Fundamentals of Metallic Corrosion in Fresh Water   -   5

Table I shows the galvanic series. This is similar to the electrochemical series shown in Table 11, except that the latter is based on thermodynamically reversible reactions, which are not practically attainable. The galvanic series is supposed to represent what happens in the real world. As in the real world, however, there are some surprises. For example, although zinc is higher than iron in the table, at 140o F the potentials are reversed. This was first recognized in the early 1950's. Prior to that time it was standard practice to galvanize the interior of water heaters, supposedly protecting the steel, but actually causing more rapid failure. Although magnesium is above aluminum in both tables, it was found that severe attack of aluminum in the vicinity of magnesium rivets occurred in the hulls of flying boats made during World War 11. Investigation revealed that the hydroxyl ions resulting from the reduction of oxygen attacked the aluminum:

Al + 30H^- = H₃A1O₃

You will note that aluminum is fairly high in the table and it may surprise you that Alcoa claims it "will not rust or rot". Aluminum holds up well when exposed to air, thanks to a continuous and highly adherent oxide layer, but is generally unsatisfactory in fresh water environments.

I have discussed the galvanic cell in considerable detail because it is well suited to explain the electrochemical nature of corrosion. However, the galvanic cell is responsible for only a small fraction of corrosion that occurs in potable waters.

The principal cause of corrosion in water is the oxygen concentration cell. This very important mechanism has not been given the attention it deserves in most publications on corrosion, perhaps because it introduces an apparent anomaly: Oxidation of metal occurs at a site where there is no oxygen.

Figure 5 depicts an oxygen concentration cell. Note that the chemical reactions involved are precisely the same as those that occur in a galvanic cell, and since voltage produced by the cell is determined by the chemical reactions, the potential of any oxygen concentration cell will be exactly the same as in a galvanic cell where the corroding metal is the anode. Polarization characteristics will depend on how this metal behaves as a cathode.

The oxygen concentration cell may be initiated by anything that will shield a small area from the dissolved oxygen in the water, such as a grain of sand or a microbial colony. Once started, the cell becomes self-perpetuating. A pit forms that is covered with a crust of metal oxide, assuring there will be no oxygen under the tubercle that is formed.

When the corroding metal is iron or steel, an additional reaction occurs. The ferrous ions produced are oxidized to ferric hydroxide:

4Fe⁺⁺ + O₂ + 10H₂O = 4Fe(OH)₃ + 8H⁺

The interior of a tubercle contains a solution of ferrous chloride and sulfate ions in concentrations greater than those in the water. Hydrogen sulfide is occasionally present. The solution is slightly acid and may have a pH of approximately 6. This liquid is covered by a black inner crust consisting of hydrous Fe₃O₄, which, being magnetic, is attracted to the iron to form porous columnar fibers. The outer crust consists of reddish brown ferric hydroxide or hydrated ferric oxide. The flow of current protects the metal in the immediate vicinity of the pit. In ferrous metals, pits generally become inactive after a period of time. When this occurs, they no longer protect the metal in their vicinity, and new pits develop. Apparently the tubercles become so impermeable that ions cannot diffuse through, and since the solution inside must maintain electrical neutrality, no additional iron ions are formed.

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