298
Fundamentals of Corrosion
of the copper had been appreciably reduced. Unprotected copper supplies a
different number of copper ions to poison fouling organisms.
In 1829, Edmund Davy was successful in protecting the iron portions of
buoys using zinc blocks; and in 1840, Robert Mallet produced a zinc alloy
that was particularly suited as a sacrificial anode. The fitting of zinc slabs
to hulls of vessels became standard practice as wooden hulls were replaced.
This provided localized protection, specifically against the galvanic action
of a bronze propeller. Overall protection of seagoing vessels was not investigated again until 1950 when the Canadian Navy determined that the proper
use of antifouling paints in conjunction with corrosion-resisting paints made
cathodic protection of ships feasible and could reduce maintenance costs.
About 1910 to 1912, the first application of cathodic protection by means of an
impressed current was undertaken in England and the United States. Since that
time, the use of cathodic protection has become widespread. There are thousands of miles of buried pipe and cables that are protected in this manner.
This form of protection is also used for water tanks, canal gates, submarines, marine piling, condensers, and chemical equipment.
9.2 Theory
Cathodic protection is achieved by applying electrochemical principles to
metallic components buried in soil or immersed in water. It is accomplished
by flowing a cathodic current through a metal–electrolyte interface, favoring
the reduction reaction over the anodic metal dissolution. This enables the
entire structure to work as a cathode.
The basis of cathodic protection is shown in the polarization diagram for
a copper-zinc cell in Figure 9.1. If polarization of the cathode is continued
using an external current beyond the corrosion potential to the open-circuit
potential of the anode, both electrodes reach the same potential and no corrosion of the zinc can take place.
Cathodic protection is accomplished by supplying an external current to
the corroding metal on the surface of which local action cells operate, as
shown in Figure 9.2. Current flows from the auxilliary anode and enters the
anodic and cathodic areas of the corrosion cells, returning to the source of
the DC current (B). Local action current will cease to flow when the entire
metal surface is at the same potential as a result of the cathodic area being
polarized by an external current to the open-circuit potential of the anodes.
As long as this external current is maintained, the metal cannot corrode.
There are two methods by which cathodic protection can be accomplished.
One is by coupling the structure with a more active metal, such as zinc or
magnesium. This produces a galvanic cell in which the active metal works as
an anode and provides a flux of electrons to the structure. The structure then
Fundamentals of Corrosion
of the copper had been appreciably reduced. Unprotected copper supplies a
different number of copper ions to poison fouling organisms.
In 1829, Edmund Davy was successful in protecting the iron portions of
buoys using zinc blocks; and in 1840, Robert Mallet produced a zinc alloy
that was particularly suited as a sacrificial anode. The fitting of zinc slabs
to hulls of vessels became standard practice as wooden hulls were replaced.
This provided localized protection, specifically against the galvanic action
of a bronze propeller. Overall protection of seagoing vessels was not investigated again until 1950 when the Canadian Navy determined that the proper
use of antifouling paints in conjunction with corrosion-resisting paints made
cathodic protection of ships feasible and could reduce maintenance costs.
About 1910 to 1912, the first application of cathodic protection by means of an
impressed current was undertaken in England and the United States. Since that
time, the use of cathodic protection has become widespread. There are thousands of miles of buried pipe and cables that are protected in this manner.
This form of protection is also used for water tanks, canal gates, submarines, marine piling, condensers, and chemical equipment.
9.2 Theory
Cathodic protection is achieved by applying electrochemical principles to
metallic components buried in soil or immersed in water. It is accomplished
by flowing a cathodic current through a metal–electrolyte interface, favoring
the reduction reaction over the anodic metal dissolution. This enables the
entire structure to work as a cathode.
The basis of cathodic protection is shown in the polarization diagram for
a copper-zinc cell in Figure 9.1. If polarization of the cathode is continued
using an external current beyond the corrosion potential to the open-circuit
potential of the anode, both electrodes reach the same potential and no corrosion of the zinc can take place.
Cathodic protection is accomplished by supplying an external current to
the corroding metal on the surface of which local action cells operate, as
shown in Figure 9.2. Current flows from the auxilliary anode and enters the
anodic and cathodic areas of the corrosion cells, returning to the source of
the DC current (B). Local action current will cease to flow when the entire
metal surface is at the same potential as a result of the cathodic area being
polarized by an external current to the open-circuit potential of the anodes.
As long as this external current is maintained, the metal cannot corrode.
There are two methods by which cathodic protection can be accomplished.
One is by coupling the structure with a more active metal, such as zinc or
magnesium. This produces a galvanic cell in which the active metal works as
an anode and provides a flux of electrons to the structure. The structure then