text: 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. 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:
  : : oxidation of iron: 2 ()
  : : reduction of oxygen: 
  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:
  : : oxidation of iron: 2 ()
  : : reduction of oxygen: 
  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:<br />
  – Iron is a reductant giving electrons while being oxidized.<br />
  – 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 cations (Me,  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 (Me) chloride (MeCl) which hydrolyzes with water to produce the corresponding metal hydroxide (Me(OH)), 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. 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.