UNIT-II

CORROSION AND ITS CONTROL METHODS

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Dr.A.Geetha

Associate Professor & HEAD

Department of Chemistry

Kongu Engineering College

UNIT – II Corrosion and its Control Methods

Corrosion: Introduction - chemical corrosion – Pilling-Bedworth rule - electrochemical corrosion and it's types – galvanic corrosion�– differential aeration corrosion with examples - galvanic series - factors influencing rate of corrosion– measurement of corrosion�(wt. loss method only).

�Control methods – sacrificial anodic protection method - corrosion inhibitors - protective coatings - pretreatment of metal surface�– metallic coating: electroplating, electroless plating and hot dipping (tinning and galvanizing) methods – non-metallic coating:�anodizing - organic coating: paints, constituents and functions - ceramic coatings

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�Corrosion�

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• The serious consequences of the corrosion process have become a problem of worldwide significance.

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• In addition to our everyday encounters with this form of degradation, corrosion causes plant shutdowns, waste of valuable resources, loss or contamination of product, reduction in efficiency, costly maintenance and expensive overdesign.

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Corrosion can be defined as the destruction of metals or alloys by the surrounding environment through chemical or Electrochemical changes.

Causes of corrosion

• 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 exists 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 are become less stable and are in the higher energy state than in the ionic state.
• So, metals have a tendency to 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.

Example : Rusting of iron (Fe₂O₃.H₂O)

Formation of green layer on copper surface

More stable

Low energy

Less stable

High energy

Types Of Corrosion

1. Dry or Chemical Corrosion

2. Wet or Electrochemical Corrosion

Dry or Chemical Corrosion

• Dry corrosion involves direct attack of atmospheric gases on metal in the absence of moisture.
• Oxygen is mainly responsible for the corrosion of most metallic substances. Other gases like CO₂,H₂S, SO₂, Cl₂ also responsible.

Example: Reaction between metal and oxygen

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Types of dry or chemical corrosion:

I. Corrosion by oxygen

II. Corrosion by other gases

III. Liquid –Metal corrosion

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I. Oxidation corrosion

• Due to direct chemical reaction of atmospheric O₂ with metal surface forming metal oxide
• Absence of moisture
• Increases with increase in temperature

Mechanism

On exposure to atmosphere., metal gets oxidized to form metal ions

M M²⁺  + 2e^-

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

½ O₂ + 2e^- O^2-

M + ½ O₂ 2M²⁺ + O^2- MO

Metal Oxide

Nature of metal oxide film

A. Stable film

• Fine grained structure
• Formation of Oxide films are impervious in nature
• Stop further oxygen attack
• Stable film acts as a protective coating. Ex: Heavy metals
• Eg: Al, Pb, Cu, Sn
• HEA

Eg : Ag, Au, Pt - no oxidation corrosion

Noble metals

• Film produced on the surface of noble metals.
• Which decompose reversibly to the metal and the oxide which is liberated in the form of oxygen.

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c. Volatile film

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Eg : Molybdenum oxide (MoO₃) - rapid corrosion

• The oxide layers formed in some cases is volatile.
• Oxide film volatilizes as soon as it is formed.
• The fresh metal surface is kept exposed all the time for further attack.
• Thus it act as a nonprotective coating. Eg: MoO₃ (Molybdenum oxide film)

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d. Porous film

The oxide layers formed in some cases have pores or cracks.

The atmospheric oxygen can easily move into the metal surface through pores of the layer.

So corrosion is a continuous process.

Eg: Alkali metals & alkaline earth metals.

Ex: Oxides of Li, Na, K, Mg, Ca, Sr, etc.

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Pilling – Bed Worth Rule

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Specific Volume Ratio = Volume of oxide layer formed

Volume of metal consumed

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• A protective and Non-Porous:

If volume of metal oxide layer is equal to or greater than the volume of metal from which it is formed. Ex: Oxides of Al, Sn, Pb, Cu, etc.

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• A Non-Protective and Porous

If volume of metal oxide layer is lesser than the volume of metal from which it is formed. Ex: Oxides of Li, Na, K, Mg, Ca, Sr, etc.

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• If Pilling-Bedworth ratio is more than 1, oxide layer is protective and non-porous.

Ex: Oxides of Al, Sn, Pb, Cu, etc.

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• If Pilling-Bedworth ratio is less than 1, oxide layer is non-protective and porous.

Ex: Oxides of Li, Na, K, Mg, Ca, Sr, etc.

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II. Corrosion by Other Gases

• Other gases such as Cl₂,SO₂,H₂S,F₂,CO₂ & NO_X. In dry atmosphere

these gases react with metal and form corrosion products.

Hydrogen Embrittlement (or) Hydrogen Corrosion

At high temperature

Fe + H2S FeS + 2H

[H] + [H] H₂

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• FeS is a Porous Layer and corrosion is a continuous process.
• Atomic hydrogen penetrates and occupies the voids of the metal. Then

develops pressure which leads to cracking of the metal.

• This is known as Hydrogen Embrittlement (or) Hydrogen Corrosion

Decarburisation:

• Atomic hydrogen is highly reactive and combine with Carbon present in metal forming CH₄ gas which leads to cracking of the metal surface.
• The process of decrease in carbon content in steel is known as decarburation of steel.

C + H CH₄

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• Observed in Nuclear power plant. In nuclear reaction where Na metal used as a coolant leads to corrosion of Cd. Liquid metal mercury dissolves most metals by forming amalgam.

� Wet or Electrochemical or Immersed Corrosion �

Occurs

• When a metal is contact with moist air or any liquid medium
• When two diff. metals are partially immersed in a soln.
• Chemically non- uniform surfaces of metals behave like electrochemical cells in the presence of water containing dissolved O₂ & CO₂
• 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:
• Evolution of H₂
• Absorption of O₂

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Hydrogen evolution Mechanism

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• This types of corrosion takes place when base metals are in contact � with acidic solutions or the solutions are completely free from

dissolved oxygen.

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• All metals above hydrogen in the electrochemical series have a � tendency to get dissolved in acidic solution with simultaneous evolution � of hydrogen gas.

Anode :

Fe Fe²⁺ + 2e^- (Oxidation)

Cathode :

2H⁺ + 2e^- H₂ (g) (Reduction)

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Over all Rxn :

Fe + 2H⁺ 2Fe²⁺ + H₂ (g)

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Oxygen Absorption Mechanism

• Base metal are in contact with neutral, aqueous or slightly alkaline solution with

some amount of dissolved oxygen.

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Anode :

Fe Fe²⁺ + 2^e- (Oxidation)

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Cathode :

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1/2O₂ + 2^e- + H₂O 2OH^- (Reduction)

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Over all rxn:

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Fe + H₂O + 1/2O₂ Fe²⁺ + 2OH^- Fe(OH)₂

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Fe(OH)₂ + O₂ + H₂O Fe(OH)₃

(or)

Fe₂O₃.3H₂O

^RUST

Difference Between Dry and Wet Corrosion

Types of Electrochemical Corrosion

• Galvanic (Bimetallic) Cell Corrosion.

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• Concentration cell corrosion (or) Differential aeration corrosion.

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Galvanic (Bimetallic) Cell Corrosion:� Galvanic Corrosion occurs at

• Electrochemically dissimilar metals must be present.
• These metals must be an electrical contact .
• The metals must be exposed to an electrolyte.

The metals higher in electrochemical series act as the anode and dissolves into the solution and the metal lower in the electrochemical series acts as the cathode.

Cathodic reaction � 1. Electrolyte in the acidic medium- Evolution of Hydrogen Gas� 2. Electrolyte in the Neutral (or) Alkaline medium – Absorption of oxygen

������������������������������������������������������������������������������������������������������������������������������������������������������������������������������������������������������������������������������������Examples: �1. Fe & Cu� Fe-Anode, Cu- cathode

(Iron undergoes corrosion whereas Copper is not affected)�Anode :� Fe Fe²⁺ + 2^e- (Oxidation)�Cathode :�1/2O₂ + 2^e- + H₂O 2OH^- (Reduction)�� Over all rxn:� Fe + H₂O + 1/2O₂ Fe²⁺ + 2OH^- Fe(OH)_2�^� Fe(OH)₂ + O₂ + H₂O Fe(OH)₃ � (or) � Fe₂O₃.3H₂O�

^RUST������2.

Other examples

2. Zn & Fe

Zn-Anode, Fe- cathode

(Zinc undergoes corrosion whereas Iron is not affected)

Anode :� Zn Zn²⁺ + 2^e- (Oxidation)�Cathode :�1/2O₂ + 2^e- + H₂O 2OH^- (Reduction)�� Over all rxn:� Zn + H₂O + 1/2O₂ Zn²⁺ + 2OH^- Zn(OH)_2�^� �3. Buried iron pipeline connected to copper plumbing�4. Steel pipe connected to Copper plumbing�5. Lead-Tin solder around copper wires.�

Measures to prevent Galavanic Corrosion

• When joining two dissimilar metals together, galvanic corrosion can be prevented by insulating the two materials from each other. Ex: when bolting flanges of dissimilar metals together, plastic washers can be used to separate the two metals.

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• Do not couple metals that are far apart in the emf series. Selection of metals should be very close in the emf series.

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• Avoid small anode–large cathode combinations.

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Concentration Cell Corrosion �or �Differential Aeration Corrosion

• Occurs due to diff. in potential between differently aerated areas.
• Part of metal exposed to air is more oxygenated part & acts as cathode
• Part of metal immersed in electrolyte is poorly oxygenated & acts as anode

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Types of Differential Aeration Corrosion

• Pitting Corrosion
• Pipeline Corrosion
• Crevice Corrosion
• Corrosion on wire fence

Pitting Corrosion�

Pitting corrosion is a localized form of corrosion by which cavities or "holes" are produced in the material. Pitting is considered to be more dangerous than uniform corrosion damage because it is more difficult to detect, predict and design against. Corrosion products often cover the pits.

Preventive measures:

1.Preparing the surfaces with mirror finish.

2. Removing all contaminants.

3. Designing and fabricating to

avoid trapped and pooled liquids.

Pipeline corrosion

Pipeline corrosion is the oxidisation and electrochemical breakdown of the structure of a pipe used to convey any substance. Pipeline corrosion occurs on both the inside and outside of any pipe and related structures, exposed to corrosive elements. Many different types of corrosion can develop in pipelines.

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Crevice Corrosion

Crevice Corrosion refers to the localized attack on a metal surface at, or immediately adjacent to, the gap or crevice between two joining surfaces. The gap or crevice can be formed between two metals or a metal and non-metallic material.

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���������� Corrosion on wire fence� The areas where the wires cross are less aerated than the other parts of the fence. The corrosion takes place at the wire crossings because the less aerated part act as anode.�

Galvanic Series

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• The series prepared by studying the corrosion of metals and alloys in a given environment like sea water.
• Different testing metals and alloys are coupled with a calomel electrode and dipped in sea water.
• The oxidation potential of different metals and alloys are determined at 25°C and tabulated in the descending order.
• Metals occupying higher positions in the series undergo corrosion in a vigorous manner.

Need for Galvanic series

• Electrochemical series does not account for the corrosion of all metals and alloys.

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Ex: Zn/Al couple

• Zn (below Al in the EMF series) is corroded
• While Al acts as cathode and protected.

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Difference between Electrochemical and Galvanic series

FACTORS INFLUENCING RATE OF CORROSION

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The nature and extent of corrosion depend on the metal and environment.

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1. Nature of Metal

a. Nature of the matal (Position of metal in emf series)

• Metals higher in the emf series possess high reactivity becomes anodic and undergoes corrosion

b. Relative areas of anode and cathode

• large cathodic area and small anodic area leads severe corrosion

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c. Nature of surface film

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• Specific Volume Ratio = Volume of oxide formed

Volume of metal

If SVR > 1; film is nonporous and protects metal from corrosion

if SVR < 1; film is porous and increases the corrosion rate

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d. Surface state of the metal

Impurity creates heterogeneity and leads galvanic corrosion and results in high corrosion rate.

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e. Hydrogen overvoltage

The metal low hydrogen overvoltage more susceptible for corrosion results in faster the cathodic reaction to increase the corrosion.

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II. Nature of Environment

a. Temperature

Corrosion rate increases with temperature. Increase in temperature increases the conductance of the medium, reduces passivity of the metal, and thereby increase the rate of corrosion.

b. Humidity

Humidity-concentration of water vapour in the atmosphere.

Corrosion rate increases with humidity

Critical humidity-sudden increase in the corrosion rate at a

particular point

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c. Corrosive gases

presence of gases like H₂S,SO₂ etc., enhances the corrosion rate

d. Suspended Solids

presence of suspended solids in environment enhances the corrosion rate

e. pH

Lower the pH of metal environment higher is the corrosion rate. Acidic environment increases corrosion rate than alkaline environment

f. Conductance of the medium

The presence of conducting species in the atmosphere increase the corrosion rate. This rate is more in a wet atmosphere (more conducting) than in dry atmosphere.

g. Polarization of electrodes

Departure of electrode potentials from their equilibrium values. Polarization at the electrodes is due to concentration changes in the electrode region.

Polarization occurs in cathode- Cathode polarization- Reduces the rate of corrosion

Polarization occurs in Anode- Anode polarization- Increase the rate of corrosion

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