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Galvanic Corrosion

There are three conditions that must exist for galvanic corrosion to occur.

  1. First there must be two electrochemically dissimilar metals present.
  2. Second, there must be an electrically conductive path between the two metals.
  3. 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 Process
Schematic representation of current flow (conventional current direction) in a simple corrosion cell. Soucre: Basic Theory, 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 iron.

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 as:

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- -> 4OH-(aq)

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.

Galvanic 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.

Natural 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:

  1. Dissimilar metals comprising the anode and cathode
  2. Inhomogeneity of a single metal, which causes one area to be anodic to another area
  3. Inhomogeneity of the electrolyte
  4. Corrosion Cell

The following examples illustrate situations in which the essential requirements of a complete cell are satisfied in a structure:

  1. Iron will be anodic to copper ground mats or to brass bolts or other brass parts
  2. An iron plate having some mill scale present may rust because the iron is anodic to the mill scale
  3. 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
  4. 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
  5. 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.
  6. 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.