CORROSION

Many of the metals we use in our daily are malleable, strong, and ductile. So when we use these metals, they are exposed to the environment. Over time, they lose their shine and become covered with a colored layer. Some of these metals even lose their strength and become brittle or weak. This happens due to the chemical interaction between metals and their environment. Thus, the spontaneous destruction of metals resulting from the interaction with the environment is called corrosion. It usually occurs on the exposed surface. If the top layer of the metal corrodes, the inner surface of the metal is also exposed. Therefore, it tends to continue to the depth of the metal.
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The 3 most common types of corrosion.They are
•general attack
•localized corrosion
•galvanic corrosion.
EXAMPLES OF CORROSION
Some of the most common examples of corrosion are:
•Tarnishing of silver Iron is coated with a brittle brown colored layer
•Brass and copper are coated with a green colored deposit
•The surface of the Aluminum Becomes Dull And Loses Its Luster.

CORROSION PREVENTION METHODS
There are some ways to guard metals from corrosion. Here we are getting to discuss a number of the foremost commonly used methods.

•BY SURFACE COATING
You can prevent corrosion of metal by coating its surface with one of the following products: By applying oil, grease, paint or varnish to the surface. By coating / depositing a thin layer of any other metal that does not corrode. For example, you can prevent the surface of iron from corroding by putting a thin layer of chromium, zinc or nickel on it. Likewise, brass or copper can be protected by covering it with a thin layer of tin. In our country, tinning brass is a very common practice.

•CONNECTION OF METAL TO A MORE ELECTROPOSITIVE METAL
You can also avoid corrosion by connecting the metal to a more electropositive metal. As long as the most electropositive metal is present, the donated metal will not be corroded. For example, you can protect iron from corrosion by connecting it to a zinc or magnesium block / plate.

•TREATMENT METAL SURFACES WITH PHOSPHORIC ACID.
It forms an insoluble phosphate coating on the surface which protects the metal from corrosion. You can also form a thin layer of chromate to prevent metals from corroding.
MECHANISM OF CORROSION
Several theories have been put forward to explain it over time. However, the modern view known as electrochemical theory seems more practical. Let us understand this theory with the help of an example of iron rust. Water containing α-oxygen and carbon dioxide acts as an electrolyte and helps the flow of electrons. Rust formation on the iron surface occurs by the following steps;
(i) At the anode region
The iron in contact with water forms an anode and oxidizes to Fe2 +. Anodic region: Fe (s) → Fe2 + (aq) + 2e– E0 Fe2 + / Fe = –0.44 V
The liberated electrons move to another part of the iron sheet. This part of the iron sheet serves as the cathode.
(ii) At the cathode region
In this cathode part of the surface, oxygen in the presence of H + ions (produced due to the ionization of water molecules) is reduced to form H2O.
Cathodic region:
O2 (g) + 4H + (aq) + 4e– → 2H2O (l)
EoO2 / H +H2O = +1.23 V
The overall reaction of the local cell is that the sum of the cathodic and anodic reactions.
Global cellular reaction:
2Fe (s) + O2 (g) + 4H + (aq) → 2Fe2 + (aq) + 2H2O (l) E0
cell = 1.67 V Fe2 + ions move in water on the surface of the leaf of iron. The presence of electrolytes (eg common salt , SO2, CO2, etc.) within the water helps carry more current through the local cell to the iron surface. The increased current flow increases the corrosion rate. Fe2 + ions are then oxidized by atmospheric oxygen to Fe3 + and form hydrated iron (III) oxide, expressed as Fe2O3.xH2O. Hydrated ferric oxide is called rust.
4Fe2 + (aq) + O2 (g) + 4H2O (l) → 2Fe2O3 (s) + 8H + Fe2O3 (s) + xH2O (l) → Fe2O3.xH2O
The H + ions produced in the above reaction further contribute to the