Galvanic Corrosion: Mechanisms, Examples, Protection

The galvanic or electrochemical corrosion is a process whereby a metal or alloy degrades more precipitously compared to the conventional oxidation. It can be said that it is an accelerated oxidation, and even, intentionally caused; as happens in cells or batteries.

This takes place under a number of conditions. First, there must be an active metal, called the anode. Also, and secondly, there must be a low-reactive noble metal called the cathode. The third and fourth conditions are the presence of a medium where electrons propagate, such as water, and ionic species or electrolytes.

Rusty iron crown. Source: Pixnio.

Galvanic corrosion is observable especially in marine environments or on the shores of beaches. The air currents raise masses of water vapor, which in turn, carry some ions; the latter end up adhering to a thin layer of water or drops that rest on the metal surface.

These conditions of humidity and salinity favor the corrosion of the metal. That is, an iron crown like the one in the image above will rust more quickly if it is exposed to the vicinity of the sea.

The ease that a metal will have to oxidize compared to another can be measured quantitatively through its reduction potentials; Tables with these potentials abound in chemistry books. The more negative you are, the more likely you are to rust.

Likewise, if this metal is in the presence of another with a very positive reduction potential, thus having a large ΔE, the oxidation of the reactive metal will be more aggressive. Other factors, such as pH, ionic strength, humidity, the presence of oxygen, and the relationship between the areas of the metal that is oxidized and that that is reduced, are also important.

Article index

  • one

    Mechanisms

    • 1.1

      Concepts and reactions

    • 1.2

      Depolarizers

    • 1.3

      Iron corrosion

  • two

    Examples

    • 2.1

      Anodic indices

  • 3

    Electrochemical corrosion protection

    • 3.1

      Sacrificial coatings

    • 3.2

      Noble coatings

  • 4

    Experiment for children

    • 4.1

      Iron plate in dissolution of copper salts

    • 4.2

      Silver oxide cleaning

  • 5

    References

Mechanisms

Concepts and reactions

Before addressing the mechanisms behind galvanic corrosion, certain concepts should be clarified.

In a redox reaction, one species loses electrons (oxidizes) while another gains them (reduces). The electrode on which oxidation occurs is called the anode; and on which the reduction occurs, cathode (in English the redcat mnemonic rule is often used to remember it).

Thus, for an electrode (a piece, screw, etc.) of a metal M, if it oxidizes it is said to be the anode:

M => M n + + ne

The number of electrons released will be equal to the magnitude of the positive charge of the resulting cation M n + .

Then another electrode or metal R (both metals must be in contact in some way), receives the released electrons; but this does not undergo a chemical reaction if it gains electrons, since it would only be conducting them (electric current).

Therefore, there must be another species in solution that can formally accept these electrons; as easily reduced metal ions, for example:

R n + + ne => R

That is, a layer of metal R would form and the electrode would therefore become heavier; while the metal M would lose mass due to the dissolution of its atoms.

Depolarizers

What if there were no metal cations that could be reduced easily enough? In that case, other species present in the medium will take the electrons: the depolarizers. These are closely related to pH: O 2 , H + , OH and H 2 O.

Oxygen and water gain electrons in a reaction expressed by the following chemical equation:

O 2 + 2H 2 O + 4e => 4OH

While the H + ions are transformed into H 2 :

2H + + 2e => H 2

That is, the OH and H 2 species are common products of galvanic or electrochemical corrosion.

Even if the metal R does not participate in any reaction, the fact that it is more noble than M promotes its oxidation; and consequently, there will be a higher production of OH ions or hydrogen gas. Because, after all, it is the difference between the reduction potentials, ΔE, one of the main drivers of these processes.

Iron corrosion

Corrosion mechanism for iron. Source: Wikipedia.

After the above clarifications, the example of iron corrosion can be addressed (top image). Suppose there is a thin layer of water in which oxygen dissolves. Without the presence of other metals, it will be the depolarizers who will set the guidelines for the reaction.

Thus, iron will lose some atoms from its surface to dissolve in water as Fe 2+ cations :

Fe => Fe 2+ + 2e

The two electrons will travel through the piece of iron because it is a good conductor of electricity. So it is known where the oxidation or the anode site started; but not where the reduction will proceed or the location of the cathodic site. The cathode site can be anywhere; and the larger its possible area, the worse the metal will corrode.

Suppose the electrons reach a point as shown in the image above. There both oxygen and water undergo the reaction already described, by which OH is released . These OH anions can react with Fe 2+ to form Fe (OH) 2 , which precipitates and undergoes subsequent oxidations that eventually transform it into rust.

Meanwhile, the anode site is cracking more and more.

Examples

In everyday life the examples of galvanic corrosion are numerous. We do not have to refer to the iron crown: any artifact made of metals can undergo the same process in the presence of humid and saline environments.

In addition to the beach, winter can also provide ideal conditions for corrosion; for example, when shoveling salts into snow on the road to prevent cars from skidding.

From the physical point of view, moisture can be retained in the welded joints of two metals, being active sites of corrosion. This is because both metals behave like two electrodes, with the more reactive one losing its electrons.

If the production of OH ions is considerable, it can even corrode the paint of the car or the device in question.

Anodic indices

One can construct his own examples of galvanic corrosion by making use of the reduction potential tables. However, the anodic index table (simplified per se) will be chosen to illustrate this point.

Anodic indices for different metals or alloys. Source: Wikipedia.

Suppose for example that we wanted to build an electrochemical cell. The metals at the top of the anodic index table are more cathodic; that is, they are easily reduced and it will therefore be difficult to have them in solution. While the metals that are at the bottom are more anodic or reactive, and they corrode easily.

If we choose gold and beryllium, both metals could not be together for long, since beryllium would oxidize extremely quickly.

And if, on the other hand, we have a solution of Ag + ions and we submerge an aluminum bar in it, it will dissolve at the same time as metallic silver particles precipitate. If this bar were connected to a graphite electrode, electrons would travel to it to electrochemically deposit silver on it as a silver film.

And if instead of the aluminum bar it were made of copper, the solution would turn bluish due to the presence of the Cu 2+ ions in the water.

Electrochemical corrosion protection

Sacrificial coatings

Suppose you want to protect a zinc sheet from corrosion in the presence of other metals. The simplest option would be to add magnesium, which would coat the zinc so that, once oxidized, the electrons released from the magnesium would reduce the Zn 2+ cations back.

However, the MgO film on the zinc would end up cracking sooner rather than later, providing high current density anode sites; that is, the corrosion of the zinc would accelerate sharply at just those points.

This electrochemical corrosion protection technique is known as the use of sacrificial coatings. The best known is zinc, used in the famous technique called galvanizing. In them, the metal M, especially iron, is coated with zinc (Fe / Zn).

Again, the zinc oxidizes and its oxide serves to cover the iron and transmit electrons to it that reduce the Fe 2+ that can be formed.

Noble coatings

Suppose again that you want to protect the same zinc sheet, but now you will use chromium instead of magnesium. Chromium is more noble (more cathodic, see table of anodic indexes) than zinc, and therefore works as a noble coating.

The problem with this type of coating is that once it cracks, it will further promote and accelerate the oxidation of the metal underneath; in this case, the zinc would corrode even more than being coated with magnesium.

And finally, there are other coatings that consist of paints, plastics, antioxidants, fats, resins, etc.

Experiment for children

Iron plate in dissolution of copper salts

A simple experiment can be devised from the same table of anode indices. Dissolving a reasonable amount (less than 10 grams) of CuSO 4 · 5H 2 O in water, a child is asked to dip into a polished iron plate. A photo is taken and the process is allowed to unfold for a couple of weeks.

The solution is initially bluish, but will begin to fade while the iron plate turns a coppery color. This is due to the fact that copper is more noble than iron, and therefore its Cu 2+ cations will be reduced to metallic copper from the ions given by the oxidation of iron:

Fe => Fe 2+ + 2e

Cu 2+ + 2e => Cu

Silver oxide cleaning

Silver objects blacken over time, especially if they are in contact with a source of sulfur compounds. Its rust can be removed by immersing the object in a tub of water with baking soda and aluminum foil. The bicarbonate provides the electrolytes that will facilitate the transport of electrons between the object and the aluminum.

As a result, the child will appreciate that the object loses its black spots and will glow with its characteristic silver color; while the aluminum foil will corrode to disappear.

References

  1. Shiver & Atkins. (2008). Inorganic chemistry. (Fourth edition). Mc Graw Hill.
  2. Whitten, Davis, Peck & Stanley. (2008). Chemistry. (8th ed.). CENGAGE Learning.
  3. Wikipedia. (2019). Galvanic corrosion. Recovered from: en.wikipedia.org
  4. Stephen Lower. (June 16, 2019). Electrochemical Corrosion. Chemistry LibreTexts. Recovered from: chem.libretexts.org
  5. The Open University. (2018). 2.4 Corrosion processes: galvanic corrosion. Recovered from: open.edu
  6. Customer Technical Service Brush Wellman Inc. (sf). A Guide to Galvanic Corrosion. Brush Wellman Engineered Materials.
  7. Giorgio Carboni. (1998). Experiments in electrochemistry. Recovered from: funsci.com

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