[AlaskanHumpback]

Hello all.

I am trying to figure what the general correlation is between the pH of a solution surrounding a piece of iron and the rate of corrosion of that iron. Any input would be greatly appreciated.

Thanks.

[cheah10]

To know how they are related, you have to know the mechanism of the reaction, and hence write out the rate equation. Have you learnt the rate equation?

[AlaskanHumpback]

To know how they are related, you have to know the mechanism of the reaction, and hence write out the rate equation. Have you learnt the rate equation?

I've looked over the general equations but I can't seem to figure out how the hydrogen Ions have any implication.

The iron is oxidized near the anodic site:
Fe(s) -> + Fe2+(aq) + 2e-
Reduction reaction on the cathodic site:
O2(g) + 2H2O(l) + 4e- -> 4OH-(aq)
The Iron ion and hydroxide ion then form Iron(II) hydroxide:
Fe2+(aq) + 2OH-(aq) -> Fe(OH)2(s)
The precipitate (ron(II) hydroxide) is then oxidized again to form rust (Fe2O3):
4Fe(OH)2(s) + O2(g) -> 2Fe2O3 •H2O(s) + 2H2O(l)

I found out that apparently at pH values below 4.0, ferrous oxide (which is not even in the equations above) is soluble. In this case the oxide dissolves as it is formed rather than depositing on the metal surface to form a protective film and slow down the reaction.
Source: http://www.corrosionist.com/what_is_effect_pH_Corrosion_rate.htm
Thanks for the reply.

[Borek]

O2(g) + 2H2O(l) + 4e- -> 4OH-(aq)

Fe2+(aq) + 2OH-(aq) -> Fe(OH)2(s)

Presence of OH^- means reaction is pH dependent.