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24.8: Electrochemical Corrosion

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    Make sure you thoroughly understand the following essential ideas. It is especially important that you know the precise meanings of all the highlighted terms in the context of this topic.



    In a sense, corrosion can be viewed as the spontaneous return of metals to their ores; the huge quantities of energy that were consumed in mining, refining, and manufacturing metals into useful objects is dissipated by a variety of different routes. The economic aspects of corrosion are far greater than most people realize; the estimated cost of corrosion in the U.S. alone was $276 billion per year. Of this, about $121 billion was spent to control corrosion, leaving the difference of $155 billion as the net loss to the economy. Utilities, especially drinking water and sewer systems, suffer the largest economic impact, with motor vehicles and transportation being a close second.

    Corrosion Cells and Reactions

    The special characteristic of most corrosion processes is that the oxidation and reduction steps occur at separate locations on the metal. This is possible because metals are conductive, so the electrons can flow through the metal from the anodic to the cathodic regions (Figure \(\PageIndex{1}\)). The presence of water is necessary in order to transport ions to and from the metal, but a thin film of adsorbed moisture can be sufficient.

    Figure \(\PageIndex{1}\): Corrosion is a two-step process. Figure \(\PageIndex{1}\): Electrochemical corrosion of iron. Corrosion often begins at a location (1) where the metal is under stress (at a bend or weld) or is isolated from the air (where two pieces of metal are joined or under a loosely-adhering paint film.) The metal ions dissolve in the moisture film and the electrons migrate to another location (2) where they are taken up by a depolarizer. Oxygen is the most common depolarizer; the resulting hydroxide ions react with the Fe2+ to form the mixture of hydrous iron oxides known as rust.


    \[\ce{Fe(s) \rightarrow Fe^{2+}(aq) + 2 e^{-}} \label{1.7.1}\]

    and the cathodic steps may invove the reduction of oxygen gas

    \[ \ce{O_2} + \ce{2 H_2O} + \ce{4e^{-}} \rightarrow \ce{4 OH^{-}} \label{1.7.2}\]

    or the reduction of protons

    \[ \ce{H^{+} + e^{-}} \rightarrow \ce{1/2 H2(g)} \label{1.7.2b} \]

    or the reduction of a metal ion

    \[\ce{M^{2+} + 2 e^{–}} \rightarrow \ce{M(s)} \label{1.7.2c}\]

    where \(\ce{M}\) is a metal.


    Figure \(\PageIndex{2}\): Schematic diagram of corrosion cells on iron.

    If one part of a metallic object is protected from the atmosphere so that there is insufficient \(\ce{O2}\) to build or maintain the oxide film, this "protected" region will often be the site at which corrosion is most active. The fact that such sites are usually hidden from view accounts for much of the difficulty in detecting and controlling corrosion.

    Figure \(\PageIndex{3}\): Pitting corrosion Most metals are covered with a thin oxide film which inhibits anodic dissolution. When corrosion does occur, it sometimes hollows out a narrow hole or pit in the metal. The bottoms of these pits tend to be deprived of oxygen, thus promoting further growth of the pit into the metal.

    In contrast to anodic sites, which tend to be localized to specific regions of the surface, the cathodic part of the process can occur almost anywhere. Because metallic oxides are usually semiconductors, most oxide coatings do not inhibit the flow of electrons to the surface, so almost any region that is exposed to \(\ce{O2}\) or to some other electron acceptor can act as a cathode. The tendency of oxygen-deprived locations to become anodic is the cause of many commonly-observed patterns of corrosion.

    Rusted-out Cars and Bathroom Stains


    Figure \(\PageIndex{4}\): Rusted out car. Severely rusted floorboards of 1990 Chrysler New Yorker. Car was damaged to the point where it became unsafe to drive. Under the rusting section between the two holes were the two brake lines (Public Domain; Bige1977 via Wikipedia).


    Dissimilar-metal corrosion can occur even if the two metals are not initially in direct contact. For example, in homes where copper tubing is used for plumbing, there is always a small amount of dissolved \(\ce{Cu^{2+}}\) in the water. When this water encounters steel piping or a chrome-plated bathroom sink drain, the more-noble copper will plate out on the other metal, producing a new metals-in-contact corrosion cell. In the case of chrome bathroom sink fittings, this leads to the formation of \(\ce{Cr^{3+}}\) salts which precipitate as greenish stains.

    Control of Corrosion


    \[\ce{M -> M^{2+} + 2 e^{-}}.\]

    Protection Method 1: Sacrificial Coatings


    Figure \(\PageIndex{5}\): Sacrificial Coatings


    Protection Method 2: Cathodic Protection


    Figure \(\PageIndex{6}\): Cathodic Protection