• 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.

• Oxygen gas.
• Water.

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, \(Fe_2O_3.3H_2O\). 
  • \( 4Fe(OH)_3(s)\rightarrow Fe_2O_3.3H_2O(s)\\ \phantom{ 4Fe(OH)_3(s)\rightarrow}Rust\)

• 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.

• 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.

• 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

• 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.

• 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.

• Oxygen gas.
• Water.

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, \(Fe_2O_3.3H_2O\). 
  • \( 4Fe(OH)_3(s)\rightarrow Fe_2O_3.3H_2O(s)\\ \phantom{ 4Fe(OH)_3(s)\rightarrow}Rust\)

• 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.

• 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.

• 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

• 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.