This type of corrosion occurs at sites where deposits allows a localized
concentration of a specific chemical, such as chloride or oxygen to be notably
different from the amount found in the bulk water environment. This corrosion
mechanism is considered a secondary reaction, whereas the primary reaction is
uniform metal wastage, or general corrosion. However, this secondary reaction
can be more devastating and unpredictable.
The sequence for metal failure due to under-deposit corrosion occurs as follows:
A deposit forms on the metal surface either from settling out of suspended
solids or precipitation of dissolved chemical species.
Under the deposit, dissolved oxygen is consumed by a primary corrosion
As the oxygen concentration under the deposit becomes depleted, and
significantly less than the oxygen in the bulk water, an electrolytic cell is
The area under the deposit becomes anodic to the surrounding area, and the
metal begins to corrode locally.
The corrosion, which began as primary corrosion, now has changed to a
secondary, differential cell corrosion system.
The rate of differential cell corrosion is proportional to the difference
in concentration of the depleted oxygen under the deposit and the oxygen
present in the bulk water.
Negatively charged chloride ions tend to migrate under the deposit to
balance the positively charged metal ions produced there. The high
concentration of chloride ions causes the area under the deposit to become
more acidic compared to the bulk solution, further enhancing the corrosion
under the deposit.
Likewise, severe concentration cell corrosion can involve the segregation
of any aggressive anions beneath deposits. Concentrations of sulfates and
chloride, in particular, are deleterious. As the corroding copper ions (Cu+)
leave the anodic surface, negatively charged chloride and sulfate ions diffuse
through the deposit to maintain neutrality, thus resulting in the
concentration of an aggressive acidic electrolyte.
A schematic of under-deposit corrosion is depicted below: