Corrosion is a galvanic process that can occur in both single metals and when two dissimilar metals are in contact. In a single metal, localized differences in the material can cause one region to act as the anode and another as the cathode. Electrons flow between these regions through the metal. For iron, the anodic reaction produces Fe2+ ions while the cathodic reaction typically reduces oxygen to hydroxide ions. These ions then react to form rust. When two metals contact, they form a corrosion cell where one metal corrodes preferentially to the other. Common examples are rusting car bodies at joints and green stains from copper plumbing contacting steel or chrome.

1. Corrosion
Corrosion is a galvanic process by which metals deteriorate through oxidation—usually but not
always to their oxides. For example, when exposed to air, iron rusts, silver tarnishes, and copper
and brass acquire a bluish-green surface called a patina.
The various metals subject to corrosion, iron is by far the most important commercially. An
estimated $100 billion per year is spent in the United States alone to replace iron-containing
objects destroyed by corrosion. Consequently, the development of methods for protecting metal
surfaces from corrosion constitutes a very active area of industrial research. In this section, we
describe some of the chemical and electrochemical processes responsible for corrosion. We also
examine the chemical basis for some common methods for preventing corrosion and treating
corroded metals.

2. Corrosion is a REDOX process.
• Under ambient conditions, the oxidation of most metals is thermodynamically spontaneous, with the notable
exception of gold and platinum. Hence it is actually somewhat surprising that any metals are useful at all in
Earth’s moist, oxygen-rich atmosphere. Some metals, however, are resistant to corrosion for kinetic reasons. For
example, aluminum in soft-drink cans and airplanes is protected by a thin coating of metal oxide that forms on
the surface of the metal and acts as an impenetrable barrier that prevents further destruction. Aluminum cans
also have a thin plastic layer to prevent reaction of the oxide with acid in the soft drink. Chromium, magnesium,
and nickel also form protective oxide films. Stainless steels are remarkably resistant to corrosion because they
usually contain a significant proportion of chromium, nickel, or both.

3. • In contrast to these metals, when iron corrodes, it forms a red-brown hydrated metal oxide
(Fe2O3⋅xH2OFe2O3⋅xH2O), commonly known as rust, that does not provide a tight protective film
(Figure 20.8.120.8.1). Instead, the rust continually flakes off to expose a fresh metal surface vulnerable to
reaction with oxygen and water. Because both oxygen and water are required for rust to form, an iron nail
immersed in deoxygenated water will not rust—even over a period of several weeks. Similarly, a nail
immersed in an organic solvent such as kerosene or mineral oil will not rust because of the absence of water
even if the solvent is saturated with oxygen.

5. • Figure 20.8.120.8.1: Rust, the Result of Corrosion of Metallic Iron. Iron is oxidized to Fe2+(aq) at an anodic
site on the surface of the iron, which is often an impurity or a lattice defect. Oxygen is reduced to water at a
different site on the surface of the iron, which acts as the cathode. Electrons are transferred from the anode to
the cathode through the electrically conductive metal. Water is a solvent for the Fe2+ that is produced initially
and acts as a salt bridge. Rust (Fe2O3•xH2O) is formed by the subsequent oxidation of Fe2+ by atmospheric
oxygen. (CC BY-NC-SA; anonymous)
• In the corrosion process, iron metal acts as the anode in a galvanic cell and is oxidized to Fe2+; oxygen is
reduced to water at the cathode.

6. • The relevant reactions are as follows:
 at cathode:
• O2(g)+4H+(aq)+4e−⟶2H2O(l)O2(g)+4H+(aq)+4e−⟶2H2O(l)
• with EoSRP=1.23V�����=1.23�.
 at anode:
• Fe(s)⟶Fe2+(aq)+2e−Fe(s)⟶Fe2+(aq)+2e−
• with EoSRP=−0.45V�����=−0.45�.
 overall:
• 2Fe(s)+O2(g)+4H+(aq)⟶2Fe2+(aq)+2H2O(l)(20.8.1)(20.8.1)2Fe(s)+O2(g)+4H+(aq)⟶2Fe2+(aq)+2
H2O(l)
• with Eocell=1.68V������=1.68�.
• The Fe2+Fe2+ ions produced in the initial reaction are then oxidized by atmospheric oxygen to produce the
insoluble hydrated oxide containing Fe3+Fe3+, as represented in the following equation:
• 4Fe2+(aq)+O2(g)+(2+4x)H2O→2Fe2O3⋅xH2O+4H+(aq)(20.8.2)(20.8.2)4Fe2+(aq)+O2(g)+(2+4x)H2O
→2Fe2O3⋅xH2O+4H+(aq)
• The sign and magnitude of Eocell������ for the corrosion process

7. • (Equation 20.8.120.8.1) indicate that there is a strong driving force for the oxidation of iron by O2 under
standard conditions (1 M H+). Under neutral conditions, the driving force is somewhat less but still
appreciable (E = 1.25 V at pH 7.0). Normally, the reaction of atmospheric CO2 with water to form H+ and
HCO3
− provides a low enough pH to enhance the reaction rate, as does acid rain. Automobile manufacturers
spend a great deal of time and money developing paints that adhere tightly to the car’s metal surface to
prevent oxygenated water, acid, and salt from coming into contact with the underlying metal. Unfortunately,
even the best paint is subject to scratching or denting, and the electrochemical nature of the corrosion process
means that two scratches relatively remote from each other can operate together as anode and cathode, leading
to sudden mechanical failure (Figure 20.8.220.8.2).

9. • Figure 20.8.220.8.2: Small Scratches in a Protective Paint Coating Can Lead to the Rapid Corrosion of Iron.
Holes in a protective coating allow oxygen to be reduced at the surface with the greater exposure to air (the
cathode), while metallic iron is oxidized to Fe2+(aq) at the less exposed site (the anode). Rust is formed when
Fe2+(aq) diffuses to a location where it can react with atmospheric oxygen, which is often remote from the
anode. The electrochemical interaction between cathodic and anodic sites can cause a large pit to form under a
painted surface, eventually resulting in sudden failure with little visible warning that corrosion has occurred.

10. • Summary
• Corrosion is a galvanic process that can be prevented using cathodic protection. The deterioration of metals
through oxidation is a galvanic process called corrosion. Protective coatings consist of a second metal that is
more difficult to oxidize than the metal being protected. Alternatively, a more easily oxidized metal can be
applied to a metal surface, thus providing cathodic protection of the surface. A thin layer of zinc protects
galvanized steel. Sacrificial electrodes can also be attached to an object to protect it.

12. • when describing about corrosion cell (electrochemical cells), where a complete circuit is formed
when (i) : there is electrical contact between anode and cathode via a metallic path (wire) and
(ii): ions can flow through the electrolyte. However, corrosion cell can also occur in a single
(isolated) metal. For example, household gates or fences are often built with a single metal
material (mild steel) and overtime, they will corrode although without coupled with another
metal.

13. Corrosion Cell - Single Metal
• Corrosion Cell - Single Metal
• Let us begin with comparison of diagram between the 2 corrosion cells. They are shown below:
A= anode, C=cathode, E=electrolyte, e=electron, MP=metallic path

14. • As can be seen, they are not that different. Corrosion cell occurring within a single metal can arise due to
many factors. Do take note that, in reality there is no single material in this world that is perfectly uniform in
characteristics. The major factors that contribute to this phenomenon can include:
1. Differences in microstructure of the material (i.e- differences in grain boundary orientation, many phases
within the material)
2. Foreign material within the material (i.e - inclusions like oxides, sulfides which promotes anodic/cathodic
site)
3. Differential aeration (i.e- lower oxygen concentration promotes anodic site [such as metal traps within soil])
4. Heat effect (i.e - welding job, metal heat treatments)
5. Mechanical work (i.e-strained areas tend to become anodic site)

15. • Let's take for example, a piece of iron that is left outside and exposed to moisture. Due to
exposure with moisture (i.e-droplet of water), the iron surface that is in contact with water
droplet becomes anode while the iron at the edge of the water droplet becomes cathode. Electron
will move from the anodic site to the cathodic site within the metal itself. The reactions for this
phenomenon are:
• Oxidation: Fe --> Fe(2+) + 2e(-)
• Reduction: O(2) + 2H(2)O + 4e(-) --> 4OH(-) [Just outside water droplet]
• The hydroxide ions [4OH(-)] can move towards the water droplet and react with Fe(2+) ions to
form iron (II) hydroxide:
• Iron (II) hydroxide: Fe(2+) + 2OH(-) --> Fe[(OH)(2)] The iron hydroxide is then oxidized to
form rust.

17. • 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 16.8.116.8.1). 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 16.8.116.8.1:
Corrosion is a two-step
process

18. • Figure 16.8.116.8.1: 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. (CC BY 3.0 Unported; Stephen
Lower)
• A corrosion system can be regarded as a short-circuited electrochemical cell in which the anodic process is
something like
• Fe(s)→Fe2+(aq)+2e−(16.8.1)(16.8.1)Fe(s)→Fe2+(aq)+2e−
• and the cathodic steps may invove the reduction of oxygen gas
• O2+2H2O+4e−→4OH−(16.8.2)(16.8.2)O2+2H2O+4e−→4OH−
• or the reduction of protons
• H++e−→12H2(g)(16.8.3)(16.8.3)H++e−→12H2(g)
• or the reduction of a metal ion
• M2++2e–→M(s)(16.8.4)(16.8.4)M2++2e–→M(s)
• where MM is a metal.

19. • Which parts of the metal serve as anodes and cathodes can depend on many factors, as can be seen from the
irregular corrosion patterns that are commonly observed. Atoms in regions that have undergone stress, as
might be produced by forming or machining, often tend to have higher free energies, and thus tend to become
anodic.
Figure 16.8.216.8.2: Schematic diagram of corrosion cells on
iron. (CC BY-NSA-NC; Anonymous by request)

20. If one part of a metallic object is protected from the atmosphere so that there is insufficient O2O2 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 16.8.316.8.3: Pitting corrosion Most metals are covered with a thin oxide film which inhibits anodic
dissolution

21. • 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. (CC BY 3.0
Unported; Stephen Lower)
• 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 O2O2 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
• Anyone who has owned an older car has seen corrosion occur at joints between body parts and under paint
films. You will also have noticed that once corrosion starts, it tends to feed on itself. One reason for this is that
one of the products of the O2 reduction reaction is hydroxide ion. The high pH produced in these cathodic
regions tends to destroy the protective oxide film, and may even soften or weaken paint films, so that these sites
can become anodic. The greater supply of electrons promotes more intense cathodic action, which spawns even
more anodic sites, and so on.

22. Figure 16.8.416.8.4: 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).

23. • A very common cause of corrosion is having two dissimilar metals in contact, as might occur near a fastener
or at a weld joint. Moisture collects at the junction point, acting as an electrolyte and forming a cell in which the
two metals serve as electrodes. Moisture and conductive salts on the outside surfaces provide an external
conductive path, effectively short-circuiting the cell and producing very rapid corrosion; this is why cars rust out
so quickly in places where salt is placed on roads to melt ice.
• 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 Cu2+Cu2+ 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 Cr3+Cr3+ salts which precipitate as greenish stains.