Why does iron rust?

in HeartSTEM3 years ago
Greetings dear friends #HeartSTEM.

As we know, iron rusting is a very important issue for today's technology, since many construction elements are made of this versatile and resistant metal. Understanding why iron rusts is the first step to try to prevent this process, which not only affects the appearance of the metal elements, but also deteriorates them, so it also represents an important economic impact.


Oxidation is the cause of the deterioration of iron elements. Source: pixabay.com.

So understanding the process of oxidation and corrosion of iron is very important for anyone. That is why in this article I present an explanation of why iron oxidizes showing how this phenomenon is revealed with a simple experiment that everyone can perform, besides allowing us to approach in a simple way the study of oxidation-reduction reactions.

So, why does iron oxidize?

We must begin by saying that, with a few exceptions such as gold, the oxidation of metals is inevitable, since they do not exist in nature in their metallic state, but in combination with some other element, forming minerals such as oxides, carbonates or sulfates.

When steel is produced in the metallurgical industry, basically large amounts of energy are used to free the iron from the rest of the elements with which it is combined. So, if it requires a large amount of energy to bring iron to its metallic state, it is because this is not its most stable form, which is why we can see oxidation as a return of metals to their natural state. In the case of iron, its most common natural form is hematite, which is an oxide with the formula Fe2O3[1].

So the driving force that causes a metal to oxidize is its tendency to return to its combined state, oxidized in the case of iron; but to return to this state requires a chemical reaction. This reaction is known as redox, or oxidation-reduction, a type of chemical reaction where the reactants exchange electrons so that the oxidation state of each increases or decreases.

During an oxidation-reduction reaction, chemical changes occur in the species involved, on the one hand the species being oxidized suffers a loss of electrons, increasing its oxidation state. These electrons flow to another species that acts as an oxidizing agent and undergoes a reduction of its oxidation state[2].

The elements in their fundamental state, or molecular in the case of gases, have an oxidation state of 0, and when oxidized this number increases. For example, iron when oxidized in the presence of oxygen goes from its oxidation state 0 to +2, due to the transfer of two electrons to oxygen which changes its oxidation state from 0 to -2.

In a general way we can represent the reaction as:

ec1.jpg

Let's see how oxidation occurs

The procedure is very simple, we will only need a couple of disposable plastic plates, a couple of common iron nails, water and a little salt. Before starting it is convenient to clean and sand the surface of both nails to remove any rust stain they may have, and thus both nails will start with a similar appearance.

We will take the two plates and we will identify, to the first one we will add common tap water and to the second one water with a little table salt; and in each plate we are going to place an iron nail (the water in each plate should be enough to leave part of the nail surface exposed).

plato clavo.jpg
Beginning of the experiment. Source: @emiliomoron.

In plate n°1, where we have the nail immersed in common water, which contains some salts and dissolved oxygen, after a few hours we begin to observe that the water in the vicinity of the nail is turning orange, which gives us indications that the oxidation reaction has started.

clavo 2.jpg
After a few hours the nail already shows signs of oxidation. Source: @emiliomoron.

This reaction takes place basically because in the nail there are some zones that function as small electrochemical cells, on the one hand, in some zones the oxidation reaction takes place (anodic zones) and in other zones the reduction reaction (cathodic zones).

In the anodic zones the oxidation of the iron takes place according to the half-reaction:

ec2.jpg

And in the cathodic zones, oxygen reduction takes place:

ec3.jpg

The following figure shows a representation of the process.

Imagen1 zones ox.jpg
Representation of the reactions of the iron oxidation process. Source: @emiliomoron.

The anodic and cathodic zones are not necessarily located in opposite places, in general the cathodic zones are usually located where the oxygen concentration is higher, and the anodic zones will preferably be located where the oxygen concentration is lower, that is on the submerged surface, producing the oxidation of the nail in this part, dissolving the iron ions in the solution.

In water, with a pH close to neutral, and in the presence of oxygen, the overall nail oxidation reaction can be represented as follows:

ec4.jpg

The iron ions Fe+2 then react with the dissolved oxygen in the water to produce a reddish orange precipitate of hydrated ferric oxide.

In the plate n°1 in which the nail is immersed in common water the process is relatively slow; however, when we compare with the process carried out in the plate n°2, where the nail is immersed in water and the dissolved salt, we can observe that the process is much faster. In the following images we can compare how the solution is colored with the ferric oxide as the experiment progresses.

Imagen2 comp clavos.jpg
Oxidation progress in both plates at: a) 4 hours, b) 12 hours and c) 24 hours. Source: @emiliomoron.

This is because oxidation is an electrochemical process, so it requires charge transfer; and as the dissolved salt increases the conductivity of the medium it favors charge transport in the solution, accelerating the process. That is why in saline environments iron corrosion is faster.


Well friends, I hope to have shared information of your interest about the oxidation process, which as we have seen, the oxidation of an iron element exposed to moisture and air does not take long to be noticed, allowing to perform this experience by which it is easy to demonstrate this process.

Thanks so much for stopping by to read the post, see you next time!


References

  1. Wikipedia.com. Óxidos de hierro.
  2. Chang, R. (2002). Chemistry. 7th edition. McGraw-Hill.
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Good experiment and nice publication.

Thank you very much!

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