Organic Corrosion Inhibitors

5 A cathodic reactant

Inside this electrochemical cell, electrons depart from anodic to cathodic sites. The charged particles, ions, move across the conducting solution to balance the electrons flow. Anions (from cathodic reactions) move toward the anode and cations (from the anode itself) drift toward the cathode. Resultantly, anode corrodes and the cathode does not. There also exists a voltage/potential difference amidst anode and cathode. Numerous discrete micro cells develop on the metal surfaces, due to the constitutional phase difference, from stress variations, coatings, and imperfection levels like dislocations, grain boundaries, kink sites, or from ionic conductivity alterations or compositional changes in the conducting solution. The corrosion process is chemically spontaneous oxidation of the metal on reaction with the cathodic reactant. Every similar cell reaction results from a pair of simultaneous anodic and cathodic reactions going on at identical rates on the surface of metal.

Figure 1.2 An electrochemical cell.

At the anode, the metal corrodes. The anodic reaction is the oxidation of a metal to its ionic form when the electric charge difference exists at the solid–liquid interface. Generally, anodic reaction is an oxidation reaction of a metal to its metal ions, which passes into conductive solution:

(1.1)

where “n” is the metallic valence, e− is the electron, M is metal, and M n+its metalion.

The cathodic reaction involves the environment and can be represented by the following reaction:

(1.2)

where R +is the positive ion present in the electrolyte, e −is the metallic electron, and R 0is the reduced species. Based on the environment, many cathodic reactions and electron consuming reactions are possible. The main reactions are as follows.

The anaerobic acidic aqueous environment

(1.3)

In the anaerobic alkaline aqueous environment

(1.4)

In the aerobic acidic aqueous environment

(1.5)

In the aerobic alkaline aqueous environment

(1.6)

Some other reactions that are most commonly present in the chemical process are following.

Metal ion reduction

(1.7)

Metal ion deposition

(1.8)

The products of the anodic and cathodic reactions react to form solid corrosion products on the surface of the metal. The Fe 2+interacts with OH −ions as:‐

(1.9)

Fe(OH) 2is reoxidized to Fe(OH) 3, an unstable product, and thus transforms to hydrated ferric oxide commonly called as red rust ( Figure 1.3).

(1.10)

(1.11)

Figure 1.3 Mechanism of rust formation.

Figure 1.4 Classified forms of corrosion.

Seldom is a single class of corrosion discovered in corroding structures. Different metals in contact and contact with different environment hardly allow only one type of corrosion to occur even within a system. Each type of corrosion is caused by their specific reaction mechanisms and has their specific monitoring, prediction, and control methods. Figure 1.4throws some light on classification of corrosion in a pictorial manner. None of the classifications is a universal standard, even the following classification is an adapted [4, 13].

This type of corrosion affects a large patch over the metal and causes overall reduction of metallic thickness subject to the fact that metal undergoing corrosion has a uniform composition and metallurgy too. What happens is that anode and cathode do not possess fixed sites; as such there are no sites preferable to corrosion, which occurs here in a uniform fashion. Corrosion rates are easily monitored by electrochemical measuring techniques or gravimetric analysis. A metal suffering from uniform corrosion can be protected using corrosion inhibitors or coatings and also by cathodic protection. Atmospheric corrosion is an example of uniform corrosion. When exposed to dry atmospheres with very less humidity, metals spontaneously tend to form an oxide film. This barrier oxide film acquires a thickness of 2–5 nm [14].

Pitting corrosion is highly destructive form and a kind of localized attack, which leads to little holes called pits in metal. Small cavities and holes, which are as deep as their diameter, are known as pits. They cause perforations by penetrating into the metal with least loss of weight [15]. Pitting is proportional to the logarithm of electrolyte’s concentration of chloride. The prerequisite for pitting to occur is that the electrolyte should be a strong oxidizer for onset of the passive state. The ferric and cupric halide ions are electron acceptors (cathodic reactants), and they do not need oxygen to initiate and propagate pitting. Other propagating factors causing pitting include localized damage chemically and mechanically to a passive oxide film, non‐metallic impurities/non‐uniformities of metal structure due to nonproportional inhibitor coverage. Pitting, however, can be evaded by reducing aggressiveness of the solution, decreasing the temperature of conductive solution, decreasing Cl− concentration and acidity.