BACKGROUND
By definition, water is described as "corrosive to a metal" if it disintegrates the metal, or if any substances in the water react with the metal at the plane of the water/metal interface. Corrosion is an electrochemical process of the wearing away or the rusting away of the material due to chemical or physical reactions. (Erosion of metal is a physical process.) Many raw water sources and treated waters are corrosive, and therefore corrosion prevention is a necessary treatment process.

We usually think of corrosion control as involving protection of equipment and pipes. The "Lead and Copper Rule" has added an additional element to the scope which is the "quantity of lead and copper which may be in the water." This rule now dictates the principal goal of corrosion control. We will discuss the methods of control in the Operator Notebook section Stabilization, more in depth.

Goal: the goal is to produce a non-corrosive water. (Which also according to the additions of the Lead & Copper Rule, is one that will not leach out lead or copper from the conveying pipes or pipe fittings.)

The Chemistry
Metals are normally found as oxides in their normal state the environment. The metal, as we most often use it, is in it’s reduced state. The metals obviously try to return to their "oxidized" natural states of matter.

Examples are:
a) Aluminum corrosion yields a white aluminum oxide deposit on the metal’s surface.
b) Iron corrosion is described as "pitting" and other reactions like "tuberculation."
c) Lead pipe carrying soft water which is high in carbon dioxide gases may corrode the lead pipe (which also forms toxic lead salts in the process).
d) Steel pits and rusts.
e) "Red water" caused by the breaking up of iron tubercules or iron precipitated nodules.

The following figures, Figures 1 through 4, describe the basic process of corrosion, as illustrated with an iron pipe:

Figure 1 Corrosion generates electricity like a battery, in what is being corroded.

Minor impurities in the metal starts the process, by creating an electrical current consisting of 2 electrons (2 e- ) in the metal. The electrons move from the portion of the metal we call the "Anode" to the place we call the "Cathode".

Figure 2

Most waters: H2O dissociates H+ + OH-
At the Cathode: two hydrogen ions combine with the two electrons to become hydrogen gas (H2 ) in the water.
At the Anode: The Fe+2 combines with 2 OH- radicals to form Fe(OH)2 ,which is ferrous hydroxide.

Figure 3

But this leaves excess hydrogen ions ( H+ ) at the anode, and hydroxyl ions ( OH- ) at the cathode, which increases the rate of pitting (corrosion) at the anode area.

Figure 4

If we add dissolved oxygen to the water we have created ferric hydroxide , better known as "rust". The ferric hydroxide precipitates to now form "tubercules", the deposits in the pipes. With low dissolved oxygen in the rusted area, it will form an encrustation of Fe(OH)2 next to this area, and an outer layer of Fe(OH3) will cover that region. This is much like a living, "growing crystal."

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