Corrosion is part of everyday life. You cross a bridge and notice the red flakes on the steel railing. The culprit? Corrosion. Even the wrought-iron of the Eiffel Tower does not escape it.
What is corrosion?
But what is corrosion? At first glance, it is dual. Corrosion is sometimes used on purpose for aesthetic reasons since its bright reddish orange color suits outdoor decoration perfectly. The flip side of the coin is that it is notorious for affecting the structural integrity of constructions and the functionality of e.g. electrical components.
A poet might describe corrosion as a destructive yet beautiful force of nature. However, we as scientific experts in the field would like to give you a more practical introduction into the phenomenon. After reading the next 388 words you will have a better understanding of both the concept of wet corrosion (in the presence of an electrolyte) and how to prevent it.
In nature metals, such as iron, are found in their most stable form. For most metals, except for the most noble metals, this is their ore (oxidized state). To extract a metal from its ore, energy is required, thereby bringing it to a higher energy state. Nature strives to the lowest energy state, so metals tend to oxidize again.
Wet corrosion: how it occurs
Wet corrosion is an electrochemical process. Taking the example of iron (Fe) in contact with oxygenated water: a number of Fe atoms will oxidize into Fe2+ leaving each time 2 electrons behind in the material until an equilibrium is reached (see equation [1]). The metal has acquired a certain potential. When electrons are removed from the metal or ions are moved away from the surface, the reaction is driven to the right (oxidation). When there are too many Fe2+ ions or too many electrons in the metal the reaction is driven to the left (reduction/metal deposition).
How electrons are consumed
Electrons can be consumed in different ways, examples of cathodic reactions consuming electrons are:
- In acid environment (aerated): O2 + 4H+ + 4e- → 2H2O
- In alkaline environment (aerated): O2 + 2H2O + 4e- → 4OH-
- In acid anaerobe environment : 2H+ + 2e- → H2(g)
- Other
This is e.g. the reason why iron oxidizes (rusts) in contact with oxygenated water, or why metals corrode in acid environments.
Corrosion can only take place when there is a closed electrical circuit (the corrosion cell) between an anode, where the oxidation takes place (i.e. rusting of steel) and the cathode where the reduction reaction takes place (consumption of electrons). Anode and cathode can be small areas of the same metal plate (due to differences in microstructure, internal stresses, …) or can be dissimilar metals in electrical contact with each other. An electrolyte is needed to close the electrical circuit by means of ion transport between anode and cathode.
The wet corrosion process: 4 connecting mechanisms
To summarize 4 connecting mechanisms are required in order to maintain the wet corrosion process:
- Anodic reaction/oxidation
- Electron transport
- Cathodic reaction/reduction reaction
- Ion transport/solubility
How to prevent corrosion
Strategies to prevent corrosion try to remove or block at least one of the 4 connecting mechanisms mentioned above. As soon as the electrical circuit is interrupted the corrosion process stops.
Examples of strategies are:
- Coatings or film forming inhibitors: by forming a barrier against the electrolyte anodic/cathodic reactions cannot take place
- Water treatment products: anodic or cathodic inhibitors constrict respectively the anodic or cathodic reaction
- Taking away the electrolyte: room climatization to avoid condensation, …
- Taking away the electrical contact between anode/cathode: galvanic separation of dissimilar metal components
Now you have had a first introduction to the concept of corrosion. Of course, corrosion is a fascinating matter, so we understand if you already have a hunger for more. If you want to learn more about corrosion, METALogic offers in depth courses online or on site that are catered to your specific questions.
