Pitting corrosion, or pitting, is a form of extremely localized corrosion that leads to the random creation of small holes in metal. The driving power for pitting corrosion is the depassivation of a small area, which becomes anodic (oxidation reaction) while an unknown but potentially vast area becomes cathodic (reduction reaction), leading to very localized galvanic corrosion. The corrosion penetrates the mass of the metal, with a limited diffusion of ions.
Another term arises, pitting factor, which is defined as the ratio of the depth of the deepest pit (from localized corrosion) to the average penetration depth (mean thickness of the corrosion layer produced by the general uniform corrosion), which can be calculated based on the weight loss and corrosion products density.
Development and kinetics of pitting According to Frankel (1998) who performed a review on pitting corrosion, it develops in three successive steps: (or nucleation) by breakdown of the passive film protecting the metal surface from oxidation, (2) growth of metastable pits (growing up to the micron scale and then repassivating), and (3) the growth of larger and stable pits.
The evolution of the pit density (number of pits per surface area) as a function of time follows a sigmoid curve with the characteristic shape of a logistic function curve, or a hyperbolic tangent. Guo et al. (2018), after a statistical analysis of hundreds of individual pits observed on carbon steel surfaces at the nano-to-micro- scales, distinguish three stages of pitting corrosion: induction, propagation, and saturation.
Mechanism The pit formation can be essentially regarded as a two step process: nucleation followed by a growth.
Depassivation of the protective layer The process of pit nucleation is initiated by the depassivation of the protective oxide layer isolating the metal substrate from a corrosive solution. The depassivation step is the least-understood step in pitting corrosion, and can be triggered by a variety of factors. Mechanical or physical damage may locally disrupt the protective layer. Crystalline defects, or impurity inclusions, pre-existing in the base metal material can also serve as nucleation points (metal sulfide inclusions in particular). The chemical conditions prevailing in the solution and the nature of the metal, or the alloy composition, of course, have a strong influence as well. Several theories have been elaborated to explain the depassivation process. Anions with weak or strong ligand properties such as chloride () and thiosulfate () respectively can complex the metallic cations (Men+) present in the protective oxide layer and so contribute to its local dissolution. Chloride anions could also compete with hydroxide ions () for the sorption onto the oxide layer and start to diffuse into the porosity or the crystal lattice of the oxide layer. Finally, according to the point-defect model elaborated by Digby Macdonald, the migration of crystal defects inside the oxide layer could explain its random localized disappearance. The main interest of the point-defect model is to explain the stochastic character of the pitting corrosion process.
Pit growth
The more common explanation for pitting corrosion is that it is an autocatalytic process driven by the random formation of small electrochemical cells with separate anodic and cathodic zones. The random local breakdown of the protective oxide layer and the subsequent oxidation of the underlying metal in the anodic zones result in the local formation of a pit where acid conditions are maintained by the spatial separation of the cathodic and anodic half-reactions. This creates a gradient of electrical potential and is responsible for the electromigration of aggressive anions into the pit.ASM Handbook, Volume 13, "Corrosion", , ASM International, 1987 For example, when a metal is exposed to an oxygenated aqueous solution containing sodium chloride (NaCl) as electrolyte, the pit acts as anode (metal oxidation) and the metal surface acts as cathode (oxygen reduction).
In the case of pitting corrosion of iron, or carbon steel, by atmospheric oxygen dissolved in acidic water (pH < 7) in contact with the metal exposed surface, the reactions respectively occurring at the anode and cathode zones can be written as follows:
Anode: oxidation of iron: 2 ()
Cathode: reduction of oxygen:
Global redox reaction:
Acidic conditions favor the redox reaction according to Le Chatelier principle because the ions added to the reagents side displace the reaction equilibrium to the right and also increase the solubility of the released cations.
Under neutral to alkaline conditions (pH > 7), the set of redox reactions given here above becomes the following:
Anode: oxidation of iron: 2 ()
Cathode: reduction of oxygen:
Global redox reaction:
The precipitation of (green rust) can also contribute to drive the reaction towards the right. However, the solubility of () is relatively high (~ 100 times that of ), but strongly decreases when pH increases because of common ion effect with the .
In the two examples given here above: – Iron is a reductant giving electrons while being oxidized. – Oxygen is an oxidant taking up electrons while being reduced.
The formation of anodic and cathodic zones creates an electrochemical cell (i.e., a small electric battery) at the surface of the affected metal. The difference in Gibbs free energy (ΔG) drives the reaction because ΔG is negative and the system releases energy (enthalpy, ΔH < 0) while increasing entropy (ΔG = ΔH - TΔS).
The transport of dissolved ions occurs into the aqueous solution in contact with the corroding metal while electrons are transported from the anode (giving ) to the cathode (accepting ) via the base metal (electrical conductor).
The localized production of positive metal (M) cations (Mn+, in the example here above) in the pit (oxidation: anode) gives a local excess of positive charges which attract the negative ions (e.g., the highly mobile chloride anions ) from the surrounding electrolyte to maintain the electroneutrality of the ion species in aqueous solution in the pit. The pit contains a high concentration of metal chloride (MCln) which hydrolyzes with water to produce the corresponding metal hydroxide (M(OH)n), and n H+ and n Cl– ions, accelerating the corrosion process.
In the case of metallic iron, or steel, the process can be schematized as follows:
Under basic conditions, such as under the alkaline conditions prevailing in concrete, the hydrolysis reaction directly consumes hydroxides ions () while releasing chloride ions:
So, when chloride ions present in solution enter in contact with the steel surface, they react with of the passive layer protecting the steel surface and form an iron–chloride complex. Then, the iron-chloride complex reacts with the anions produced by the water dissociation and precipitates ferrous hydroxide () while releasing chloride ions and new ions available to continue the corrosion process.
In the pit, the oxygen concentration is essentially zero and all of the cathodic oxygen reactions take place on the metal surface outside the pit. The pit is anodic (oxidation) and the locus of rapid dissolution of the metal.Princeton.edu, pitcorrosion. The metal corrosion initiation is autocatalytic in nature however its propagation is not.
This kind of corrosion is often difficult to detect and so is extremely insidious, as it causes little loss of material with the small effect on its surface, while it damages the deep structures of the metal. The pits on the surface are often obscured by corrosion products. Pitting can be initiated by a small surface defect, being a scratch or a local change in the alloy composition (or local impurities, e.g. metallic sulfide inclusions such as MnS or NiS), or a damage to the protective coating. Polished surfaces display a higher resistance to pitting.
Capillary electrophoresis in the pit In order to maintain the solution electroneutrality inside the pit populated by cations released by oxidation in the anodic zone (e.g., in case of steel), anions need to migrate inside the narrow pit. It is worth to notice that the electromobilities of thiosulfate () and chloride () anions are the highest after these of and ions in aqueous solution. Moreover, the molar conductivity of thiosulfate ions is even higher than that of chloride ions because they are twice negatively charged (weak base reluctant to accept a proton). In capillary electrophoresis, thiosulfate moves faster than chloride and eluates before this latter. The high electromobility of both anions could also be one of the many factors explaining their harmful impact for pitting corrosion when compared with other much less damaging ion species such as and .
Susceptible alloys and environment conditions Pitting corrosion is defined by localized attack, ranging from microns to millimeters in diameter, in an otherwise passive surface and only occurs for specific alloy and environmental combinations. Thus, this type of corrosion typically occurs in alloys that are protected by a tenacious (passivating) oxide film such as stainless steels, nickel alloys, aluminum alloys in environments that contain an aggressive species such as chlorides (Cl–) or thiosulfates (S2O32–). In contrast, alloy/environment combinations where the passive film is not very protective usually will not produce pitting corrosion. A good example of the importance of alloy/environment combinations is carbon steel. In environments where the pH value is lower than 10, carbon steel does not form a passivating oxide film and the addition of chloride results in uniform attack over the entire surface. However, at pH greater than 10 (alkaline) the oxide is protective and the addition of chloride results in pitting corrosion.
Besides chlorides, other anions implicated in pitting include thiosulfates (S2O32−), fluorides and iodides. Stagnant water conditions with low concentrations of dissolved oxygen also favor pitting. Thiosulfates are particularly aggressive species and are formed by partial oxidation of pyrite (, a ferrous disulfide), or partial sulfate reduction by microorganisms, a.o. by sulfate reducing bacteria (SRB). Thiosulfates are a concern for corrosion in many industries handling sulfur-derived compounds: sulfide ores processing, oil wells and pipelines transporting soured oils, kraft paper production plants, photographic industry, methionine and lysine factories.
Influence of redox conditions
Although in the aforementioned example, oxic conditions were always considered with the reduction of dissolved in the cathodic zones, pitting corrosion may also occur under anoxic, or reducing, conditions. Indeed, the very harmful reduced species of sulfur (, , , , , S0 and ) can only subsist under reducing conditions. Moreover, in the case of steel and stainless steel, reducing conditions are conducive to the dissolution of the protective oxide layer (dense γ-) because is much more soluble than , and so reducing conditions contribute to the breakdown of the protective oxide layer (initiation, nucleation of the pit). Reductants exert thus an antagonist effect with respect to the oxidants (chromate, nitrite) used as corrosion inhibitors to induce steel repassivation via the formation of a dense γ- protective layer. Pitting corrosion can thus occur both under oxidizing and reducing conditions and can be aggravated in poorly oxygenated waters by differential aeration, or by drying/wetting cycles.