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    Corrosion PDF

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    John

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    Corrosion PDF

    Uploaded by

    John
    173 views3 pages

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    Corrosion.pdf

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    173 views3 pages

    Corrosion PDF

    Uploaded by

    John

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    © All Rights Reserved
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    of 3

    Corrosion

    We all must have observed newly bought iron, silver or coper articles appears very shiny but
    with passage of time they get dull. This is due to the layer of metal oxide that develops on their
    surface. Rusting of iron, silver jewellery getting tarnished or copper articles getting covered by
    green layer. Metals react with atmospheric oxygen and produces metal oxides that are basic in
    nature because they react with water to form bases.

    o In case of rusting of iron, the iron reacts with the oxygen present in air and moisture
    and develops rust (hydrated iron (III) oxide).

    Fig. Iron chain left in moist air got rusted


    Corrosion causes damage to car bodies, bridges, iron railings, ships and to all objects made of
    metals, especially those of iron.

    o In case of rusting of copper, the metallic copper reacts with oxygen, carbon-dioxide and
    atmospheric moisture and develops a green coloured coating of copper hydroxide and
    copper carbonate.

    Fig. Copper developing green coloured rust on exposure to moist air


    o In case of tarnishing of silver articles, the metallic silver reacts with hydrogen sulphide
    or sulphur present in air and gets tarnished.
    Fig. Tarnished silver Vs polished silver

    In corrosion, a metal is oxidised by loss of electrons to oxygen and formation of oxides.


    Corrosion of iron (commonly known as rusting) occurs in presence of water and air. The
    chemistry of corrosion is quite complex but it may be considered essentially as an
    electrochemical phenomenon. At a particular spot (Fig. 3.13) of an object made of iron,
    oxidation takes place and that spot behaves as anode and we can write the reaction.

    Anode: 2 Fe (s) → 2 Fe2+ + 4 e– E°(Fe2+ /Fe) = – 0.44 V

    Electrons released at anodic spot move through the metal and go to another spot on the metal
    and reduce oxygen in the presence of H+ (which is believed to be available from H2CO3 formed
    due to dissolution of carbon dioxide from air into water. Hydrogen ion in water may also be
    available due to dissolution of other acidic oxides from the atmosphere). This spot behaves as
    cathode with the reaction:

    Cathode: O2(g) + 4 H+(aq) + 4 e– → 2 H2O (l) E°(H+/O2 / H2O) =1.23 V

    The overall reaction being:

    2Fe(s) + O2(g) + 4H+(aq) → 2Fe2 +(aq) + 2 H2O (l) E(cell) =1.67 V

    The ferrous ions are further oxidised by atmospheric oxygen to ferric ions which come out as
    rust in the form of hydrated ferric oxide (Fe2O3. x H2O) and with further production of hydrogen
    ions.

    Prevention of corrosion

    One of the simplest methods of preventing corrosion is to prevent the surface of the metallic
    object to come in contact with atmosphere. This can be done by covering the surface with paint
    or by some chemicals (e.g. bisphenol). Another simple method is to cover the surface by other
    metals (Sn, Zn, etc.) that are inert or react to save the object. An electrochemical method is to
    provide a sacrificial electrode of another metal (like Mg, Zn, etc.) which corrodes itself but saves
    the object.

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