Fundamentals of Metallic Corrosion in Fresh Water   -   7

During an investigation of the occurrence of heavy metals from corrosion of household plumbing, it was found that generally the concentration of metals decreases as the plumbing ages. The one exception was cadmium from galvanized plumbing, which was not detected in relatively new homes, but was found in significant concentrations from galvanized plumbing over 50 years old. It was demonstrated that the zinc used to galvanize the pipe contained approximately a half of a percent of cadmium. When the Cd - Zn alloy dissolved, the dissolved cadmium ions plated on the remaining alloy, thus gradually increasing the concentration of this toxic metal.

By now, if you have followed this discussion, you may be under the impression that all natural waters containing dissolved oxygen are severely corrosive. That this is not the case is due largely to the fact that many naturally occurring waters are capable of coating the cathode area of the metal with a thin layer of calcium carbonate. In order for this to occur, three conditions must be met: (1) the water must have a Langelier Index close to zero, (2) it must contain a significant bicarbonate ion concentration, and (3) it must be flowing over the metal surface.

Water has a Langelier Index of zero when it is in equilibrium with calcium carbonate [2]. When water is passed through a column of crushed limestone (a crystalline form of calcium carbonate), it has a negative index if some of the limestone is dissolved, a positive index if some calcium carbonate is precipitated, and a zero index if there is no change.

In chemical terms (See Appendix A) it can be shown that:

pH_s = log K₂/K_s - log (Ca⁺⁺) - log (HCO₃)

where pHs is the pH at which the water is just saturated with CaCO₃, (Ca⁺⁺) and (HC0₃) are the concentrations of calcium and bicarbonate respectively, K₂ is the second ionization constant of carbonic acid, and K_s is the solubility product of calcium carbonate. The values of the constants depend upon the temperature and the degree of mineralization (ionic strength) of the water.

At room temperature (25^oC) and moderate mineralization (400+-mg/l total dissolved solids) the above equation becomes:

pH_s = 11.85 - log (Ca⁺⁺) - log (HCO₃)

where both the calcium and bicarbonate are expressed as calcium carbonate.

To illustrate the use of the above equation, consider a typical ground water in the Santa Clara Valley. Calcium hardness is 100 mg/l and alkalinity is 200 mg/I. Both hardness and alkalinity are usually expressed as CaCO₃. Hence the saturation pH is:

11.85 - log (100) - log (200) = 11.85 - 2 - 2.3 = 7.55

This is close to the pH value actually found in the water. Since the coastal mountains surrounding the valley contain dolomite (a calcium magnesium carbonate), one would expect the ground water to be saturated with calcium carbonate.

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