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Rusting

 
1.6 Rusting
 
Definition of Rusting of Iron
A metal corrosion that occurs to iron.
 
Process of Rusting of Iron
  • When iron metalrusts, a layer of reddish-brown iron oxide isformed on the iron surface that is easily cracked and permeable.
  • Thus,rusting occurs continuously and damage the structure of the iron.
 
Conditions for Rusting of An Iron
  • Oxygen gas.
  • Water.
 
This image is a flowchart that explains the redox reaction of rusting. It consists of four steps: 1. Oxidation 2. Reduction 3. Formation of a chemical cell 4. Formation of rust Each step is represented by a rectangular box with a number at the top. The boxes are connected by arrows indicating the sequence of the steps. The title of the flowchart is ‘Redox Reaction of Rusting,’ and the logo ‘Pandai’ is present on the right side. The overall design uses blue and red colors.
 
Redox Reaction of Rusting
Step Explaination

Oxidation
  • Iron atom releases two electrons to form iron(II) ion, \(Fe^{2+}\). 
    • \(Fe(s)\rightarrow Fe^{2+}(aq)+2e^-\\\)
  • The released electrons would flow through the iron to the surface. 
  • Iron is an electrical conductor, thus the electrons are able to flow through the iron. 
  • There is oxygen gas at the surface. 

Reduction
  • The released electrons from the oxidation process are received by oxygen gas. 
  • Since oxygen gas received electrons, oxygen is reduced to form hydroxide ion, \(OH^-\). 
    •  \(O_2(g)+2H_2O(l) +4e^-\rightarrow 4OH^-(aq)\)

Formation of a chemical cell
  • The part of the system that undergoes oxidation is the negative terminal of the chemical cell. 
  • The part of the system that undergoes reduction is the positive terminal of the chemical cell. 
    • \(2\times [Fe(s)\rightarrow Fe^{2+}(aq)+2e^-]\\ O_2(g)+2H_2O(l) +4e^-\rightarrow 4OH^-(aq)\\ \text{Equals to}\\2Fe(s)+O_2(g)+2H_2O(l)\rightarrow 2Fe^{2+}(aq)+4OH^-(aq) \)

Formation of rust
  • The iron(II) ions, \(Fe^{2+}\) flow from the negative terminal to the positive terminal. 
  • At the positive terminal, the iron(II) ions, \(Fe^{2+}\) reacts with the hydroxide ion, \(OH^-\) to form iron(II) hydroxide, \(Fe(OH)_2\). 
    • \(Fe(aq)+2OH^-(aq)\rightarrow Fe(OH)_2(s)\\\)
  • The iron(II) hydroxide is oxidised to form iron(III) hydroxide, \(Fe(OH)_3\). 
    • \(4Fe(OH)_2(s)+O_2(g)+2H_2O(l)\rightarrow 4Fe(OH)_3(s)\\\)
  • The iron(III) hydroxide, \(Fe(OH)_3\). then forms rust. 
  • Rust is hydrated iron(III) hydroxide, \(Fe_2O_3.3H_2O\). 
    • \( 4Fe(OH)_3(s)\rightarrow Fe_2O_3.3H_2O(s)\\ \phantom{ 4Fe(OH)_3(s)\rightarrow}Rust\)
 
Formation of Rust
This image illustrates the process of rust formation on iron. It shows a piece of iron submerged in water, with rust forming on the surface. The iron piece is connected to a positive terminal on both sides, while the center is connected to a negative terminal. Oxygen (O₂) is also present above the water. Rust is labeled at the points where it forms on the iron.
 
Definisi of Metal Corrosion
  • Oxidation of metal through the action of air or oxygen gas, water and/or electrolyte.
  • The oxidised metal would release electron(s) to form an ion.
 
Metal Corrosion
  • In general, the more electropositive the metal is, the easier it is for the metal to corrode.
  • For example, corrosion of iron, Fe is faster than copper, Cu.
 
Prevention of Metal Corrosion
Using A More Electropositive Metal
  • Metal corrosion can be prevented by touching a more electropositive metal with the metal being protected.
  • The more electropositive metal releases electrons easier.
  • For example; a layer of zinc protecting an iron, but there is a scratch on the layer of zinc.
    • Zinc is more electropositive than iron.
    • Zinc is oxidised by releasing electrons.
    • Electrons flow to the surface of iron where there are water and oxygen.
  • Instead of iron being oxidised, the zinc layer acts as a sacrificial layer of protection as shown below:
Zinc Layer as A Sacrificial Layer of Protection
This image illustrates the process of galvanic corrosion protection. It shows a piece of iron with a zinc coating. A water drop is present on the surface, causing a reaction. The zinc (Zn) reacts to form Zn²⁺ ions and releases electrons (e⁻). These electrons are then used in the reduction of oxygen (O₂) and hydrogen ions (H⁺) to form water (H₂O). The labels highlight the zinc coating, water drop, and iron substrate.

 
Using A Less Electropositive Metal
  • A layer of less electropositive metal (e.g. tin) protecting iron can still prevent corrosion of iron.
  • However, if the protective layer is scratched, the corrosion of iron is enhanced instead of prevented.
  • The more electropositive metal releases electrons easier.
  • For example; a layer of tin protecting an iron, but there is a scratch on the layer of tin.
    • Iron is more electropositive than tin.
    • Iron is oxidised by releasing electrons.
    • Electrons flow to the surface of iron where there are water and oxygen.
  • Instead of tin being oxidised, the iron is oxidised at a faster rate because iron is more electropositive than tin.
 

 

 

 

 

 

 

 

The views expressed in the notes here are those of the contributor and don’t necessarily represent the views of Pandai.

Rusting

 
1.6 Rusting
 
Definition of Rusting of Iron
A metal corrosion that occurs to iron.
 
Process of Rusting of Iron
  • When iron metalrusts, a layer of reddish-brown iron oxide isformed on the iron surface that is easily cracked and permeable.
  • Thus,rusting occurs continuously and damage the structure of the iron.
 
Conditions for Rusting of An Iron
  • Oxygen gas.
  • Water.
 
This image is a flowchart that explains the redox reaction of rusting. It consists of four steps: 1. Oxidation 2. Reduction 3. Formation of a chemical cell 4. Formation of rust Each step is represented by a rectangular box with a number at the top. The boxes are connected by arrows indicating the sequence of the steps. The title of the flowchart is ‘Redox Reaction of Rusting,’ and the logo ‘Pandai’ is present on the right side. The overall design uses blue and red colors.
 
Redox Reaction of Rusting
Step Explaination

Oxidation
  • Iron atom releases two electrons to form iron(II) ion, \(Fe^{2+}\). 
    • \(Fe(s)\rightarrow Fe^{2+}(aq)+2e^-\\\)
  • The released electrons would flow through the iron to the surface. 
  • Iron is an electrical conductor, thus the electrons are able to flow through the iron. 
  • There is oxygen gas at the surface. 

Reduction
  • The released electrons from the oxidation process are received by oxygen gas. 
  • Since oxygen gas received electrons, oxygen is reduced to form hydroxide ion, \(OH^-\). 
    •  \(O_2(g)+2H_2O(l) +4e^-\rightarrow 4OH^-(aq)\)

Formation of a chemical cell
  • The part of the system that undergoes oxidation is the negative terminal of the chemical cell. 
  • The part of the system that undergoes reduction is the positive terminal of the chemical cell. 
    • \(2\times [Fe(s)\rightarrow Fe^{2+}(aq)+2e^-]\\ O_2(g)+2H_2O(l) +4e^-\rightarrow 4OH^-(aq)\\ \text{Equals to}\\2Fe(s)+O_2(g)+2H_2O(l)\rightarrow 2Fe^{2+}(aq)+4OH^-(aq) \)

Formation of rust
  • The iron(II) ions, \(Fe^{2+}\) flow from the negative terminal to the positive terminal. 
  • At the positive terminal, the iron(II) ions, \(Fe^{2+}\) reacts with the hydroxide ion, \(OH^-\) to form iron(II) hydroxide, \(Fe(OH)_2\). 
    • \(Fe(aq)+2OH^-(aq)\rightarrow Fe(OH)_2(s)\\\)
  • The iron(II) hydroxide is oxidised to form iron(III) hydroxide, \(Fe(OH)_3\). 
    • \(4Fe(OH)_2(s)+O_2(g)+2H_2O(l)\rightarrow 4Fe(OH)_3(s)\\\)
  • The iron(III) hydroxide, \(Fe(OH)_3\). then forms rust. 
  • Rust is hydrated iron(III) hydroxide, \(Fe_2O_3.3H_2O\). 
    • \( 4Fe(OH)_3(s)\rightarrow Fe_2O_3.3H_2O(s)\\ \phantom{ 4Fe(OH)_3(s)\rightarrow}Rust\)
 
Formation of Rust
This image illustrates the process of rust formation on iron. It shows a piece of iron submerged in water, with rust forming on the surface. The iron piece is connected to a positive terminal on both sides, while the center is connected to a negative terminal. Oxygen (O₂) is also present above the water. Rust is labeled at the points where it forms on the iron.
 
Definisi of Metal Corrosion
  • Oxidation of metal through the action of air or oxygen gas, water and/or electrolyte.
  • The oxidised metal would release electron(s) to form an ion.
 
Metal Corrosion
  • In general, the more electropositive the metal is, the easier it is for the metal to corrode.
  • For example, corrosion of iron, Fe is faster than copper, Cu.
 
Prevention of Metal Corrosion
Using A More Electropositive Metal
  • Metal corrosion can be prevented by touching a more electropositive metal with the metal being protected.
  • The more electropositive metal releases electrons easier.
  • For example; a layer of zinc protecting an iron, but there is a scratch on the layer of zinc.
    • Zinc is more electropositive than iron.
    • Zinc is oxidised by releasing electrons.
    • Electrons flow to the surface of iron where there are water and oxygen.
  • Instead of iron being oxidised, the zinc layer acts as a sacrificial layer of protection as shown below:
Zinc Layer as A Sacrificial Layer of Protection
This image illustrates the process of galvanic corrosion protection. It shows a piece of iron with a zinc coating. A water drop is present on the surface, causing a reaction. The zinc (Zn) reacts to form Zn²⁺ ions and releases electrons (e⁻). These electrons are then used in the reduction of oxygen (O₂) and hydrogen ions (H⁺) to form water (H₂O). The labels highlight the zinc coating, water drop, and iron substrate.

 
Using A Less Electropositive Metal
  • A layer of less electropositive metal (e.g. tin) protecting iron can still prevent corrosion of iron.
  • However, if the protective layer is scratched, the corrosion of iron is enhanced instead of prevented.
  • The more electropositive metal releases electrons easier.
  • For example; a layer of tin protecting an iron, but there is a scratch on the layer of tin.
    • Iron is more electropositive than tin.
    • Iron is oxidised by releasing electrons.
    • Electrons flow to the surface of iron where there are water and oxygen.
  • Instead of tin being oxidised, the iron is oxidised at a faster rate because iron is more electropositive than tin.
 

 

 

 

 

 

 

 

The views expressed in the notes here are those of the contributor and don’t necessarily represent the views of Pandai.
Chapter : Redox Equilibrium
Topic : Investigate prevention of metal rusting
Form 5 Chemistry
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