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Rusting of Iron
2022-08-22 23:45:45

Rusting of Iron

What is a metal?

Metals and non-metals are the two main types used to categorise the materials surrounding us. Materials with the properties of being strong, flexible, shiny, fusible, and ductile are said to be metals. Metals include things like iron, steel, aluminium, gold, copper, etc.

Rusting of Iron

1. What is Rusting of Iron?

Let's picture our favourite colour being used to paint our bike. In addition to making the bike seem nice, the paint also prevents rusting, which is a very vital function.

Rusting happens whenever iron and steel are subjected to moisture or water and air or oxygen. Iron and steel corrosion is referred to as rusting. Corrosion is a phrase used to explain how metal starts to break down when it is revealed to the environment. When we say a metal "corrodes," we mean it has reacted with environmental water and oxygen to generate the corresponding oxides.

2. Why do metals rust?

  • A redux reaction is technically two sequential reactions in one: reduction and oxidation, which causes corrosion. When some molecules gain electrons, reduction takes place. The oxidation procedure, on the other hand, happens whenever some molecules or atoms shed their electrons.
  • Rusting, which occurs in iron, is a typical illustration of corrosion.
  • Three things are necessary:
    • Anode: An iron in this instance that readily releases electrons.
    • The oxygen in this scenario serves as the cathode, which readily receives electrons.
    • Finally, an electrolyte solution, which in this case is water, transports ions between thecathode and the anode.
  • The electrolyte solution, which is water, aids in the formation of rust. Rust is the result of iron converting to hydrated ferric oxide due to this redox process.
  • Iron and other corrosive metals are powerless to stop it. Some metals have a thermodynamically preferred state, which means they would much rather always be rusted. Simply said, we humans alter it to suit our own wants.

3. Experiment to demonstrate rusting

  • A straightforward experiment will help us to visualise this.
  • In four test tubes, put an iron nail (say Tb1, Tb2, Tb3, and Tb4).
  • Tb1 should be completely submerged with ordinary tap water.
  • Regular tap water should be added to the Tb2, along with half a spoon of sodium chloride. Shake the Tb2 side to side until the common salt completely dissolves.
  • Fill the Tb3 with just-boiling water and add a thin coating of oil.
  •  Any oxygen cannot dissolve further due to this oily coating. Add a small amount of calcium chloride, which draws out moisture or water, to the Tb4.

The nail is exposed to water and air or oxygen in test tube one (Tb1)

The nail in test tube two (TB2) is exposed to air, moisture, and salt.

The nail is in contact with water in the third test tube, Tb3.

The nail in test tube number four (TB4) is exclusively exposed to oxygen.

Let's use a rubber bong to stop all four test tubes, then let it stand for a few days.

Observations

  • Both the first and second test tubes' nails are corroded. In particular, the second test tube's nail has rusted more than the first tube's nail. At the bottom of these tubes, we can see a reddish-brown precipitate where this rust had partially flaked off.
  • In the third and fourth test tubes, the nails are undamaged.

Let's examine the above findings:

  • The iron has undergone an oxidation reaction forming hydrated iron (III) oxide or rust.
  • The catalysts, salt and acid act to speed up the process, which is why the nail in the second test tube developed more rust than the nail in the first.

4. Chemical process of rusting

  • Stage 1 of rusting: Iron is oxidised into iron (II) ion
    The iron atoms in the water droplet's centre will now begin to lose two electrons, forming iron (II) ions. From the water droplet's centre, electrons go outward and surround it. Iron is considered to have been oxidised to iron (II)ions because it lost its electrons.
    Equation: Fe→Fe2++ 2e-
  • Stage2 of rusting: Oxygen is reduced to hydroxide ions
    In this stage, oxygen accepts the electrons in the presence of water to form hydroxide ions.
    (Since the electrons from the middle move towards the edge, there exists an abundance of oxygen, and this oxygen accepts the electron at the edge of the water droplet). This process produces hydroxide ions.
    Equation: O2+2H2O+4e-→4OH-
  • Stage 3 of rusting: Iron (II) ions react with hydroxide ions
    Both the iron (II) ions and hydroxide ions react together to form iron (II) hydroxide, which is a solid.
    Equation: Fe2++OH-→Fe (OH)2
  • Stage 4 of rusting: Iron (II) hydroxide is oxidised
    Iron (II) hydroxide is oxidised by oxygen into iron (III) hydroxide. Whenever there is an oxidation process in rusting, oxygen is always accompanied by water.
    Equation: 4Fe (OH)2+O2+2H2O→4Fe (OH)3
  • Stage 5 of rusting: Iron (III) hydroxide decomposes
    We know that rust is hydrated iron (III) oxide. In this stage, iron (III) hydroxide decomposes into iron (III) oxide.
    Equation: Fe (OH)3→Fe2O3. xH2O

5. Parameters that must prevail for iron to rust

Rust can only occur under two circumstances.

  • Water or moisture
  • Oxygen or air

6. Damages caused by rusting

  • The quality of metals is compromised.
  • Metal becomes less effective.
  • Maintenance requirements and material costs rise.

7. Prevention and Treatment

  • If we have to regularly replace goods made of steel and iron; rusting can be very pricey.
  • If we coat the regions that are exposed to oxygen and moisture, rusting can be avoided. For instance, it is done on a lot of bicycles and autos.
  • Placing a layer of oil or grease is another way to stop rust. For instance, this is done on numerous chains for bicycles. This lowers friction and delays the corrosion process by lubricating the moving parts.
  • Galvanizing is a protective technique that offers an additional means of preventing corrosion. When exposed to oxygen and moisture, zinc reactivates more quickly than iron, according to a comparison. Zinc will therefore corrode more quickly than iron. By using this characteristic, zinc can be utilised to shield iron by being placed on top of it. So, sacrificial protection is another name for this technique.
    Ships' corrosion is stopped using this technique.

8. Sample Questions

1.What can be done to stop iron from rusting?

  • By painting the regions that are in contact with oxygen, air, and water.
  • Pacing an oil or grease layer
  • Galvanizing is a cost-effective way of protection.

2. Describe rusting.

Steel or iron that has been exposed to oxygen and moisture will corrode by rusting.

3. What types of harm does rusting cause?

  • The purity of metals is impacted.
  • Metal becomes less effective.
  • Maintenance requirements and material costs rise.

4. What is the redox reaction of rusting of iron?

  • Iron is oxidised into iron (II) ion: Fe→Fe2++ 2e-
  • Oxygen is reduced to hydroxide ions: O2+2H2O+4e-→4OH-
  • Iron (II) ions react with hydroxide ions: Fe2++OH-→Fe (OH)2
  • Iron (II) hydroxide is oxidised: 4Fe (OH)2+O2+2H2O→4Fe (OH)3
  • Iron (III) hydroxide decomposes: Fe (OH)3→Fe2O3. xH2O

5. What are the favourable conditions for rusting?

The favourable conditions for rusting are:

  • Water
  • Oxygen
  • Impurity

6. Describe metals.

Materials with the properties of being hard, malleable, lustrous, fusible, and ductile are said to be metals.

7. What are some examples of metals?

Some examples of metals are iron, aluminium, gold, copper etc.

8. What is the process of rusting?

Initially, the iron atoms in the centre of the water droplet will start to give away two electrons to form iron (II) ions. Electrons travel from the centre of the water droplet towards the edge and all around. The oxygen accepts the electrons in the presence of water to form hydroxide ions. Both the iron (II) ions and hydroxide ions react together to from iron (II) hydroxide. Iron (II) hydroxide is oxidised by oxygen into iron (III) hydroxide. Finally, iron (III) hydroxide decomposes into iron (III) oxide.

9. What is corrosion?

The term corrosion is used to describe the breakdown of the metal when exposed to the environment. When we say a metal corrodes, we mean that it has reacted with water and oxygen from the environment to form their respective oxides.