The document discusses corrosion as the deterioration of materials due to environmental reactions, emphasizing its prevention through methods such as surface coatings and galvanization. It highlights the electrochemical nature of corrosion in aqueous environments, detailing anodic and cathodic processes and specific types of corrosion, including pitting and galvanic corrosion. Additionally, it describes methods for determining corrosion rates and offers insights into concentration cell corrosion and prevention strategies.

1. CORROSION
PREVENTION
AND CONTROL

2. Corrosion
– it is not simply “rust”
– means "the deterioration of a
material or its properties due to a reaction
of that material with its chemical
environment."

3. Deterioration of materials is cause by
-sun exposure
-oxygen and moisture
-Wind and other environmental elements.

4. Corrosion of metals can be prevented by:
 Surface coating
 Galvanization : Prevention of
corrosion of iron by Zn coating.
 Applying, oil, grease, paint or varnish
on the surface.
 By forming insoluble phosphate or
chromate coating

5. -Oil
-Painting
-Chromate Coating

6. 5.1 Significance and Purpose
5.1 Significance and Purpose
- The basic purpose of corrosion control is to maintain
the soundness and integrity of a structure.
-the main purposes of Corrosion Prevention and
Control (CPC) is to reduce cost of corrosion effects on
Facilities and Infrastructure (F&I) and maximize life-
cycle return on investment (ROI) increasing
sustainability and durability.

7. Electrochemical nature of aqueous corrosion
5.2Electrochemical nature of aqueous
corrosion
Aqueous Corrosion (corrosion in water media) is
electrochemical in nature. This means that the
corrosion reaction involves electrons (e-
).
It is divided into Anodic and Cathodic reactions.

8. An Anodic Reaction is an
oxidation process through which
the valence of a metal increases
from zero to a more positive
value.
Aqueous Corrosion involves the
loss of metal to the aqueous
phase. When this occurs, the
metal releases a number of
electrons into the Metal/Alloy.
M M
→ n+
+ ne-
In
this reaction, Mn+
is
a soluble corrosion
product that does
not protect the
surface of the metal.

9. A Cathodic Reaction is
a reduction process during
which the valence of the
species is reduced.
The electrons produced
during the anodic corrosion
reaction must be consumed at
the same rate by a
corresponding cathodic
reaction.

10. Corrosion can be defined as the deterioration
of materials by chemical processes. Of these,
the most important by far is electrochemical
corrosion of metals, in which the oxidation
process M M
→ +
+ e–
is facilitated by the
presence of a suitable electron acceptor,
sometimes referred to in corrosion science as
a depolarizer.

11. Electrochemical corrosion
of iron
Corrosion often begins at a location
(1) where the metal is under stress (at
a bend or weld) or is isolated from
the air (where two pieces of metal are
joined or under a loosely-adhering
paint film.) The metal ions dissolve in
the moisture film and the electrons
migrate to another location (2) where
they are taken up by adepolarizer.
Oxygen is the most common
depolarizer; the resulting hydroxide
ions react with the Fe2+
to form the
mixture of hydrous iron oxides
known as rust.

13. Pitting corrosion
Pitting corrosion
Most metals are covered
with a thin oxide film which
inhibits anodic dissolution.
When corrosion does occur,
it sometimes hollows out a
narrow hole or pit in the
metal. The bottoms of these
pits tend to be deprived of
oxygen, thus promoting
further growth of the pit
into the metal.

15. 5.3 Corrosion Rate Determination
5.3 Corrosion Rate Determination
I. Weight loss measurements
II. Electrochemical tests

16. I. Weight loss measurements
The simplest way of measuring the corrosion rate of
a metal is to expose the sample to the test medium
(e.g. sea water) and measure the loss of weight of
the material as a function of time.

17. II. Electrochemical tests
- Electrochemical tests methods can be used to
characterize corrosion mechanisms and predict
corrosion rates.

18. 5.4 Galvanic and concentration cell corrosion
5.4 Galvanic and concentration cell
corrosion
Galvanic 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.

19. Example:
A common example of galvanic corrosion is
the rusting of corrugated iron sheet, which
becomes widespread when the protective
zinc coating is broken and the underlying
steel is attacked.

20. Preventing Galvanic Corrosion
Electrically insulate the two metals from each other.
If they are not in electrical contact, no galvanic
coupling will occur.
Aluminum anodes mounted
on a steel-jacketed structure

21. Ensure there is no contact with an
electrolyte. This can be done by
using water-repellent compounds
such as greases, or by coating the
metals with an impermeable
protective layer, such as a suitable
paint, varnish, or plastic.

22. Concentration cell corrosion is a limited form
of a galvanic cell that has two equivalent
half-cells of the same composition differing
only in concentrations. It produces small
voltage as it attempts to reach
chemical equilibrium, which occurs when
the concentration of reactant in both half-
cells are equal.

23. Three general
types of
concentration
cells

24. Metal ion concentration cells
In the presence of water, a high concentration of
metal ions will exist under faying surfaces and a low
concentration of metal ions will exist adjacent to the
crevice created by the faying surfaces. An
electrical potential will exist between the two points.
The area of the metal in contact with the high
concentration of metal ions will be cathodic and will
be protected, and the area of metal in contact with
the low metal ion concentration will be anodic and
corroded.

25. Oxygen concentration cells
Water in contact with the metal surface will
normally contain dissolved oxygen. An oxygen cell
can develop at any point where the oxygen in the
air is not allowed to diffuse uniformly into the
solution, thereby creating a difference in oxygen
concentration between two points. Corrosion will
occur at the area of low-oxygen concentration,
which are anodic.

26. Active-passive cells
If a metal is protected against corrosion by a tightly
adhering passive film (usually an oxide) and salt
deposits on the surface in the presence of water, the
active metal beneath the film will be exposed to
corrosive attack in areas where the passive film is
broken. An electrical potential will develop between
the large area of the passive film (cathode) and the
small area of the exposed active metal (anode).
Rapid pitting of the active metal will result.

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