There are three conditions that must exist for galvanic corrosion
- First there must be two electrochemically dissimilar metals
- Second, there must be an electrically conductive path between
the two metals.
- And third, there must be a conductive path for the metal ions
to move from the more anodic metal to the more cathodic metal.
If any one of these three conditions does not exist, galvanic
corrosion will not occur.
Corrosion as an Electrochemical
Schematic representation of current flow
(conventional current direction) in a simple corrosion cell.
, M. Tullmin.
A piece of bare iron left outside where it is exposed
to moisture will rust quickly. It will do so even more quickly
if the moisture is salt water. The corrosion rate is enhanced
by an electrochemical process in which a water droplet becomes
a voltaic cell in contact with the metal, oxidizing the
Considering the sketch of a water droplet (after Ebbing),
the oxidizing iron supplies electrons at the edge of the
droplet to reduce oxygen from the air. The iron surface
inside the droplet acts as the anode for the process.
Electrochemical corrosion involves two half-cell reactions;
an oxidation reaction at the anode and a reduction reaction
at the cathode. For iron corroding in water with a near
neutral pH, these half cell reactions can be represented
Fe(s) -> + Fe2+(aq) + 2e-
The electrons can move through the metallic iron to the
outside of the droplet where
O2(g) + 2H2O(l) + 4e-
Within the droplet, the hydroxide ions can move inward
to react with the iron(II) ions moving from the oxidation
region. Iron(II) hydroxide is precipitated.
Fe2+(aq) + 2OH-(aq) -> Fe(OH)2(s)
Rust is then quickly produced by the oxidation of the precipitate.
4Fe(OH)2(s) + O2(g) -> 2Fe2O3
H2O(s) + 2H2O(l)
The rusting of unprotected iron in the presence of air
and water is then inevitable because it is driven by an
electrochemical process. However, other electrochemical
processes can offer some protection against corrosion. For
magnesium rods can be used to protect underground steel
pipes by a process called cathodic protection.
In electrochemical corrosion, a galvanic cell is created when
two different metals, or different areas on the same metal, are
coupled by means of an electrical or ion-conducting electrolyte.
The result is an electrochemical reaction. In essence, electrochemical
corrosion is reserved for those processes where a current flows
between anodic and cathodic areas situated at different parts
of a metallic surface or between two different metals of the same
or different material.
The driving force for corrosion is a potential difference between
the different materials. The bimetallic driving force was discovered
in the late part of the eighteenth century by Luigi Galvani in
a series of experiments with the exposed muscles and nerves of
a frog that contracted when connected to a bimetallic conductor.
The principle was later put into a practical application by Alessandro
Volta who built, in 1800, the first electrical cell, or battery:
a series of metal disks of two kinds, separated by cardboard disks
soaked with acid or salt solutions. This is the basis of all modern
wet-cell batteries, and it was a tremendously important scientific
discovery, because it was the first method found for the generation
of a sustained electrical current.
Cells, General Chemistry, Richard A. Paselk, Department
of Chemistry, Humboldt State University.
|To allow the flow of electrons to continue
there must be a connection between the solutions in the beakers
so that the charges can be neutralized with a counter flow
of ions. The connecting ionic fluid is referred to as a salt-bridge.
Other arrangements are possible such as semi-permeable membranes
etc. Such an arrangement is called a Galvanic cell.
Galvanic cells (or Voltaic cells) are cells in which
the overall redox reactions occur spontaneously (equilibrium
favors products) as written. They can serve as a source
of electric power (as a battery).
Anodes vs. Cathodes.
Anode: the electrode at which the oxidation half reaction
occurs - electrons are taken up from the reaction at this
electrode. Thus at the zinc electrode Zn(II) ions would
be released, leaving two electrons behind for each ion.
Note that this electrode should have a negative charge!
(The chemistry is putting excess electrons into the electrode.)
Cathode: the electrode at which the reduction half reaction
takes place - electrons are released to the reaction at
this electrode. (Think cathode ray tube, CRT, shooting
electrons.) Note that the chemical reaction is taking
electrons from this electrode so it will be positive!
Thus a copper electrode would provide electrons to Ag(I)
ions to plate out silver metal.
The principle was also engineered into the useful protection
of metallic structures by Sir Humphry Davy and Michael Faraday
in the early part of the nineteenth century. The sacrificial corrosion
of one metal such as zinc, magnesium or aluminum is a widespread
method of cathodically protecting metallic structures.
In a bimetallic couple, the less noble material will become the
anode of this corrosion cell and tend to corrode at an accelerated
rate, compared with the uncoupled condition. The more noble material
will act as the cathode in the corrosion cell. Galvanic corrosion
can be one of the most common forms of corrosion as well as one
of the most destructive.
The relative nobility of a material can be predicted by measuring
its corrosion potential. The well known galvanic series lists
the relative nobility of certain materials in sea water. A small
anode/cathode area ratio is highly undesirable. In this case,
the galvanic current is concentrated onto a small anodic area.
Rapid thickness loss of the dissolving anode tends to occur under
these conditions. Galvanic corrosion problems should be solved
by designing to avoid these problems in the first place. Galvanic
corrosion cells can be set up on the macroscopic level or on the
microscopic level. On the microstructural level, different phases
or other microstructural features can be subject to galvanic currents.
Corrosion Cells, Corrosion Doctors.
Metallic corrosion is an electrochemical process consisting of
at least two partial reactions that occur on the metal surface.
At anodic areas, metal dissolution or corrosion occurs. At cathodic
areas, reactions such as the reduction of oxygen, or the evolution
of hydrogen occur. Electronic charge moves from the anodic site
through the metal to the cathodic site where electrons take part
in the cathodic reactions. Movement of ionic charge through the
aqueous phase between the two sites completes the electrochemical
circuit and corrosion proceeds. Corrosion may be prevented or
controlled by modification of the corrosive environment....
The environment for many structures provides conditions favoring
formation of natural corrosion cells. The metal or metals of a
structure serve as anode, cathode, and the necessary metallic
conductor between the two. Water, either as such or as moisture
in soil, provides the electrolyte required to complete the cell
circuit. Such cells develop their driving force or electrical
potential from differing conditions at the interfaces between
metal and electrolyte of the anode and cathode. These differences
fall into three categories:
- Dissimilar metals comprising the anode and cathode
- Inhomogeneity of a single metal, which causes one area to
be anodic to another area
- Inhomogeneity of the electrolyte
- Corrosion Cell
The following examples illustrate situations in which the essential
requirements of a complete cell are satisfied in a structure:
- Iron will be anodic to copper ground mats or to brass bolts
or other brass parts
- An iron plate having some mill scale present may rust because
the iron is anodic to the mill scale
- An apparently homogeneous iron plate may rust because tiny
areas of the surface contain impurities or grain stresses which
cause them to be anodic to other areas of the surface
- Weld areas of a welded pipe may rust because the weld metal
is of different composition, may contain impurities, or may
cause stress which make it anodic to nearby metal areas
- Corrosion may be observed on the bottom of a pipeline while
the top remains nearly undamaged. This may be attributable to
higher oxygen concentration in the soil moisture (electrolyte)
at the top of the pipe, leaving the bottom anodic. The soil
being undisturbed at the bottom of the pipe provides a lower
oxygen content and a lower resistance to current flow than is
present in the backfill covering the top of the pipe.
- Exposed iron areas in contact with concrete. Encased or embedded
iron may rust because the concrete creates a different and special
electrolytic environment which causes the exposed iron to become
anodic to the embedded iron.