Form of insidious localized corrosion in which a pit develops at the anode site
Severe pitting corrosion problems caused by chloride ions on a truss beam of the Nandu River Iron Bridge (Hainan Province, China) leading to the complete rupture of a metallic element.
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 (resulting due to corrosion) to the average penetration, which can be calculated based on the weight loss.
Development and kinetics of pitting
According to Frankel (1998) who performed a review on pitting corrosion, it develops in three successive steps: (1) initiation (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.[1]
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.[2] 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.[2]
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 the aggressive solution. The depassivation of the protective oxide layer is the less properly understood step in pitting corrosion and its very local and random appearance probably its most enigmatic characteristic. Mechanical or physical damages may locally disrupt the protective layer. Crystalline defects, or impurity inclusions, pre-existing in the base metal material can also serve as nucleation points (especially metal sulfide inclusions). The chemical conditions prevailing in the solution and the nature of the metal, or the alloy composition, are also important factors to take into consideration. Several theories have been elaborated to explain the depassivation process. Anions with weak or strong ligand properties such as chloride (Cl− ) and thiosulfate (S 2O2− 3) 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 (OH− ) 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.[3][4][5] The main interest of the point-defect model is to explain the stochastic character of the pitting corrosion process.
Pit growth
Schematic diagram showing the mechanism of localized corrosion with anodic zone (Fe oxidized into Fe inside the pit) and cathodic zone (O2 reduced into OH elsewhere outside the pit) developing on a metal immersed into an aqueous solution containing dissolved oxygen. Here, the pH conditions are neutral or alkaline (presence of OH ions in solution). The transport of ions occurs into the aqueous solution while electrons are transported from the anode to the cathode via the base metal (electrical conductor).
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.[6] 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:
Global redox reaction: 2 Fe + O2 + 4 H+ → 2 Fe2+ + 2 H2O
Acidic conditions favor the redox reaction according to Le Chatelier principle because the H+ ions added to the reagents side displace the reaction equilibrium to the right and also increase the solubility of the released Fe2+ cations.
Under neutral to alkaline conditions (pH > 7), the set of redox reactions given here above becomes the following:
Global redox reaction: 2 Fe + O2 + 2 H2O → 2 Fe(OH)2
The precipitation of Fe(OH)2 (green rust) can also contribute to drive the reaction towards the right. However, the solubility of Fe(OH)2 (Fe2+) is relatively high (~100 times that of Fe3+), but strongly decreases when pH increases because of common ion effect with the OH−.
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 e−) to the cathode (accepting e−) via the base metal (electrical conductor).
The localized production of positive metal cations (Men+, Fe2+ 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 anionsCl− ) 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 (Me) chloride (MeCln) which hydrolyzes with water to produce the corresponding metal hydroxide (Me(OH)n), and n H+ and n Cl– ions, accelerating the corrosion process.[7]
In the case of metallic iron, or steel, the process can be schematized as follows:[8]
Fe2+ + Cl− → [FeCl complex]+
[FeCl complex]+ + 2 H2O → Fe(OH)2 + 2 H+ + Cl−
Under basic conditions, such as under the alkaline conditions prevailing in concrete, the hydrolysis reaction directly consumes hydroxides ions (OH− ) while releasing chloride ions:
[FeCl complex]+ + 2 OH− → Fe(OH)2 + Cl−
So, when chloride ions present in solution enter in contact with the steel surface, they react with Fe2+ of the passive layer protecting the steel surface and form an iron–chloride complex. Then, the iron-chloride complex reacts with the OH− anions produced by the water dissociation and precipitates ferrous hydroxide (Fe(OH)2) while releasing chloride ions and new H+ 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.[9] 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),[10][11] or a damage to the protective coating. Polished surfaces display a higher resistance to pitting.[12]
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., Fe2+ in case of steel), anions need to migrate inside the narrow pit. It is worth to notice that the electromobilities of thiosulfate (S 2O2− 3) and chloride (Cl− ) anions are the highest after these of H+ and OH− 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 SO2−4 and NO−3.
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.[13]
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 (FeS2, 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.
Although in the aforementioned example, oxic conditions were always considered with the reduction of dissolved O2 in the cathodic zones, pitting corrosion may also occur under anoxic, or reducing, conditions. Indeed, the very harmful reduced species of sulfur (H2S, HS− , S2− , HS–S− , − S–S− , S0 and S 2O2− 3) can only subsist under reducing conditions.[14] Moreover, in the case of steel and stainless steel, reducing conditions are conducive to the dissolution of the protective oxide layer (dense γ-Fe 2O 3) because Fe2+ is much more soluble than Fe3+ , 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-Fe 2O 3 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.
Under strongly reducing conditions, in the absence of dissolved oxygen in water, or pore water of the ground, the electron acceptor (oxidizing agent) at the cathodic sites, where reduction occurs, can be the protons (H+) of water itself, the protons of hydrogen sulfide (H2S), or in acidic conditions in case of severe pyrite oxidation in a former oxic atmosphere, dissolved ferric ions (Fe3+ ), known to be very potent oxidizers. The presence of harmful reduced species of sulfur and microbial activity feeding the sulfur cycle (sulfide oxidation possibly followed by bacterial sulfate reduction) have also to be taken into account. Strictly abiotic (i.e. inorganic) corrosion processes are generally slower under anoxic conditions than under oxic conditions, but the presence of bacteria and biofilms can aggravate the degradation conditions and causes unexpected problems. Critical infrastructures and metallic components with very long service life may be susceptible to pitting corrosion: for example the metallic canisters and overpacks aimed to contain vitrified high-level radioactive waste (HLW) and spent nuclear fuel and to confine them in a water-tight enveloppe for several tenths of thousands years in deep geologic repositories.
Different types of corrosion inhibitor exist. Among them, oxidizing species such as chromate (CrO2− 4) and nitrite (NO− 2) were the first used to re-establish the state of passivation in the protective oxide layer. In the specific case of steel, the Fe2+ cation being a relatively soluble species, it contributes to favor the dissolution of the oxide layer which so loses its passivity. To restore the passivity, the principle simply consists to prevent the dissolution of the oxide layer by converting the soluble divalent Fe2+ cation into the much less soluble trivalent Fe3+ cation. This approach is also at the basis of the chromate conversion coating used to passivatesteel, aluminium, zinc, cadmium, copper, silver, titanium, magnesium, and tin alloys.[15]:p.1265[16]
As hexavalent chromate is a known carcinogen, its aqueous effluents can no longer be freely discharged into the environment and its maximum concentration acceptable in water is very low.
Nitrite is also an oxidizing species and has been used as corrosion inhibitor since the 1950's.[17][18][19] Under the basic conditions prevailing in concrete pore water nitrite converts the relatively soluble Fe2+ ions into the much less soluble Fe3+ ions, and so protects the carbon-steel reinforcement bars by forming a new and denser layer of γ-Fe 2O 3 as follows:
2 Fe2+ + 2 NO−2 + 2 OH− → Fe2O3 + 2 NO + H2O
Corrosion inhibitors, when present in sufficient amount, can provide protection against pitting. However, too low level of them can aggravate pitting by forming local anodes.
A single pit in a critical point can cause a great deal of damage. One example is the explosion in Guadalajara, Mexico on 22 April 1992, when gasoline fumes accumulated in sewers destroyed kilometers of streets. The vapors originated from a leak of gasoline through a single hole formed by corrosion between a steel gasoline pipe and a zinc-plated water pipe.[20]
Firearms can also suffer from pitting, most notably in the bore of the barrel when corrosive ammunition is used and the barrel is not cleaned soon afterwards.[21] Deformities in the bore caused by pitting can greatly reduce the firearm's accuracy.[22] To reduce pitting in firearm bores, most modern firearms have a bore lined with chromium.[23]
In laboratories, pitting corrosion may damage equipment, reducing its performance or longevity. Fume hoods are of particular concern, as the material constitution of their ductwork must suit the primary effluent(s) intended for exhaust.[25] If the chosen vent material is unsuitable for the primary effluent(s), consequent pitting corrosion will prevent the fume hood from effectively containing harmful airborne particles.[26]
Electrochemistry is the branch of physical chemistry concerned with the relationship between electrical potential difference and identifiable chemical change. These reactions involve electrons moving via an electronically-conducting phase between electrodes separated by an ionically conducting and electronically insulating electrolyte.
Rust is an iron oxide, a usually reddish-brown oxide formed by the reaction of iron and oxygen in the catalytic presence of water or air moisture. Rust consists of hydrous iron(III) oxides (Fe2O3·nH2O) and iron(III) oxide-hydroxide (FeO(OH), Fe(OH)3), and is typically associated with the corrosion of refined iron.
In chemistry and manufacturing, electrolysis is a technique that uses direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell. The voltage that is needed for electrolysis to occur is called the decomposition potential. The word "lysis" means to separate or break, so in terms, electrolysis would mean "breakdown via electricity".
Electroplating, also known as electrochemical deposition or electrodeposition, is a process for producing a metal coating on a solid substrate through the reduction of cations of that metal by means of a direct electric current. The part to be coated acts as the cathode of an electrolytic cell; the electrolyte is a solution of a salt of the metal to be coated; and the anode is usually either a block of that metal, or of some inert conductive material. The current is provided by an external power supply.
Redox is a type of chemical reaction in which the oxidation states of a substrate change. Oxidation is the loss of electrons or an increase in the oxidation state, while reduction is the gain of electrons or a decrease in the oxidation state.
Extractive metallurgy is a branch of metallurgical engineering wherein process and methods of extraction of metals from their natural mineral deposits are studied. The field is a materials science, covering all aspects of the types of ore, washing, concentration, separation, chemical processes and extraction of pure metal and their alloying to suit various applications, sometimes for direct use as a finished product, but more often in a form that requires further working to achieve the given properties to suit the applications.
Corrosion is a natural process that converts a refined metal into a more chemically stable oxide. It is the gradual deterioration of materials by chemical or electrochemical reaction with their environment. Corrosion engineering is the field dedicated to controlling and preventing corrosion.
In physical chemistry and engineering, passivation is coating a material so that it becomes "passive", that is, less readily affected or corroded by the environment. Passivation involves creation of an outer layer of shield material that is applied as a microcoating, created by chemical reaction with the base material, or allowed to build by spontaneous oxidation in the air. As a technique, passivation is the use of a light coat of a protective material, such as metal oxide, to create a shield against corrosion. Passivation of silicon is used during fabrication of microelectronic devices. Undesired passivation of electrodes, called "fouling", increases the circuit resistance so it interferes with some electrochemical applications such as electrocoagulation for wastewater treatment, amperometric chemical sensing, and electrochemical synthesis.
In chemistry, a reducing agent is a chemical species that "donates" an electron to an electron recipient. Examples of substances that are common reducing agents include the alkali metals, formic acid, oxalic acid, and sulfite compounds.
Sulfide (British English also sulphide) is an inorganic anion of sulfur with the chemical formula S2− or a compound containing one or more S2− ions. Solutions of sulfide salts are corrosive. Sulfide also refers to chemical compounds large families of inorganic and organic compounds, e.g. lead sulfide and dimethyl sulfide. Hydrogen sulfide (H2S) and bisulfide (SH−) are the conjugate acids of sulfide.
The chloralkali process is an industrial process for the electrolysis of sodium chloride (NaCl) solutions. It is the technology used to produce chlorine and sodium hydroxide, which are commodity chemicals required by industry. Thirty five million tons of chlorine were prepared by this process in 1987. The chlorine and sodium hydroxide produced in this process are widely used in the chemical industry.
A silver chloride electrode is a type of reference electrode, commonly used in electrochemical measurements. For environmental reasons it has widely replaced the saturated calomel electrode. For example, it is usually the internal reference electrode in pH meters and it is often used as reference in reduction potential measurements. As an example of the latter, the silver chloride electrode is the most commonly used reference electrode for testing cathodic protection corrosion control systems in sea water environments.
Electrogalvanizing is a process in which a layer of zinc is bonded to steel in order to protect against corrosion. The process involves electroplating, running a current of electricity through a saline/zinc solution with a zinc anode and steel conductor. Such Zinc electroplating or Zinc alloy electroplating maintains a dominant position among other electroplating process options, based upon electroplated tonnage per annum. According to the International Zinc Association, more than 5 million tons are used yearly for both hot dip galvanizing and electroplating. The plating of zinc was developed at the beginning of the 20th century. At that time, the electrolyte was cyanide based. A significant innovation occurred in the 1960s, with the introduction of the first acid chloride based electrolyte. The 1980s saw a return to alkaline electrolytes, only this time, without the use of cyanide. The most commonly used electrogalvanized cold rolled steel is SECC, acronym of "Steel, Electrogalvanized, Cold-rolled, Commercial quality". Compared to hot dip galvanizing, electroplated zinc offers these significant advantages:
Bacterial anaerobic corrosion is the bacterially-induced oxidation of metals. Corrosion of metals typically alters the metal to a form that is more stable. Thus, bacterial anaerobic corrosion typically occurs in conditions favorable to the corrosion of the underlying substrate. In humid, anoxic conditions the corrosion of metals occurs as a result of a redox reaction. This redox reaction generates molecular hydrogen from local hydrogen ions. Conversely, anaerobic corrosion occurs spontaneously. Anaerobic corrosion primarily occurs on metallic substrates but may also occur on concrete.
Anaerobic corrosion is a form of metal corrosion occurring in anoxic water. Typically following aerobic corrosion, anaerobic corrosion involves a redox reaction that reduces hydrogen ions and oxidizes a solid metal. This process can occur in either abiotic conditions through a thermodynamically spontaneous reaction or biotic conditions through a process known as bacterial anaerobic corrosion. Along with other forms of corrosion, anaerobic corrosion is significant when considering the safe, permanent storage of chemical waste.
Concrete degradation may have many different causes. Concrete is mostly damaged by the corrosion of reinforcement bars due to the carbonatation of hardened cement paste or chloride attack under wet conditions. Chemical damages are caused by the formation of expansive products produced by various chemical reactions, by aggressive chemical species present in groundwater and seawater, or by microorganisms. Other damaging processes can also involve calcium leaching by water infiltration and different physical phenomena initiating cracks formation and propagation. All these detrimental processes and damaging agents adversely affects the concrete mechanical strength and its durability.
The Schikorr reaction formally describes the conversion of the iron(II) hydroxide (Fe(OH)2) into iron(II,III) oxide (Fe3O4). This transformation reaction was first studied by Gerhard Schikorr. The global reaction follows:
Green rust is a generic name for various green crystalline chemical compounds containing iron(II) and iron(III) cations, the hydroxide (HO− ) anion, and another anion such as carbonate (CO2− 3), chloride (Cl− ), or sulfate (SO2− 4), in a layered double hydroxide structure. The most studied varieties are
Galvanic corrosion is an electrochemical process in which one metal corrodes preferentially when it is in electrical contact with another, in the presence of an electrolyte. A similar galvanic reaction is exploited in primary cells to generate a useful electrical voltage to power portable devices. This phenomenon is named after Italian physician Luigi Galvani (1737–1798).
Mineral processing and extraction of metals are very energy-intensive processes, which are not exempted of producing large volumes of solid residues and wastewater, which also require energy to be further treated and disposed. Moreover, as the demand for metals increases, the metallurgical industry must rely on sources of materials with lower metal contents both from a primary and/or secondary raw materials. Consequently, mining activities and waste recycling must evolve towards the development of more selective, efficient and environmentally friendly mineral and metal processing routes.
↑ Yu, Xiang; Al-Saadi, Saad; Zhao, Xiao-Ling; Rahman, R.K. Singh (October 2021). "Electrochemical Investigations of Steels in Seawater Sea Sand Concrete Environments". Materials. 14 (19): 5713. Bibcode:2021Mate...14.5713Y. doi:10.3390/ma14195713. PMC8510146. PMID34640111. It is known that mild steel rebar can form a protective passivation layer in the strong alkaline environment of concrete (when pH > 10). However, this layer can be destroyed by chloride ions and other harmful ions, diffusing from the ocean or industrial environment through the concrete matrix to the surface of the steel bars, causing rapid corrosion of the steel bars.
↑ K.H. Jürgen, Buschow, Robert W. Cahn, Merton C. Flemings, Bernhard Ilschner, Edward J. Kramer, and Subhash Mahajan (2001): Encyclopedia of Material – Science and Technology, Elsevier, Oxford, UK.
↑ Joseph H Osborne (2001): "Observations on chromate conversion coatings from a sol–gel perspective". Progress in Organic Coatings, volume 41, issue 4, pages 280-286. doi:10.1016/S0300-9440(01)00143-6
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