
Large anode connected to a small cathode results in slow
corrosion rate.
Electrolyte solution properties (pH, oxygen content, temperature,
flow rate etc).
Galvanic corrosion can be observed in the following
examples:
• Steel pipe connected to copper.
• Zinc coating on mild steel.
• Tin coating on copper vessel.
• Lead- antimony solder around copper wires.

 Differential aeration corrosion is a type of corrosion that occurs when
oxygen concentrations vary across a metal's surface. The varying
concentration of oxygen creates an anode and a cathode on the metal's
surface. Oxidation then occurs because an anode and a cathode have
been established on the surface.
 In differential aeration corrosion, the area with the higher oxygen
concentration becomes the cathode. The area with the lower oxygen
concentration becomes the anode. Consequently, the portion of the
metal that has the lower oxygen concentration is the portion subject to
corrosion.
 At the anode (less O2 concentration),
M --------- Mn+ + ne-
 At the cathode (more O2 concentration),
H2O + ½ O2 + 2e ------- 2OH-
 Some examples where varying concentrations of oxygen may be found are
metals that are buried, certain joint types, crevices and cracks. Metals
that are partially submerged in water are also subject to differential
aeration corrosion because the oxygen concentration in the water is
typically different from the oxygen concentration in the atmosphere.

 It takes place when a drop of electrolyte is in
contact with the metal surface. The metal
surface covered by is in contact with lesser
amount of air than the uncovered metal
surface. Thus, metal covered by drop
becomes anodic and corroded whereas the
uncovered metal becomes cathode.

Waterline corrosion is a case of differential aeration corrosion,
more prevalent in cases such as ocean going ships, water storage
steel tanks, etc., in which a portion of the metal is always under
water. The waterline corrosion takes place due to the formation
of differential oxygen concentration cells. The part of the metal
below the water line is exposed only to dissolved oxygen while
the part above the water is exposed to higher oxygen
concentration of the atmosphere. Thus, part of the metal below
the water acts as anode and undergoes corrosion and part above
the waterline is free from corrosion. A distinct brown line is
formed just below the water line due to the deposition of rust.

Pitting Corrosion is the localized corrosion of a metal
surface confined to a point or small area, that takes the
form of cavities. Pitting corrosion is one of the most
damaging forms of corrosion.
• Pitting corrosion is usually found on passive metals and
alloys such aluminium alloys and stainless alloys when
the ultra-thin passive film (oxide film) is chemically or
mechanically damaged and does not immediately re-
passivate.
• The resulting pits can become wide and shallow or
narrow and deep which can rapidly perforate the wall
thickness of a metal.

 For a defect-free "perfect" material, pitting corrosion is caused
by the environment (chemistry) that may contain aggressive
chemical species such as chloride. Chloride is particularly
damaging to the passive film (oxide) so pitting can initiate at
oxide breaks.
 • The environment may also set up a differential aeration
cell (a water droplet on the surface of a steel, for example)
and pitting can initiate at the anodic site (center of the water
droplet).
 • For a homogeneous environment, pitting is caused by the
material that may contain inclusions or defects. In most
cases, both the environment and the material contribute to
pit initiation.
 • The environment (chemistry) and the material (metallurgy)
factors determine whether an existing pit can be repassivated
or not.. An existing pit can also be repassivated if the material
contains sufficient amount of alloying elements such as Cr,
Mo, Ti, W, N, etc...

1. Ratio of anodic and cathodic area:
When two metals are in contact and dipped in a corrosive environment,
the corrosion of the anodic metal is directly proportional to the
ratio of anodic to cathodic area of the metals. The process of
corrosion is more rapid and the destruction of the metal is fasterif
the anodic area of the metal is small. The current density is high for
smaller anodic area and the demand for electrons at larger cathodic
area is met by severe and faster corrosion rate at the anodic area.

 2. Electrode potential values of the metals:
The rate and severity of corrosion of metals depends on
the difference of electrode potentials or the position of
the metals in the galvanic series. When dissimilar metals
are in contact, dipped in a corrosive medium, the
corrosion at the anodic region occurs with a faster rate,
because the anodic metal is higher up in the galvanic
series. Greater the electrode potential difference of the
metals, faster is the rate of anodic metal.
e.g., corrosion of zinc metal is severe and faster than
iron metal when it is in contact with copper in a
medium. This is due to the greater electrode potential
difference between Zn and Cu than Fe and Cu.

3. pH of the medium:
The rate of corrosion of metal is much faster in an
acid medium than alkaline or neutral condition.
Corrosion rates almost always increase with
decreasing pH(increasing acidity). This is due to the
increase in the concentration of H+ ions and
increasing the solubility of most potentially corrosive
products.
4. Humidity and Temperature:
A humid environment favours the corrosion of
metals. Iron does not rust easily when exposed to dry
air, however, it undergoes rusting when exposed to
humid conditions. The atmospheric gases like CO2,
SO2 etc gets dissolved in water and produces a
medium that sets up an electrochemical cell in
metal. The rate of corrosion of metal enhances with
the rise in temperature.

5. Nature of the corrosion products:
The corrosion rate of a metal proceeds with a faster
rate, if the corrosion product formed is soluble in the
medium to which it is immersed. If the corrosion
product is insoluble, then it is likely to cover the
entire metal surface as continuous thin film coating
inhibiting the corrosion process.
In an oxygen environment, if the tarnished metal
gets a weak, thin, porous and non adherent film of
oxide the metal is susceptible for corrosion.
However, in few cases of metals, the oxide film so
formed is nonporous, tough and adherent on the
surface of the metal, which protects the metal from
further corrosion.

 If the surface of the metal is protected by
any means, then the corrosion can be
controlled. The following methods are
usually adopted to control corrosion.
 Proper selection and designing of metal
equipment, fabrication etc.
 Apply protective coatings like metal/
inorganic/ organic coatings.
 Cathodic and anodic protection
 Corrosion inhibitors.

 Galvanization (or galvanizing as it is most
commonly called in that industry) is the
process of applying a protective zinc coating
to steel or iron, to prevent rusting. The most
common method is hot-dip galvanizing, in
which parts are submerged in a bath of
molten zinc.

Galvanising protects the base metal in three
ways:
 It forms a coating of zinc which, when intact,
prevents corrosive substances from reaching the
underlying steel or iron.
 The zinc serves as a sacrificial anode so that
even if the coating is scratched, the exposed steel
will still be protected by the remaining zinc.
 The zinc protects its base metal by corroding
before iron. For better results, application of
chromates over zinc is also seen as an industrial
trend.

 Hot-dip galvanizing deposits a thick, robust
layer of zinc iron alloys on the surface of a
steel item. In the case of automobile bodies,
where additional decorative coatings of paint
will be applied, a thinner form of galvanizing
is applied by electrogalvanizing. The hot-dip
process generally does not reduce strength
on a measurable scale, with the exception of
high-strength steels where hydrogen
embrittlement can become a problem. This
deficiency is a consideration affecting the
manufacture of wire rope and other highly-
stressed products.

 The protection provided by hot-dip galvanizing is insufficient
for products that will be constantly exposed to corrosive
materials such as acids, including acid rain in outdoor uses.
For these applications, more expensive stainless steel is
preferred. Some nails made today are galvanized.
Nonetheless, electroplating is used on its own for many
outdoor applications because it is cheaper than hot-dip zinc
coating and looks good when new. Another reason not to use
hot-dip zinc coating is that for bolts and nuts of size M10 (US
3/8") or smaller, the thick hot-dipped coating fills in too
much of the threads, which reduces strength (because the
dimension of the steel prior to coating must be reduced for
the fasteners to fit together). This means that for cars,
bicycles, and many other light mechanical products, the
practical alternative to electroplating bolts and nuts is not
hot-dip zinc coating, but making the fasteners from stainless
steel.

“ The process of coating molten zinc on the
base metal (iron) surface by hot dipping is
known as galvanization”.
1. The base metal is cleaned and degreased using organic
solvents.
2. The metal is then treated with dil. H2SO4( pickling) for
10 minutes at about 800 C to remove any rust or scales.
3. The metal is further treated with flux material- ZnCl2
and NH4Cl for best adhesion property.
4. Finally, the metal part is dipped in hot molten zinc at
430- 4700 C.
5. The excess zinc is removed from the surface of the
coated metal by rolling, wiping or air blow techniques.

 Degreasing A hot (90oC) caustic bath is used to
remove - Oil, grease, paint another organic
compounds If these contaminating materials are not
removed, the next stage of the process (pickling) is
affected Check if any paint is on item - some paints
are not easily removed and may require abrasive
blasting before galvanizing Some pipes and pipe
fittings have a black varnish coating that cannot be
removed in the caustic bath.
 Acid pickling The steel items are immersed in
hydrochloric acid to remove - rust, mill scale and
other metal oxides The steel surface must be
perfectly clean of these oxides for the molten zinc to
react with the steel Very heavy rust may not be easily
removed by pickling - badly rusted items should be
abrasive blasted or mechanically cleaned first.

 Pre-fluxing : Steel items are water rinsed after pickling
and immersed in hot (70-80oC) zinc ammonium
chloride (ZAC) solution. The ZAC solution conditions
the clean steel surface ready for hot dip galvanizing.
Good pre-treatment = good quality galvanizing
 The steel is immersed in molten zinc (temp 4500C)
The clean steel surface reacts with the molten zinc to
form a zinc-iron alloy which is very strongly bonded to
the surface. The hot dip galvanized coating forms in 3-5
minutes, depending on the steel thickness.
 Quenching : After hot dip galvanized steel item is
removed from the galvanizing bath, it is immediately
quenched in a sodium dichromate solution The
dichromate quenching cools the item so that it can be
quickly handled and conditions the surface of the
galvanized coating to maintain its bright appearance.
Dichromate quenching will reduce white rusting
problems. White rust forms when rainwater reacts with
newly galvanized steel to form zinc hydroxide.

Corrosion ppt part 1

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Corrosion ppt part 1