How does rust come about?
As rust This is the name given to the corrosion product that arises from iron or steel through oxidation with oxygen in the presence of water. Rust is generally made up of iron (II) oxide, iron (III) oxide and water of crystallization, according to the empirical formula:
(x, y, z positive ratios).
The oxidation products that form on the surface of iron at high temperatures are called scale and consist of iron oxides of different oxidation states. Particularly when forging red-hot iron, hammer blows flake off thin gray-black iron oxide layers from the surface, which are known as Hammer blow are designated.
In the case of other metals, such as zinc, chromium, aluminum or nickel, some of which are also less noble than iron, only the topmost atomic layers oxidize to a barely visible oxide layer that shields the metal below from further reaction with oxygen (see also Passivation).
In the case of iron, however, the corrosion at the rust / material interface does not come to a standstill, because the electrical conductivity of the (moist) rust that has already formed and its oxygen permeability promote further corrosion at the rust / material boundary.
Rust forms loose structures of low strength. The oxidation causes an increase in mass and volume. The latter leads to tension and flaking of the rust layer (see figure on the right). Reinforcing steel does not rust if it is well encapsulated and embedded in the concrete. The alkaline environment of concrete offers additional protection. If, however, water and air gain access to the steel, the concrete bursts open as a result of the increase in the volume of the rust and the decay is accelerated.
Electrochemical model of rust formation
A drop of water (blue) lies on an iron surface (gray), surrounded by air (white), see diagram on the right. According to the voltage series of the elements, positively charged iron atoms diffuse into the aqueous environment, the electrons remain in the metal and charge it negatively (see (1) in the schematic drawing):
At first one could assume that the iron reacts with the hydrogen of the water according to (Hydrogen corrosion):
- Dissociation of the water into protons (autoprotolysis)
The negative charge of the metal and the boundary layer of positively charged iron ions prevent rapid reaction with protons.
If oxygen is present, it can take over the transport of the electrons. In the diagram above, oxygen diffuses from the outside into the water droplets. The difference in concentration in the water droplet creates a potential difference between (2) and (3). The anodic area (2) and cathodic area (3) form a galvanic cell with the water as the electrolyte, which is referred to as a corrosion or local element due to the short circuit between (2) and (3).
In (3) the electrons react with water and oxygen to form hydroxide ions:
The hydroxide ions form iron (III) hydroxide (4) with the iron ions. With the participation of oxygen and water, poorly soluble iron (III) oxide hydroxide precipitates, which is deposited on the iron surface at (5) (Oxygen corrosion):
Simplified, the sum equation from (A) and (B) reads:
If iron comes into contact with another metal, a local element is also created at the contact point, which leads to corrosion of the less noble metal.
Heavily rusted metals can be brushed to remove rust. Rust converters are often used in repairs, e.g. B. used by cars.
Light rust can also be washed off with a weak acid. Dilute phosphoric acid such as that found in cola is suitable, for example. So that the acid does not attack the metal, it must be rinsed off with plenty of water. The metal must be dried thoroughly and protected from further corrosion.
Three strategies for corrosion protection can be derived from the model:
Keep away from oxygen
Example: Iron heating pipes do not rust on the inside if the water is routed in a closed system without air access. In addition, the solubility of oxygen decreases as the temperature of the water rises.
However, these reactions can also be prevented by various protective measures. One example of this is passivation - the coating with such less noble metals that form a stable oxide layer.
If metals are coated with less noble metals, this is called electroplating, galvanizing or chrome plating. To prevent the ingress of water, metals can be painted or given a plastic coating.
Keeping away from moisture that acts as an electrolyte
Example 1: In countries with low humidity, there is practically no rust damage to cars.
Example 2: Protective layers of grease, paint or metal coatings shield iron from the environment (hot-dip galvanizing, tinplate), see pictures on the left and right.
Example 3: An iron alloy with a chromium content of more than 12% (=Stainless steel) is protected from oxidation by the chromium oxide layer.
Reduction of the potential difference in local elements
Example 1: Hot-dip galvanizing provides long-term protection for iron against rust. If the coating is damaged, zinc and iron form a local element when exposed to water (similar to a battery). Zinc, the less noble metal, corrodes and protects iron from oxidation.
In the case of a coating with a more noble metal (e.g. tin for tinplate), the opposite occurs. The iron rusts, possibly covered by the protective layer (see picture of the beverage can, hash browns). The presence of a more noble metal even promotes oxidation. The local element made of iron and the more noble metal prevents the protective negative charging of the iron (see above).
Example 2: Iron pipes are electrically connected to a so-called sacrificial anode made of a less noble metal. As in the first example, iron is protected at the expense of the sacrificial anode, provided that both are in contact via an electrolyte, for example moist soil.
Example 3: Instead of a sacrificial anode, an electrically conductive electrode (e.g. graphite) also protects if it is kept at a positive potential relative to the iron via an external DC voltage source.
- Werner Schatt: Introduction to materials science. German publishing house for basic industry, Leipzig 1991. ISBN 3-342-00521-1
- Günter Schulze, Hans-Jürgen Bargel: Materials science. Schroedel, Hanover 1978, Springer, Berlin 72000. ISBN 3-540-66855-1
Categories: Chemical Reaction | Corrosion | iron
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