1. Engineering Chemistry
Dr. Thathsara D. Maddumapatabandi
thathsaram@sltc.ac.lk
Module code:
Credits: 3.00
Mid Exam: 30%
Quizzes: 10%
Laboratory Experiments: 10%
Final Exam: 50%
Recommended Text books:
1. Chang, R. and Goldsby, K.A. (2015). Chemistry, 12th edition. NY: McGraw Hill.
2. Agarawal, S. (2015): Engineering Chemistry, Fundamentals and Applications, 2nd Edition, Cambridge
3. Brown L. S. and Holmes T. A. (2011): Chemistry for Engineering Students, 2nd Edition, Cengage Learning.
4. Yen T. F. (2008):Chemistry for Engineers, Imperial College Press, London.
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2. Course Outline
1) Handling numbers in chemistry
2) Atomic theory & electron configuration
3) Theory of chemical bonding
4) Chemical Kinetics
5) Chemical Equilibrium
6) Chemical Thermodynamics
7) Electrochemistry
8) Corrosion
9) Water & water treatment
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3. 3
We see many examples of corrosion around us. Rust on iron,
tarnish on silver, and the green patina formed on copper and
brass are a few of them. Corrosion causes enormous damage to
buildings, bridges, ships, and cars.
Metals degenerate in the presence of moisture and air.
The process of spontaneous deterioration or disintegration of metals
(except gold and platinum) caused by direct chemical or indirect
electrochemical attack by reaction with the environment.
What is Corrosion?
4. 4
Why does corrosion occur?
• Metals exist in nature in combined forms like oxides, sulphides, sulphates,
carbonates, etc.
• These combined states (ores) are thermodynamically stable states of the metal.
• Energy is supplied to extract such metals from their ores. The extracted metal
is at a higher energy level and hence it is in a thermodynamically unstable
state. Metals try to get back to their stable states by combining with other
elements, and in this process, corrosion occurs and oxides, sulphides,
chlorides, sulphates, and so forth are formed.
• Although corroded metals are thermodynamically more stable than pure metals,
corrosion affects the useful properties of the metals like malleability, ductility and
electrical conductivity.
5. 5
A region of the metal’s surface serves as the anode, where oxidation
occurs:
The electrons given up by iron reduce atmospheric oxygen to water at the
cathode, which is another region of the same metal’s surface:
6. 6
Note that this reaction occurs in an acidic medium; the H ions are supplied in
part by the reaction of atmospheric carbon dioxide with water to form H2CO3.
The Fe2+ ions formed at the anode are further oxidized by oxygen:
This hydrated form of iron(III) oxide is known as rust. The amount of water
associated with the iron oxide varies,
so we represent the formula as Fe2O3 .xH2O.
8. 8
Theories/Mechanism of Corrosion
Dry or chemical corrosion
Wet or electrochemical corrosion
Dry or chemical corrosion
Also called direct corrosion, it occurs because of direct chemical action of the
environment on the metal surface in the absence of moisture or a liquid
electrolyte.
It generally occurs at a temperature higher than 100 °C when it is not possible to
develop a surface water layer. It is generally of three types:
9. 9
Oxidation corrosion It is brought about by the direct action of oxygen on the
metal surface at high or low temperature. At low temperature, alkali and
alkaline earth metals are oxidized and at high temperatures except Ag, Au and
Pt, all other metals are oxidized. It generally takes place in the absence of
moisture.
Corrosion by other gases In this, dry gases like H2, Cl2, F2, SO2 and CO2 directly
attack the metal surface. Rate of corrosion depends upon the chemical affinity
between the metal and gas and also on the nature of the film formed.
Liquid metal corrosion It occurs when an anhydrous liquid attacks the metal
surface.
When a liquid metal flows over a solid metal at high temperature it weakens the
solid metal because of
• Its dissolution in liquid metal.
• Penetration of liquid metal into solid metal. For example, sodium metal
(coolant) leads to the corrosion of cadmium in a nuclear reactor.
10. 10
Wet or electrochemical corrosion
It occurs when A metal is in contact with a conducting liquid or
When metals of different reactivities are in contact with a solution.
Rusting of iron is the most common example of electrochemical
corrosion.
This type of corrosion occurs under wet or moist conditions, and is more
common than dry corrosion.
11. 11
Factors Influencing Corrosion
The rate and extent of corrosion depend on the following
• Nature of the metal
Purity of the metal, Physical state of the metal etc
• Nature of the corroding environment.
Temperature, Humidity, Presence of impurities in atmosphere, Effect of pH etc
12. 12
Protection against Corrosion (Corrosion
Control)
The various methods by which corrosion can be controlled and prevented
are as follows
1. Protection by proper designing
2. Material selection
3. Modifying the environment
4. Modification of properties of the metal
5. Use of inhibitors
6. Cathodic protection or electrochemical protection
7. Use of protective coating.
14. 14
Cathodic protection
Sacrificial Anodic protection :
• The metal structure can be safe from corrosion by connecting it with a
copper(Cu) wire to a more active metal so that all the corrosion is
concentrated at the more active metal .
• As the more active metal is sacrificed in the process of saving metal from
corrosion . Hence it is known as sacrificial anode .
• Metals commonly used as sacrificial anodes are Magnesium , Zinc and
their alloys .
• Zinc is used as sacrificial anode in good electrolytes such as sea water .
• Magnesium is used as a sacrificial anode in high resistivity electrolytes
such as soil .
Applications :
1) Protection of underground cables and pipe lines from soil corrosion .
2) Protection of ships and boats from marine corrosion sheets of Magnesium/Zinc
are hung around the SHIP HULL .
3)The formation of rusty water or water corrosion is prevented by insertion of
Magnesium/Zinc sheets or rods into the domestic water tank .
15. 15
Cathodic protection of an iron storage tank (cathode) by magnesium,
a more electropositive metal (anode). Because only the
magnesium is depleted in the electrochemical process, it is
sometimes called the sacrificial anode.
19. 19
The tendency for iron to oxidize is greatly reduced when it is alloyed with certain
other metals. For example, in stainless steel, an alloy of iron and chromium, a layer
of chromium oxide forms that protects the iron from corrosion.
An iron container can be covered with a layer of another metal such as tin or
zinc. A “tin” can is made by applying a thin layer of tin over iron. Rust formation is
prevented as long as the tin layer remains intact. However, once the surface has been
scratched, rusting occurs rapidly. If we look up the standard reduction potentials,
according to the diagonal rule, we find that iron acts as the anode and tin as the
cathode in the corrosion process:
The protective process is different for zinc-plated, or galvanized, iron. Zinc is more
easily oxidized than iron
Please read the separate pdf document for Protection against
Corrosion (Corrosion Control) uploaded on LMS
when a piece of iron or an article made of iron is left in the open, it develops a reddish-brown coating. If this is left unattended for a long time, then the metal becomes weak and brittle and breaks off. A similar phenomenon is observed in copper that develops a greenish coating; similarly silver loses its luster and so on.
Rusting of iron and formation of a green layer of basic carbonate on the surface of copper as discussed above are the common examples of corrosion. Note that the term
rusting applies to the corrosion of iron and iron base alloys. Nonferrous alloys corrode but do not rust.
Because standard reduction potential for the reduction of Fe2+ is less positive than that for the reduction of O2, Fe can be oxidized by O2.
Metallic corrosion is not limited to iron. Consider aluminum, a metal used to make many useful things, including airplanes and beverage cans. Aluminum has a
much greater tendency to oxidize than iron does; we see that Al has a more negative standard reduction potential than Fe. Based on this fact alone, we might
expect to see airplanes slowly corrode away in rainstorms, and soda cans transformed into piles of corroded aluminum. These processes do not occur because the layer of
insoluble aluminum oxide (Al2O3) that forms on its surface when the metal is exposed to air serves to protect the aluminum underneath from further corrosion. The rust that
forms on the surface of iron, however, is too porous to protect the underlying metal.
Some of the metals directly react with oxygen in the absence of moisture.
Alkali and alkaline earth metals react with qxygen at room temperature and form corresponding oxides, while some metals react with oxygen at higher temperature.
Metals like Ag, Au and Pt are not oxidized as they are noble metals.
During oxidation of a metal, metal oxide is formed as a thin film on the metallic surface which protects the metal from further corrosion.
If diffusion of either oxygen or metal is across this layer of metal oxide plays an important role in the process of corrosion.
Wet or electrochemical theory of corrosion
It’s a common type of corrosion of metal in aqueous corrosive environment. 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.
A number of methods have been devised to protect metals from corrosion. Most
of these methods are aimed at preventing rust formation. The most obvious approach
is to coat the metal surface with paint. However, if the paint is scratched, pitted, or
dented to expose even the smallest area of bare metal, rust will form under the paint
layer. The surface of iron metal can be made inactive by a process called passivation.
A thin oxide layer is formed when the metal is treated with a strong oxidizing agent
such as concentrated nitric acid. A solution of sodium chromate is often added to
cooling systems and radiators to prevent rust formation.
Underground pipelines are often protected against corrosion by making the pipeline the cathode of a voltaic cell.Pieces of an active metal such as Mg are buried along the pipeline and connected to it by wire as shown in figure.
In mist soil where corrosion can occur, the active metal serves as the anode, and the pipe experiences cathodic protection.
Underground pipelines are often protected against corrosion by making the pipeline the cathode of a voltaic cell. Pieces of an ative metal such as Mg are buried along the pipeline and connected to it by wire as shown in the figure. In moist soil where corrosion can occur the active metal serves as the anode, and the pipe experiences cathodic protection.
The rusting of
iron exposed to the weather—even
support iron underneath concrete—
is an example of uniform corrosion.
Iron is oxidized, and oxygen from
the air is reduced. Water is needed
for ion mobility between the anodic
and cathodic regions, and the
presence of ionic salts speeds the
reaction considerably
Iron is often covered with a coat of paint or another metal such as Sn or Zn to protect its surface against corrosion. Covering the surface with paint or Sn is simply a means of preventing oxygen and water from reaching the iron surface. If the coating is broken and the iron is exposed to oxygen and water corrosion will begin.
Protecting a metal from corrosion by making it the cathode in an electrochemical cell is known as cathodic protection. The metal that is oxidized while protecting the cathode is called the sacrificial anode.
