Basic physics behind rusting of Iron. Does-metals also corrode? How to prevent rusting?

1. MODERN DELHI PUBLIC SCHOOL
SECTOR 87 FARIDABAD
RUSTING
OF
IRON
[CHEMISTRY PROJECT FILE]
BY :
NAME - UTKARSH VARSHNEY
CLASS - 12th
(Master Ability)
ROLL NO. -

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MODERN DELHI PUBLIC SCHOOL
SECTOR 87 – FARIDABAD
This is to certify that UTKARSH VARSHNEY a student of class
12th
has successfully completed his research on the topic
RUSTING OF IRON
under the guidance of Mr VIKRAM BHAMBHU during the
year (2017 – 18) in partial fulfilment of Chemistry practical
conducted by Central Board of Secondary Education (CBSE)
___________________ ___________________
Signature of Examiner Signature of Teacher

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ACKNOWLEDGEMENT
[It would be my utmost pleasure to express
my sincere thanks to my chemistry teacher
Mr VIKRAM BHAMBHU in providing me a
helping hand and his valuable guidance,
support and supervision all through this
project. I would also like to thanks my
parents and my brother (UJJWAL VARSHNEY
– IIT KANPUR) for helping me in the editing
part and presentation part.]

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INDEX
S.NO TOPIC
1 INTRODUCTION
2 GALVANIC CORROSION
3 MECHANISM OF RUSTING
4 RUSTING IN NON METALS
4.1 RUSTING IN GLASS
4.2 PREVENTIONS
5 EXPERIMENT (RUSTING OF IRON)
5.1 REQUIREMENT
5.2 PROCEDURE
5.3 OBSERVATION TABLE
5.4 CONCLUSION
6 FACTOR PROMOTING RUST
7 MEATHODS OF PREVENTION
8 BIBLOGRAPHY

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INTRODUCTION
Rust is an iron oxide, usually
red oxide formed by
the redox reaction
of iron and oxygen in the
presence of water or air
moisture. Several forms of
rust are distinguishable
both visually and
by spectroscopy,
and form under
different
circumstances. Rust
consists of
hydrated iron (III)
oxides Fe2O3·nH2O and iron (III) oxide hydroxide (FeO(OH), Fe(OH)3).
Given sufficient time, oxygen and water, any iron mass will eventually convert
entirely to rust and disintegrate. Surface rust is flaky and friable, and provides no
protection to the underlying iron, unlike the formation of patina on copper surfaces.
Rusting is the common term for corrosion of iron and its alloys, such as steel. Many
other metals undergo equivalent corrosion, but the resulting oxides are not
commonly called rust.
Other forms of rust exist, like the result of reactions between iron and chloride in an
environment deprived of oxygen – rebar used in underwater concrete pillars is an
example – which generates green rust.

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GALVANIC RUSTING
DEFINATION - Galvanic corrosion (also called bimetallic corrosion) is an
electrochemical process in which one metal corrodes preferentially when it is in
electrical contact with another, in the presence of an electrolyte.
EXAMPLE - A common example of galvanic corrosion occurs in corrugated iron, a
sheet of iron or steel covered with a zinc coating. Even when the
protective zinc coating is broken, the underlying steel is not attacked. Instead, the
zinc is corroded because it is less noble; only after it has been consumed can rusting
of the base metal occur in earnest. By contrast, with a traditional tin can, the
opposite of a protective effect occurs: because the tin is more noble than the
underlying steel, when the tin coating is broken, the steel beneath is immediately
attacked preferentially.
Statue of Liberty - A spectacular
example of galvanic corrosion
occurred in the Statue of
Liberty when regular maintenance
checks in the 1980s revealed that
corrosion had taken place between
the outer copper skin and
the wrought iron support structure.
Although the problem had been
anticipated when the structure was
built by Gustave Eiffel to Frédéric
Bartholdi's design in the 1880s, the
insulation layer of shellac between
the two metals had failed over time and resulted in rusting of the iron supports. An
extensive renovation requiring complete disassembly of the statue replaced the
original insulation with PTFE. The structure was far from unsafe owing to the large
number of unaffected connections, but it was regarded as a precautionary measure
to preserve a national symbol of the United States.

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MECHANISM OF RUSTING
OVERVIEW – The theory of rust can be explained by taking the example of
rusting of iron. The theory is called electrochemical theory because it explains
the formation of rust on the basis of formation of electrochemical cells on the
surface of the metal.
The overall rusting involves the following steps:
(i) Oxidation occurs at the anodes of each electrochemical cell. Therefore, at each
anode neutral iron atoms are oxidised to ferrous ions.
At anode:
Thus, the metal atoms in the lattice pass into the solution as ions, leaving electrons
on the metal itself. These electrons move towards the cathode region through the
metal.
(ii) At the cathodes of each cell, the electrons are taken up by hydrogen ions
(reduction takes place). The ions are obtained either from water or from acidic
substances (e.g. in water
Or
At cathode:
The hydrogen atoms on the iron surface reduce dissolved oxygen.
Therefore, the overall reaction at cathode of different electrochemical cells may be
written as,
(iii) The overall redox reaction may be written by multiplying reaction at anode by
2 and adding reaction at cathode to equalise number of electrons lost and gained -

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Oxidation half reaction:
Reduction half reaction:
Overall cell reaction:
The ferrous ions are oxidised further by atmospheric oxygen to form rust.
It may be noted that salt water accelerates corrosion. This is mainly due to the fact
that salt water increases the electrical conduction of electrolyte solution formed on
the metal surface. Therefore, rusting becomes more serious problem where salt
water is present.

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RUSTING IN NON METALS
INTRODUCTION – Corrosion is often thought of as the oxidation of metals such
as iron, but ceramics also corrode, or react with their environment. Concrete, for
example, generally is very stable, but it contains calcium hydroxide and calcium
aluminate, which are attacked by sulphates, such as calcium sulphate often
present in ground water. Tungsten carbide, usually highly resistant to corrosion,
is destroyed in less than a week of contact with sulphuric acid, H2SO4.
CERAMICS – Most of the ceramics material are almost immune to corrosion. The
strong ionic/covalent bonds that hold them together leave very little free chemical
energy in the structure. So, they can be thought of as already corroded.
An example of corrosion protection in ceramics is the lime added to soda-lime glass
to reduce its solubility in water.
POLYMERS - Corrosion on polymers, both plastics and rubber materials, is in many
cases similar to metals but in other cases it looks very different. Corrosion attacks on
polymers are often hard to discover, the material may look normal but can in fact be
embrittled and have lost its mechanical strength.
Mechanical stressed polymers applied in chemical environments may initiate cracks
on the surfaces. These cracks can thereafter propagate through the material either
as a result of the mechanical stresses or in combination with continuing chemical
attack. Corrosion of polymers can be divided into either chemical reaction or
physical interaction.
CHEMICAL REACTION - Polymers consist of a network with molecular chains mainly
consisting of carbon, hydrogen and oxygen. Corrosion by chemical reaction changes
the configuration of the polymer chains. Listed below are some of the environments
that cause chemical reactions in polymers.
1. Heat: Chain scission will occur when polymers are exposed to heat above a specified
temperature limit, which is unique for each type of polymer.

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2. UV- radiation: In the presence of oxygen, UV-radiation can cause a breakdown of the
polymer chains.
3. Ozone: Attacks from ozone on unsaturated polymers (e.g. natural rubber) under stress,
causes characteristic cracks.
4. Water: Absorption of water at elevated temperatures causes hydrolysis of certain groups in
a polymer chain (e.g. urethane and ester groups). Hydrolysis weakens the polymer since
the backbone structure is altered.
PHYSICAL INTERACTIONS - Physical effects on polymers are caused by interaction with
the environment. This may lead to swelling, dissolving or leakage of additives. The
interaction is dependent on diffusion of substances into the polymer, and the
process is in some cases reversible.
Organic substances usually affect polymers through physical interaction, while
substances like strong acids or bases normally result in an irreversible breakdown of
polymers.
RUSTED GLASS

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PREVENTION OF RUSRING IN GLASS - Laboratory scale procedures and practical tests
were used to study the problem of glass ware corrosion (Permanent Filling) in
domestic mechanical dishwashers. Result of these tests showed glassware
corrosion to be caused by alkali attack of the glass structure. It was also found
that sequestrates such a sodium triphosphates greatly accelerate the corrosive
action of alkali. Silicates, certain metals and metal oxides were found to inhibit
glassware corrosion. The basis for this inhibiting effect is believed to be
adsorption on the glass surface of reaction products of these materials in an
alkaline solution.
RUSTED GLASS

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EXPERIMENT – RUSTION OF IRON NAIL
OBJECTIVE: TO STUDY THE EFFECT OF METAL
COUPLING ON RUSTING OF IRON.
AIM – 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 electro-positive metal like zinc, magnesium or aluminium rusting is prevented
but if on the other hand, it is coupled with less electro –positive metals like copper,
the rusting is facilitated.
APPARATUS - COUPLING ON RUSTING OF IRON

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EQUIPMENTS REQUIRED
S.NO EQUIPMENTS
1 Two Petri dishes
2 Four test-tube
3 Four iron nails
4 Beaker
5 Sand paper
6 Wire gauge
7 Gelatine
8 Copper, Zinc and Magnesium strips
9 Potassium ferricyanide solutions
10 Phenolphthalein
Phenolphthalein Potassium Ferricyanide Solution
solutions

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PROCEDURE
1. At first we have to clean the surface of iron nails with the help of
sand paper.
2. 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
3. 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
4. Keep the covered Petri dishes for one day or so.
5. 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 colourless phenolphthalein to pink.
Zinc strip wrapped around one nail.

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OBSERVATION TABLE
S.NO METAL PAIR COLOUR OF
PATCH
NAIL RUSTS
OR NOT
1 IRON- ZINC PINK NO
2 IRON- MAGNESIUM PINK NO
3 IRON- COPPER BLUE YES
CONCLUSION
It is clear from the observation that coupling of iron
with more electropositive metals such as zinc and
magnesium resists corrosion and rusting of iron.
Coupling of iron with less electropositive metals such as
copper increases rusting.

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FACTORS PROMOTING RUSTING
Four elements need to be present for corrosion to occur and collectively
referred to as the corrosion cell: an anode (+), a cathode (-), a metallic conductor
and an electrolyte. Changing the potency of the electrolyte affects the rate of
corrosion. Corrosion rates are determined by a variety of factors; however, five
factors do play an overwhelmingly important role in determining corrosion
rates.
Oxygen: Like water, oxygen increases the rate of corrosion. Corrosion can take
place in an oxygen-deficient environment, but the rate of the corrosion reaction
(and destruction of the metal) is generally much slower. In immersed conditions,
if an electrolyte is in contact with one area of metal containing more oxygen than
the electrolyte in contact with another area of the metal, the higher oxygen-
concentration area is cathodic relative to the remaining surface. An oxygen
concentration cell then forms, which results in rapid corrosion.
Temperature: Corrosion reactions are electrochemical in nature and usually
accelerate d with increasing temperature; therefore, corrosion proceeds faster
in warmer environments than in cooler ones.
Chemical Salts: Chemical salts increase the rate of corrosion by increasing the
efficiency (conductivity) of the electrolyte. The most common chemical salt is
sodium chloride, a major element of seawater. Sodium chloride deposited on
atmospherically exposed surfaces also acts as a hygroscopic material (i.e., it
extracts moisture from the air), which then increases the corrosion in non-
immersed areas.
Humidity: Humidity and time-of-wetness play a large role in promoting and
accelerating corrosion rates. Time-of-wetness refers to the length of time an
atmospherically exposed substrate has sufficient moisture to support the
corrosion process. The wetter the environment, the more corrosion is likely to
occur.
Pollutants: Acid rain (a chemical by-product from manufacturing and
processing plants), and chlorides (in coastal areas) promote corrosion. Acid
gases, such as carbon dioxide, can also dissolve in a film of moisture in contact
with the metal.

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METHODS OF PREVENTION OF RUSTING
1. Barrier Coatings - One of the easiest and cheapest ways to prevent corrosion
is to use barrier coatings like paint, plastic, or powder. Powders, including
epoxy, nylon, and urethane, are heated to the metal surface to create a thin film.
Plastic and waxes are often sprayed onto metal surfaces. Paint acts as a coating
to protect the metal surface from the electrochemical charge that comes from
corrosive compounds. Today’s paint systems are actually a combination of
different paint layers that serve different functions. The primer coat acts as an
inhibitor, the intermediate coat adds to the paint’s overall thickness, and the
finish coat provides resistance to the environmental factors.
2. Hot-Dip Galvanization- This corrosion prevention method involves dipping
steel into molten zinc. The iron in the steel reacts with the zinc to create a
tightly-bonded alloy coating which serves as protection. The process has been
around for more than 250 years and has been used for corrosion protection of
things like artistic sculptures and playground equipment. Compared to other
corrosion prevention methods, galvanization is known for lower initial costs,
sustainability, and versatility.
3. Alloyed Steel (Stainless) - Alloyed steel is one of the most effective corrosion
prevention methods around, combining the properties of various metals to
provide added strength and resistance to the resulting product. Corrosion-
resistant nickel, for example, combined with oxidation-resistant chromium
results in an alloy that can be used in oxidized and reduced chemical
environments. Different alloys provide resistance to different conditions, giving
companies greater flexibility.
4. Cathodic Protection - Cathodic protection protects against galvanic
corrosion, which occurs when two different metals are put together and exposed
to a corrosive electrolyte. To prevent this, the active sites on the metal surface
need to be converted to passive sites by providing electrons from another
source, typically with galvanic anodes attached on or near the surface. Metals
used for anodes include aluminium, magnesium, or zinc.

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BIBLIOGRAPHY
1. CLASS 12TH CHEMISTRY
NCERT
2. www.wikipedia.org
3. www.icbse.com
4. www.technopedia.com
5. www.slideshare.net