When a droplet of water containing a little dissolved oxygen falls on an iron pipe, the iron under the droplet oxidizes:
Fe(s) Fe2+(aq) + 2 e-
The electrons are quickly snatched up by hydrogen ions and oxygen at the edge of the droplet to produce water:
4e- + 4 H+(aq) + O2(aq) 2 H2O(l)
These equations imply that increasing the concentration of hydrogen ions will cause more iron to oxidize.
More acidic water is expected to increase corrosion so long as there is a sufficient supply of oxygen. If the pH is very low, and there isn't enough oxygen, the hydrogen ions will snatch up the electrons anyway, making hydrogen gas instead of water:
2 H+(aq) + 2 e- H2(g)
But where's the rust? The equations above tell only a small part of the story.
Hydrogen ions are being consumed by the process. As the iron corrodes, the pH in the droplet rises. Hydroxide ions (OH-) appear in water as the hydrogen ion concentration falls. They react with the iron(II) ions to produce insoluble iron(II) hydroxides:
Fe2+(aq) + 2 OH-(aq)
The iron(II) ions also react with hydrogen ions and oxygen to produce iron(III) ions.
Fe2+(aq) + 4 H+(aq) + O2(aq)
4 Fe3+(aq) + 2 H2O(l)
The iron(III) ions react with hydroxide ions to produce hydrated iron(III) oxides (also known as iron(III) hydroxides).
Fe3+(aq) + 3 OH-(aq)
These can dry to make plain iron(III) oxide, Fe2O3. This is the red, powdery stuff we call "rust".
Since these processes involve hydrogen ions or hydroxide ions, they will be affected by changes in pH.
If you have other ions like calcium or carbonate present, they make a variety of precipitates that mix in with the iron hydroxide precipitates
to produce a crusty, gnarled coating which can slow corrosion under some circumstances by cutting the iron off from the acid, water, and air supply.
Author: Fred Senese email@example.com