PGE 519 – Corrosion Principles,
Prevention and Control
Presentation 1 – Introduction to Corrosion
COURSE CONTENTS
• Definition and Classification of Corrosion
• Mechanism of Corrosion of Metals
• The Electrochemical Theory of Corrosion
• Forms of Corrosion
• Stress Corrosion and Corrosion Fatigue
• Methods of Corrosion Control and
Prevention
CORROSION
• Corrosion can be defined as the destruction or deterioration of
material because of reaction with its environment.
• The study of corrosion must include the following:
• Knowing what the reaction involves (reactants).
• The products of the reaction.
• The mechanism of reaction in order to avoid severely corrosive
conditions and adopt protective measures against it.
• 5-20% of the ferrous metals produced annually are destroyed by
corrosion.
EFFECTS OF CORROSION ON MATERIALS
• Appearance: Badly corroded or rusted material (or equipment)
in a plant will leave a poor impression on the observer. Aesthetic
values of equipment are lost and appearance of equipment
becomes repulsive to the user.
• Cost: Corrosion leads to direct expenses linked to the cost of
protective measures and replacement of corrosive material. A
good share of the material may be recovered as scrap but
a large part must be replaced by new metal.
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• Plant Shut Downs: Plants are frequently shut down or
portion of process stopped due to unexpected corrosion
failures.
• Contamination of Products: Many products such as foods
and drugs are often contaminated by corrosion products.
• Effect on Safety and Reliability: Corrosion products can
contaminate consumable items such as water, milk and
other dairy products thereby risking human health.
CLASSIFICATION OF CORROSION
PROCESSES
• Chemical Corrosion: This is a corrosion in which a metal interacts
with a medium which does not conduct electricity. Example,
interaction between a metal and oxygen (particularly at high
temperature), halogens, hydrogen, Sulphur, gases, etc.
• Electrochemical Corrosion: This is a corrosion in which oxidation-
reduction processes takes place at local areas where one area
serves as the anode and the other as the cathode. Example
corrosion of steel in water, dissolution of insoluble anodes, corrosion
of pipelines carrying current conducting fluids.
• Biochemical Corrosion: This is caused by the activities of
microorganisms using metals as culture medium. Soil, stagnant
waters and certain organic products usually favour biocorrosion.
MECHANISM OF CORROSION OF
METALS
• The presence of moisture and oxygen are important factors
in the process.
• Very pure water can attack iron or steel in the presence of O2
but the corrosion is slow and usually tends to stifle itself due
to the low solubility of the oxides, of iron which are found
on the surface.
• In the presence of a salt solution (such as sodium chloride)
and an electric current the corrosion can be greatly
accelerated.
THE BASIC WET CORROSION
CELL
• The Anode: It usually corrodes by loss of electrons from
electrically neutral atoms to form discrete ions. This ions
may remain in the solution or react to form insoluble
corrosion products. The corrosion reaction of a metal is
usually expressed by the simplified reaction:
𝑀 → 𝑀𝑍+ + 𝑍𝑒−
𝑀 − 𝑀𝑍+ → 𝑍𝑒−
• The Cathode: The cathode does not normally corrode. Two
important reactions which may occur at the cathode depending
on the pH of the solution are:
pH < 𝐻+ + 𝑒− → 𝐻 (𝑎𝑡𝑜𝑚)
7 2𝐻 → 𝐻2 (𝑔𝑎𝑠)
pH > 2𝐻2𝑂 + 𝑂2 + 4𝑒− → 4𝑂𝐻−(Oxygen reduction, neutral
7 or
basic)
• Electrolyte: Is a solution which must of necessity conduct
electricity.
• Electrical Connection: The anode and cathode must be in
electrical contact for a current to flow in the corrosion cell.
Example Daniel Cell.
Zn / Zn2+ // Cu2+ / Cu
THE ELECTROCHEMICAL THEORY OF
CORROSION
• The theory is based on the fact that all the metals corrode when
they discharge positive ions into solution.
• This leaves the metal with a negative charge or potential.
• The tendency for this to occur is based on the relative
positions of metals in the electrochemical series.
• Metals which are higher up will dissolve while those lower
will be displaced from the aqueous solutions of their salt.
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• At the Anode
𝑍𝑛 − 2𝑒− → 𝑍𝑛2+(𝑂𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛)
𝐹𝑒𝐶𝑙2 → 𝐹𝑒2+ + 2𝐶𝑙−
• At the cathode
2𝐶𝑙− + 2𝑒− → 𝐶𝑙2(𝑅𝑒𝑑𝑢𝑐𝑡𝑖𝑜𝑛)
• The electrons produced at the anode flow to the cathode
to neutralize the gas ions thus forming chlorine gas
(Cl2) which is liberated at the cathode.
