The document discusses the corrosion and rusting of metals like iron. It explains that rusting is an electrochemical process where iron loses electrons and oxygen gains electrons, forming hydroxide ions. The rate of corrosion is affected by water and electrolytes. Different metals coupled to an iron nail will affect its rusting - a more electropositive metal like zinc will prevent rusting, while a less electropositive metal like magnesium will facilitate rusting. The document also outlines the reactions involved in rust formation and discusses how rust changes over time based on conditions.

2. AKNOWLEDGEMENT
I undertook this Project work , as the part of my XII-Chemistry project .I had tried
to apply my best of knowledge and experience gained during study and class work
experience.
I would like to extend my sincere thanks and gratitude to my teacher Mr.Sakthivel
Murugan
I would like to take the opportunity to extend my sincere thanks and gratitude to our
parents for being a source of inspiration and providing time and freedom to develop
this project.

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5. INTRODUCTION
Metal corrosion is the most common form of corrosion. The
corrosion occurs at the surface of the metal in forms of
chemical or electrochemical reactions. This process significantly
reduces the strength, plasticity, toughness and other
mechanical properties of the metallic material. However,
because of the metal and its alloys is still the most important
pipe and structure materials, the cost of corrosion grows
significantly with the growth of industries. Thus many scientists
focus on the research of corrosion control in order to reduce the
cost of replacing the rusting metal material.

7. MECHANISM OF CORROSION OF METAL
General Principle of Corrosion: Reaction is the fundamental
reaction during the corrosion process, which the electron can
flow from certain areas on the metal surface to other areas
through a solution which can conduct electric currents. Basically,
both anodic and cathode reactions have to balance each other
out, resulting in a neutral reaction. Both anodic and cathodic
reactions occur simultaneously at the same rates. What’s more,
the site of these electrodes may consist of either two different
kinds of metals, or they may be on different areas of the same
piece of metal, resulting a potential difference between the two
electrodes, so that the oxidation reaction of the metal at the
anode and formation of negative ions at the cathode can take
place at the same time.

8. Similar electrical potentials may also be developed between two
areas of a component made of a single metal as result of small
differences in composition or structure or of differences in the
conditions to which the metal surface is exposed. That part of a
metal which becomes the corroding area is called the “anode” ; that
which acts as the other electrode of the battery is called “cathode”
which does not corrode, but is an important part of the system. In
the corrosion systems commonly involved, with water containing
some salts in solution as the electrolyte. Corrosion may even take
place with pure water, provided that oxygen is present. In such cases
oxygen combines with the hydrogen generated at the cathode,
removing it and permitting the reaction to go on.
RUSTING: AN ELECTROCHEMICAL

9. MECHANISM
Rusting may be explained by an electrochemical mechanism. In
the presence of moist air containing dissolved oxygen or carbon
dioxide, the commercial iron behave as if composed of small
electrical cells. At anode of cell, iron passes into solution as
ferrous ions. The electron moves towards the cathode and form
hydroxyl ions. Under the influence of dissolved oxygen the
ferrous ions and hydroxyl ions into form, i.e., hydrated ferric
oxide.
METHODS OF PREVENTION OF CORROSION
AND RUSTING
Some of the methods used:-

10. ➜Barrier Protection: In this method, a barrier film is introduced
between iron surface and atmospheric air. The film is obtained by
painting, varnishing etc.
➜Galvanization: The metallic iron is covered by a layer of more
reactive metal such as zinc. The active metal losses electrons in
preference of iron. Thus, protecting from rusting and corrosion.
AIM OF THIS PROJECT
In this project the aim is to investigate effect of the metals coupling
on the rusting of iron. Metal coupling affects the rusting of iron. If
the nail is coupled with a more electropositive metal rusting is
prevented but if on the other hand, it is coupled with less electro –
positive metals the rusting is facilitated.

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12. At first we have to clean
the surface of iron nails
with the help of sand
paper.
After that we have to
wind zinc strip around
one nail, a clean copper
wire around the second
and clean magnesium
strip around the third
nail.
Then to put all these
three and a fourth nail
in Petri dishes so that
they are not in contact
with each other.
Then to fill the Petri
dishes with hot agar-
agar solution in such a
way that only lower half
of the nails are covered
with the liquids.
Keep the covered Petri
dishes for one day or
so.
The liquids set to a gel
on cooling.
Two types of patches
are observed around
the rusted nail, one is
blue and the other pink.
Blue patch is due to the
formation of potassium
Ferro-ferricyanide
where pink patch is due
to the formation of
hydroxyl ions which
turns colorless
phenolphthalein to pink.

13. ASSOCIATED REACTIONS
The rusting of iron is an electrochemical process that begins with the
transfer of electrons from iron to oxygen. The iron is the reducing
agent(gives up electrons) while the oxygen is the oxidizing agent
(gains electrons). The rate of corrosion is affected by water and
accelerated by electrolytes, as illustrated by the effects of road salt
on the corrosion of automobiles.
The key reaction is the reduction of oxygen
𝑶𝟐 + 𝟒𝒆− + 𝟐𝑯𝟐𝑶⟶𝟒𝑶𝑯−
Because it forms hydroxide ions, this process is strongly affected by
the presence of acid. Indeed, the corrosion of most metals by
oxygen is accelerated at low pH.
Providing the electrons for the above reaction is the
oxidation of iron that may be described as follows:

14. 𝑭𝒆⟶𝑭𝒆𝟐+ + 𝟐𝒆−
The following redox reaction also occurs in the presence of water
and is crucial to the formation of rust: 𝟒𝑭𝒆𝟐+ + 𝑶𝟐⟶𝟒𝑭𝒆𝟑+ +
𝟐𝑶𝟐−
In addition, the following multistep acid-base reactions affect
the course of rust formation:
𝑭𝒆𝟐+ + 𝟐𝑯𝟐𝑶⟶𝑭𝒆(𝑶𝑯)𝟐+ 𝟐𝑯+ 𝑭𝒆𝟑+ +
𝟑𝑯𝟐𝑶⟶𝑭𝒆(𝑶𝑯)𝟑 + 𝟑𝑯+ as do the following
dehydration equilibria:
𝑭𝒆(𝑶𝑯)𝟐 ⟶𝑭𝒆𝑶 + 𝑯𝟐𝑶
𝑭𝒆(𝑶𝑯)𝟑 ⟶𝑭𝒆𝑶(𝑶𝑯) + 𝑯𝟐𝑶
𝟐𝑭𝒆𝑶(𝑶𝑯) ⟶𝑭𝒆𝟐𝑶𝟑 + 𝑯𝟐𝑶
From the above equations, it is also seen that the corrosion products are
dictated by the availability of water and oxygen. With limited dissolved

15. oxygen, iron(II) containing materials are favored, including FeO and black
lodestone or magnetite (𝑭𝒆𝟑𝑶𝟒). High oxygen concentrations favor ferric
materials with the nominal formulae 𝑭𝒆(𝑶𝑯)𝟑−𝒙𝑶𝒙/𝟐.
The nature of rust changes with time, reflecting the slow rates of the
reactions of solids.
Furthermore, these complex processes are affected by the presence of
other ions, such as 𝑪𝒂𝟐+, both of which serve as an electrolyte, and
thus accelerate rust formation, or combine with the hydroxides and
oxides of iron to precipitate a variety of Ca-Fe-O-OH species. Onset of
rusting can also be detected in laboratory with the use of ferroxyl
indicator solution. The solution detects both 𝑭𝒆𝟐+ ions and hydroxyl
ions. Formation of Fe2+ ions and hydroxyl ions are indicated by blue and
pink patches respectively.

16. RUSTING ALL AROUND.

17. COLOUR OF
THE PATCH
NAILS RUST
IRON-ZINC
IRON-MAGNESIUM

18. IRON-COPPER
IRON-NAIL

20. Thanks!