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                       CORROSION

• Corrosion is a galvanic process by which metals deteriorate through oxidation usually but not always to their oxides.

• For example when exposed to air, iron (Fe)is rusts, Silver (Ag) tarnishes , and Copper (Cu) and brass acquire a bluish green surface called a patina.

• Of the various metals subject to corrosion, Iron is by fast the most important commercially. 

 

                  

RUSTING OF IRON TARNISHING IN COPPER               TARNISHING IN SILVER

Corrosion is a redox process. The oxidation of most of metals is thermodynamically spontaneous with the notable exception of gold and platinum.

  RUSTING OF IRON

Rust is a combination of several different oxides of iron. 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.

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DIAGRAME: -PROCESS OF RUSTING

Prevention from Rusting

Methods used to prevent Rusting of Iron are as follows:

• Alloying – Since Rusting of Iron is a chemical process that happens because the metal is attaining more stable chemical state, alloying (mixing) the iron with other stable metals or alloys can slow down the process of rusting.

• Galvanizing – Galvanizing a metal object means to coat the surface of that object with a layer of metallic zinc. Also, it is an inexpensive procedure. In conclusion, it will provide it with protection against rusting.

• Coating and Painting – Coating the surface of a metal object with a layer of either Paint or Varnish will break the contact between the surface and atmospheric oxygen making it consequently immune to rusting.

• Humidity Control – Controlling the humidity of the environment is also a solution. Therefore, the chances of the metal object rusting will reduce.

 

Harmful Effects of Corrosion

• Lose of efficiency.

• Contamination of product.

• Damage of metallic equipment’s.

• Inability to use metallic materials.

• Lose of valuable materials such as blockage of pipes, mechanical damage of underground water pipes.

  CATALYST

Catalysis is the increase in the rate of a chemical reaction due to the participation of an additional substance called a catalyst. With a catalyst, reactions occur faster and require less activation energy. Because catalysts are not consumed in the catalyzed reaction, they can continue to catalyze the reaction of further quantities of reactant. Often only tiny amounts are required.

 

Types of catalysts

 The catalysts have been divided into different types according to their behavior and pattern of action.

 1.  Positive catalyst

 A catalyst which enhances the speed of the reaction is called positive catalyst and the phenomenon is known as positive catalysis. Various examples are given below:

i. Decomposition of H2O2 in presence of colloidal platinum

2H2O2 --- Pt --- > 2H2O   +    O2

ii. Decomposition of KClO3 in presence of manganese dioxide.

2KClO3  ---- MnO2 -- > 2KCl   +   3O2

2.  Negative Catalyst

 There are certain substances which, when added to the reaction mixture, retard the reaction rate instead of increasing it. These are called negative catalysts or inhibitors and the phenomenon is known as negative catalysis. The examples are given below.

 

i. The oxidation of sodium sulphite by air is retarded by alcohol.

            2 Na2SO3      +       O2---- Alcohol  -- >   2 Na2SO4

ii. The decomposition of hydrogen peroxide decreases in presence of glycerin.

2 H2O2  --- Glycerine --- > 2 H2O  +   O2

3. Auto catalyst

 In certain reactions, it is observed that one of the products formed during the reaction acts as a catalyst for that reaction. Such type of catalyst is called auto catalyst and the phenomenon is known as auto catalysis.

 In the oxidation of oxalic acid by potassium permanganate, one of the products MnSO4 acts as a auto-catalyst because it increases the speed of the reaction.

4. Induced Catalyst

When one reactant influences the rate of other reaction, which does not occur under ordinary conditions, the phenomenon is known as induced catalysis.

 Sodium arsenite solution is not oxidized by air. If, however, air is passed through a mixture of the solution of sodium arsenite and sodium sulphite, both of them undergo simultaneous oxidation. Thus sulphite has induced the arsenite and hence is called induced catalyst.

Promoters: The activity of a catalyst can be increased by addition of a small quantity of a second material. A substance which, though itself not a catalyst, promotes the activity of a catalyst is called a promoter. Some examples of the promoters are given below.

 i. In the Haber's process for the synthesis of ammonia, traces of molybdenum increase the activity of finely divided iron which acts as a catalyst.

N2     +   3 H2  < - Fe---  -- +Mo- - > 2NM3

Catalytic Poisons

A substance which destroys the activity of the catalyst is called a poison and the process is called catalytic poisoning. Some of the examples are-

 (i) The platinum catalyst used in the oxidation of SO2 in contact process is poisoned by arsenious oxide.

SO2     +   O2 < - Pt ---  -- poisoned by As2O3- - >  2 SO3

(ii) The iron catalyst used in the synthesis of ammonia in Haber process is poisoned by H2S

N2     +   3 H2  < - Fe ---  -- poisoned by H2S- - >  2 NH3

Active Centres

The catalytic surface has unbalanced chemical bonds on it. The reactant gaseous molecules are adsorbed on the surface by these free bonds. This accelerates the rate of the reaction. The distribution of free bonds on the catalytic surface is not uniform. These are crowded at the peaks, cracks and corners of the catalyst. The catalytic activity due to adsorption of reacting molecules is maximum at these spots. These are, therefore, referred to as the active centers. If a catalyst has more active centers, then its catalytic activity is increased.