Rust

Rust corrosion is defined as the product formed of iron or steel by oxidation with oxygen in the presence of water. Rust is porous and does not protect against further degradation, other than the oxide layer of many metallic materials as chromium, aluminum or zinc. Based on these properties, the metals in the groups -ferrous metals ( rust ) and non-ferrous metals ( will not rust ) distinguished. The weathering of ferrous materials in air and water to rust caused worldwide annual losses in the billions.

• 4.1 Keep away from oxygen
• 4.2 Keep away from moisture
• 4.3 reduction of the potential difference in local elements

Survey

Chemically rust iron (III ) oxide and water of crystallization is generally composed of iron (II ) oxide, together. Molecular formula:

(x, y, z are positive numbers ratio )

Rust is therefore a water-containing oxide of iron, a chemical compound which is one of the oxides, and addition of water and contains hydroxide ion ( oxide hydrate ). It is produced by the oxidation of iron, without higher temperatures would be required. Rust is thus similar to the connection Braunstein ( hydrous manganese dioxide), which is also to be regarded as hydrated oxide of a transition metal.

Rust forms loose structure of low strength. The oxidation causes an increase in mass and volume. The latter leads to tension and to flaking of the rust layer (see illustrations).

For corrosion protection of ferrous materials are coated with protective layers, fitted with sacrificial anodes or afterwards with phosphoric acid removes rust ( corrosion protection).

Reinforcing steels do not rust if they are well encapsulated embedded in the concrete. Additional protection is provided, the alkaline environment of concrete. But if getting water and air access to the steel, the concrete bursts due to the increase in volume of rust and decay is accelerated because it have water and air even better access.

Electrochemical model of the formation of rust

The formation of rust (corrosion) of iron begins

• By the attack of an acid ( acid corrosion) or
• Of oxygen and water (oxygen corrosion) on the metal surface.

Acid corrosion

In the case of an acid corrosion ( hydrogen corrosion), the protons ( hydrogen ions ) of the acid to escape the metal electrons: iron reacts with hydrogen ions in the water (at A) to ferrous cations:

The hydrogen ions (oxidizing agent ) react to this hydrogen gas ( redox reaction ), as they absorb the electrons of the metal ( reduction of the oxidizing agent ). The reaction scheme of the overall reaction is therefore:

Oxygen corrosion

In the case of oxygen corrosion ( weathering of iron to rust) oxygen acts as the oxidant: He picks up electrons.

In the schematic representation of rusting (see picture) is located on an iron surface (gray) a drop of water (blue), surrounded by air (white). According to the series of the elements to diffuse the positively charged iron ions in the aqueous environment, the electrons remain in the metal charge it negatively, see (1) in the diagram.

Neutral water contains 10-7 mol / L hydrogen ions ( autoionization ):

The negative charge of the metal and the boundary layer of positively charged iron ions on the iron surface generally prevent a rapid reaction with protons: oxygen - and air-free water does not attack the iron metal.

However, if oxygen is present, it takes over the transport of electrons. It diffuses from the outside in water drops (see diagram ). The difference in concentration in the water drops now generates a potential difference between (2) and (3). The anode region (2 ) and the cathodic zone (3) form with the water as electrolyte, a galvanic cell, a redox reaction is running.

The electrons react with water and oxygen to form hydroxide ions, see in ( 3) in the diagram:

The hydroxide ion form, with the iron ion, iron ( II) hydroxide (4).

Iron (II ) hydroxide is olive green to gray-green, and is implemented in the presence of water and air to ferric ions. Together with the hydroxide ions formed at the second redox russet iron ( III) hydroxide:

Simplifies the overall reaction scheme is thus:

Due water delivery thereof to form sparingly soluble iron (III ) oxide hydroxide, which deposits on the iron surface in (5):

In addition, the following actions take place:

The mixture initially formed from iron ( II ) hydroxide and iron ( III ) hydroxide is partially water discharge to a stable mixture of iron (II ) oxide, iron (III ) oxide and water of crystallization implemented, which is colloquially referred to as rust:

Accelerating factors in the formation of rust

When iron is in contact with another metal, is formed at the contact point a local cell, which leads to the corrosion of the less noble metal. The rusting process is also accelerated by the presence of salts, since they increase the conductivity of the water. The migration of ions in the water is important for the corrosion process, otherwise the circuit would be interrupted and the corrosion came very quickly to a halt (see salt bridge in a normal electrochemical cell).

Oxidation and corrosion processes, similar to the formation of rust

Anhydrous oxidation products that are formed at high temperatures on the surface of iron are called scale. They are unlike rust from water or hydroxide - free iron oxides of different oxidation states. Especially when forging hot iron burst by hitting them from the surface thin gray black iron oxide layers, which are referred to as a hammer blow.

With other metals such as zinc, chromium, aluminum or nickel, which are sometimes also less noble than iron, oxidizing only the uppermost atomic layers to a barely visible oxide film that shields the underlying metal from further reaction with oxygen (see passivation).

In the presence of air and water but also weathering and corrosion processes can occur, for example, to copper patina. However, when iron is the corrosion of the grid / material interface is not stopped, because the electrical conductivity of the already formed (wet) the grate and its oxygen permeability favor the further corrosion of the boundary grid / material.

At temperatures> 180 ° C are formed on surfaces of ferrous materials when exposed to water vapor at high temperature protective layers of magnetite ( Fe3O4). It is produced by reaction of metallic iron with water molecules to form hydrogen. For pipes in high-pressure boilers with locally very high heat load this reaction can proceed amplified and is sometimes one of the causes of pipe bursters. High pH values ​​of the water, particularly in the presence of alkali metal ions, in addition to accelerate this reaction.

Rust removal

Heavily rusted metals can be removed by brushing or grinding of rust. One of the most effective methods for the removal of rust is blasting with sand or similar materials that are free of silica. This method is used mainly in the art in front of a painting. There is not enough sand blast method, also the pneumatic needle scalers can be used. The complete removal of rust down to bare metal is one of the requirements that a corrosion resistant coating can be achieved.

Light rust can be washed with a weak acid. Good example is dilute phosphoric acid. Thus, the acid does not attack the metal, they must then be rinsed with plenty of water. The metal must be thoroughly dried and protected from further corrosion. Phosphoric acid is also used as a rust converter and is used in various blends for the repair of cars about.

In all of these methods to remove rust from the rust is removed, the rusty removal is lost.

Corrosion protection

From the model, three strategies for corrosion protection can be derived:

Keep away from oxygen

For example, heating pipes of iron do not rust internally, if the water is conducted in a closed system without access of air. In addition, the solubility of oxygen decreases with increasing heating of the water and reaches 111.6 ° C a minimum. Above this temperature, however, beyond the solubility for oxygen increases significantly again.

It is these reactions, but also by various other protective measures to prevent. One example is the passivation of: the coating with such noble metals that form a stable oxide layer. A metal may also be provided by electroplating, galvanizing or chrome plating with another metal as a protective layer against oxidation. Other protective coatings are non-diffusive and non-porous paints and coatings with plastics and spun concrete.

Keep away from moisture

Since water acts as an electrolyte in the reaction for rust development, keeping dry is a good counter-strategy. So there is practically example, in countries with low humidity, no rust damage to cars.

Another option are protective layers of grease, paint, chrome or metal supports that iron from the surrounding shield ( hot-dip galvanizing, tin plate ). Once this protective layer is destroyed, the Rostungsprozess begins to see pictures left and right.

Stainless steel is a ferrous alloy with a chromium content of more than 12 %, and is protected by the chromium oxide film from oxidation.

Reduction of the potential difference in local elements

Example 1: Hot dip galvanizing protects iron lasting protection against rust. If there is damage to the coating, forming zinc and iron in the presence of water, a local element (similar to a battery). Zinc corrodes as the less noble metal and keeps the iron from oxidation. In most zinc dust color, the zinc can be electrodeposited on the other hand do not work, since it is isolated from the binder. Only zinc dust paints with electrically conductive binder or zinc dust paints based on epoxy resin with a suitable pigment volume concentration (PVC), in which touch the zinc particles protect from corrosion.

With a coating of a more noble metal (for example tin tinplate ) enters the opposite case. The rusty iron, possibly obscured by the protective layer ( see picture of the can ). The presence of a noble metal even promotes oxidation. The local element of iron and the more noble metal prevents the protective negative charging of the iron ( see above).

Example 2: Iron tubes are electrically connected to a so-called sacrificial anode made ​​of a noble metal. As in the first example, iron is protected at the expense of the anode, if both are on an electrolyte, for example in damp soil contact.

Example 3: Instead of a sacrificial anode and an electrically conductive electrode protection (for example graphite), when held by an external DC power source to a positive potential relative to the iron. This is then called cathodic protection, which is used in pipelines and in bridge construction.