Corrosion in aqueous solutions proceeds by an electrochemical process, and anodic and cathodic electrochemical reactions mmust occur simultaneously. No net overall, charge builds up on the metal as a result of corrosion since the rate of the anodic and cathodic reactions are equal.
Anodic reactions involve oxidation of metal to its ions, e.g. for steel the following reaction occurs.
Fe > Fe2+ + 2e (1)
The cathodic process involves reduction and several reactions are possible. In acidic water, where hydrogen ions (H+) are plentiful, the following reaction occurs.
2H+ + 2e > H2 (2)
In alkaline solutions, where hydrogen ions are rare, the reduction of water will occur to yield alkali and hydrogen.
2H2O + 2e > H2 + 2OH- (3)
However, unless the water is deaerated reduction of oxygen is the most likely process, again producing alkali at the surface of the metal.
O2 + 2H2O + 4e > 4OH- (4)
Reactions (1) and (2) are shown schematically in Fig 1 where anodic and cathodic sites are nearby on the surface of a piece of metal. The rate of these two reactions can be changed by withdrawing electrons or supplying additional electrons to the piece of metal. It is an established principle that if a change occurs in one of the factors under which a system is in equilibrium, the system will tend to adjust itself so as to annul, as far as possible, the effect of that change. Thus, if the electrons is withdrawn from the piece of metal the rate of reaction (1) will increase to attempt to offset the action and the dissolution of iron will increase, whereas reaction (2) will decrease. Conversely, if additional electrons supplied from an external source to the piece of metal, reaction (1) will decrease to give reduced corrosion and reaction (2) will increase. The latter case will apply to cathodic protection. Thus, to prevent corrosion it’s has to continue to supply electrons to the steel from an external source to satisfy the requirements of the cathodic reaction. Note that the anodic and cathodic processes are inseparable. Reducing the rate of the anodic process will allow the rate of the cathodic process to increase.
These principles may be expressed in a more quantitative manner by plotting the potential of the metal against the logarithm of the anodic and cathodic reaction rates expressed as current densities. Typical anodic and cathodic curves are illustrated in Fig 2. The corrosion current, Icorr, and the corrosion potential, Ecorr, occur at the point of intersection of the anodic and cathodic curves, i.e. where anodic and cathodic reactions rates are equal. If electrons are “pumped” into the metal to make it more negative the anodic dissolution of iron is decreased to a negligible rate at a potential EI, whereas the rate of the cathodic current is increased to I1. Hence, a current I1 must be supplied from an external source to maintain the potential at E1 where the rate of dissolution of the iron is at a low value. If the potential is reduced to E2 (Fig 2) the current required from the external source will increase to I2. Further protection of the metal is insignificant, however, and the larger current supplied from the external source is wasted. The metal is then said to be over-protected.
In aerated neutral or alkaline solutions the cathodic corrosion process is usually the reduction of oxygen. The kinetics of this cathodic process are controlled by the rate at which oxygen can diffuse to the surface of the metal, which is slower than the rate of consumption of oxygen by the cathodic reaction. Thus, the rate of this reaction does not increase as the potential of the metal is made more negative but remains constant unless the rate of supply of oxygen to the surface of the metal is increased by, for example, increase fluid flow rate. The influence of flow velocity on cathodic protection parameters is illustrated in Fig 3. A current of I1 is initially required to maintain the metal at the protection potential E1. However, if the flow rate is increased the limiting current for the reduction of oxygen is increased (dotted line) and the current required to maintain the metal at the protection potential is increased by &DeltaI. Thus, the current density required to maintain the correct protection potential will vary with service conditions. Clearly, cathodic current density is not a good guide as to whether a structure is cathodically protected. The correct protection potential must be maintained if corrosion is to be prevented.
If the structure is over-protected and the potential is reduced to a potential region where reduction of water (reaction 3) can take place, further current will be required from the external source and current will be wasted. In Fig 3 reducing the potential from E1 to E2 will increase the current required from the external source from I1 to I2 as a result of an increased rate of reduction of water.
Excessive negative potentials can cause accelerated corrosion of lead and aluminium because of the alkaline environments created at the cathode. These alkaline conditions may also be detrimental to certain paint systems, and may cause loss of the paint film. Hydrogen evolution at the cathode surface may, on high-strength steels, result in hydrogen embrittlement of the steel, with subsequent loss of strength. It may also cause disbanding of any insulating coating: the coating would then act as an insulating shield to the cathodic protection currents.