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Rusting
Rusting
Conditions for rusting of an iron
Oxygen gas
Water
Redox reaction of rusting
Step
Explanation
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
Metal corrosion
Definition: 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.
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.
Zinc layer as a sacrificial layer of protection
Prevention of metal corrosion 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.
Chapter : Redox Equilibrium
Topic : Rusting
Form 5
Chemistry
View all notes for Chemistry Form 5
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