Metal corrosion is a complex phenomenon that can be defined as a chemical alteration of a solid metal at the interface with the external environment (wet atmosphere, solution, etc.). The attack can extend in depth if the products formed are powdery (rust for example), but it can be limited to a thin protective layer (alumina on aluminum for example).In addition to surface action, corrosion may lead to a decrease in the mechanical properties of the metal (embrittlement).The economic impact of corrosion is considerable.For Iron, for example, corrosion losses are estimated at one quarter of the world’s steel production.
CONCEPT OF CORROSION
In reality, corrosion is a superficial deterioration of a material caused by electrical, chemical or mechanical phenomena.With the exception of some noble metals (Gold, Platinum for example) with a high redox potential, thermodynamics indicates a spontaneous tendency for metals to produce chemically more stable compounds (oxides for example).In fact the attack of the metal is a process of oxidation reduction of the electrochemical type which sees its origin in the character of electrical conduction of the two phases in presence: an electronic conduction in the metal phase and an ionic conduction in the electrolyte.
Electrochemical reactions allow charge transfers at the interface between the metal (electrode) and the electrolyte.
The corrosion of a metal such as iron is the result of a process that can be schematized by the following equation :
Fe → Fe2+ + 2 e– (1)
This equation simply means that the atoms of the metal (Fe) under the action of a corrosion process transform into positive ions (Fe2+) which leave the metal to pass into the medium.There is a change in the state of the material related to the dissolution reaction of the metal.This change of state is accompanied by the release of electrons (2 e- in the case of iron).The preservation of environmental neutrality requires that the electrons emitted by the dissolution reaction (1) be consumed in a second reaction which may be, for example :
2 H+ + 2 e– →H2 (2) acidic environment
½ O2 + H2O + 2 e– → 2OH– (2’) neutral and basic environment
The second reaction (2 and 2’) is related to the nature of the medium with which the material is in contact (electrolyte).
During a corrosion process, the surface of the material is therefore necessarily the seat of two reactions :
Dissolution or oxidation reaction (anodic reaction);
The decomposition reaction of the environment : «reduction» (cathodic reaction).
The process of corrosion, which superimposes anodic and cathodic reactions, is accompanied by the circulation of an electric current (circulation of electrons in the metal and ions in the electrolyte) between the two areas of the interface.
The electrons released by the anode (release of a Fe2+ ion in solution) are consumed by the cathodic process.To do this, the electrons flow through the metal from the anode to the cathode whose electrical potential is greater than that of the anode.
In order for a corrosion process to occur, it is therefore necessary to have :
a metal surface with anodic and cathodic zones, thus heterogeneity of potential;
an electrolyte for ion transport to close the electrical circuit to allow electron circulation.The electrolyte is provided by the medium (submerged soil, water, etc…).
When corrosion is widespread, the different points of the metal surface are successively cathodic and anodic and the loss of thickness is generally uniform.In this case, the corrosion rate can be assessed, for example, by measuring the thickness or weight loss, thus allowing to estimate the service life of the part concerned.
Unlike generalized corrosion, localized corrosion results from the localization of anodic areas in certain areas of the interface, for reasons of heterogeneity in the metal (metal defects, etc.) or in the electrolyte.Corrosion penetration is all the more rapid because the ratio of the anodic surface to the cathodic surface is low (insufficient protection by sacrificial anodes, etc…).
Pourbaix diagrams, also called potential-[pH] diagrams, indicate on an E-PH plane the areas of existence or predominance of an element.They are constructed from thermodynamic data.Figure 2 provides an example of a simplified Iron Pourbaix diagram.
According to the diagram in Figure 2, iron can be found in three thermodynamic states according to its potential and the pH of the solution in which it is immersed :
Thermodynamic passivation if it is in the stability domains of ferrous or ferric hydroxides (Fe2O3);
Activity or corrosion in ferrous and ferric ion stability domains (Fe3+, Fe2+, and HFe to extreme pH);
Immunity in the area of the diagram corresponding to the stability of the iron (Fe).
In the state of immunity, the metal and the surrounding environment are in a state of thermodynamic stability: they cannot react and therefore corrosion is not possible.This situation occurs naturally with noble metals (a potential far greater than that of the hydrogen electrode.Cf. Table 1).
In the state of passivity, the metal and the environment are not in a state of thermodynamic stability, but the metal is naturally covered with a protective film that isolates it from the external environment.This film (passive layer) must be thermodynamically stable vis-a-vis the external environment and shall not present any weakness locally.Otherwise, the metal corrodes locally.In common practice, stainless steel and titanium, for example, owe their resistance to corrosion to their passive behaviour.
In the state of activity, the metal is not thermodynamically stable, and is not covered with a protective film: it reacts with its environment and corrodes.
With reference to the electrochemical corrosion process shown in Figure 1, cathodic protection consists of lowering the electrochemical potential of a metal so that only cathodic reactions can occur on its surface.This potential reduction is achieved by injecting an external current at the metal surface to polarize it.