Metals Corrosion; The rate of corrosion interaction is related to several kinetic and thermodynamic factors, such as temperature, pH, fluid velocity and the concentration of the affecting medium.
The most important factors that are involved in corrosion reactions are:
The effect of temperature on metals corrosion
Increasing the temperature causes an increase in the amount and speed of corrosion. Even when the temperature of different parts of a certain part is different. Generally, the part that has a higher temperature becomes more anodic than other parts.
Temperature plays a dual role in oxygen corrosion, increasing temperature reduces oxygen solubility.
Related content: What is corrosion?
Impact of potential differences on metal corrosion
Due to the potential difference between the electrodes (for example, the two metals zinc and iron in brine), The most active metal is anode and it corrodes. This is the case when homogeneous and bonded metals are in a common electrolyte environment. In this reaction, the action of metals corrosion will not happen and the other metal will be protected.
Heat treatment often severely affects the behavior of metals corrosion and alloys.
The effect of “surface condition” on metals corrosion.
Surface conditions are another cause of corrosion of metals. The onset of reaction and the rate of corrosion vary greatly between polished and clean metal surfaces compared with rough surfaces or surfaces with surface films or other foreign materials.
The effect that speed can have on metals corrosion …
The destructive effects of mechanical abrasion alone are not discussed but are considered for the following reasons.
- Early corrosion films (which in most cases prevent corrosion from progressing) are lost due to wear.
- The surface of activated metal is re-exposed to subsequent corrosion.
Related article: What is cathodic protection?
So far, relatively little research has been done on the effect of radiation on metal corrosion. Radiation is one of the topics that should be given more attention.
However, experiments have been performed on damage caused by environments containing atomic radiation metals. These experiments the amount and intensity of corrosion of metals increases.
The presence of various impurities in the environment is one of the most important causes of metals corrosion. These impurities have various effects on the manner and amount of corrosion
The effect of metals corrosion usually increase over time. In some cases, there is a linear relationship between them. However, in some cases, the amount of corrosion may decrease over time. Which is often due to the nature of the jumps and deposits that form on the surface of metals.
The rate of material degradation in industrial parts and equipment intensifies under the following conditions:
- Conditions that are affected by tensile stresses and are also exposed to corrosive environments.
- When tensions are within (or above) their elastic limit.
Studies have shown that stress is an effective factor in chemical reactions and oxidation of materials. Hence it needs to be carefully considered.
Metallic properties can be exacerbated by metals corrosion .
Metallurgical properties and characteristics of materials play an important role in their corrosion. Crystalline structure, grain boundaries, specific mechanical properties of metals and alloys, casting methods, heat treatment and chemical composition of alloys are important and effective factors that must be considered.
Other factors that can affect the metals corrosion
The presence of some environmental factors and conditions causes great complexity in corrosion studies. Including:
- The difference in blowing (aeration) at the contact surfaces of the electrolytes causes the formation of anodic regions.
- Concentration differences (or pH differences) at different points of metal surfaces in the electrolyte or corrosive environment causes anodic and cathodic regions.
The biological effects and the presence of microorganisms or microorganisms in the corrosion process have been studied and their damage has been revealed.
In some cases, by creating layers with obstacles on metal surfaces, it causes the production of concentration difference cells. In other cases, by absorbing hydrogen from the metal surface and therefore removing it as a resistant agent in corrosion cells leads to the destruction of metals.
Among them are “sulfate-reducing bacteria” that produce iron sulfides in areas close to the cathode points and accelerate or metals corrosion .
Related article: Iron corrosion and prevention methods
Effect of corrosion medium concentration
Increasing the concentration of corrosive environment leads to an increase in the corrosion rate of active metals. But it has a subtle effect on the corrosion rate of a metal with a patio property.
In some cases, with a large increase in the concentration of the environment, the rate of corrosion of metals increases with the behavior of the patio. Lead has this property and it is believed that lead sulfate, which is a protective shell at low concentrations of sulfuric acid, can be soluble in concentrated acid.
The term environmental effects refers to mechanisms that are specific to a particular metal in a particular environment. These mechanisms are often the mechanisms of cracking.
Effect of oxygen concentration and oxidants on corrosion of metals
One of the important causes of corrosion in corrosive aqueous environments is the amount of oxygen dissolved in the aqueous solution, so that with increasing oxygen and oxides, the rate of corrosion constantly increases.
The effect of oxygen and pH is also very effective on the rate of corrosion metals. Existence of air blowing (aeration) at adjacent surface with electrolytes causes anodic and cathodic regions.
For example, in the case of solutions that come in contact with air oxygen, air oxygen sometimes does not reach the entire metal surface evenly. This condition causes the dissolved oxygen concentration in one area of the solution to increase and decrease in another area. In such solutions, blown dissociation cells will form and cause corrosion.
Metals corrosion in water environment
Water as a good electrolyte is always one of the main causes of rust and corrosion of metals. Corrosion of metals in water is due to the fact that water reacts with carbon dioxide in the air to form carbonic acid, which is a weak acid. Carbonic acid is a stronger electrolyte than water and accelerates the rusting of metals.
Metals corrosion in seawater
Seawater is one of the corrosive environments for metals that marine industries and desalination plants are most expected to. Contrary to popular belief, seawater is not just a solution of mineral salts, but contains a mixture of salts, dissolved gases, suspended solids and organic elements, all of which can affect metals.
The rate of metals corrosionin seawater depends not only on the type of metal or alloy used, but also on the depth, temperature, amount and type of dissolved gases, organic and inorganic compounds and many biological factors of seawater.
Related content: Corrosion management
Marcel Pourbaix has developed a unique and concise method of summarizing the thermodynamic information of metals corrosion for a given metal in an efficient-pH potential diagram.
These diagrams indicate certain areas of potential and pH where the metal is corroded and other areas where corrosion is protected.
Such diagrams are commonly referred to as “pourbaix diagrams”, but are sometimes also called equilibrium diarams; Because these diagrams are used for situations where the metal is in equilibrium with its environment. Purebred patterns are available for more than 70 different metals.
An example of a pourbaix diagram is shown in the figure below, which shows the pourbaix diagram for aluminum. The horizontal axis in the diagram is the pH of the aqueous solution, which is a measure of the chemical environment. The vertical axis is the potential of the E electrochemical environment.
In a pourbaix diagram, there are three possible types of straight:
1.Horizontal line, which are for reactions involving only the electrode potential of the E (but not the pH)
2.Vertical lines, which are for reactions involving only pH (but not electrode potential E)
3.Slanted lines, which pertain to reactions involving both the electrode potential E and the pH.
Pourbaix diagrams also contain regions or parts between different lines in which specific chemical compounds or species are thermodynamically stable. Aluminum pourbiax diagrams in the figure below, the different regions of each of species Al (solid). The ions in them are stable, it shows. The ions in them are stable, it shows.
When a stable species is an ion dissolved, that area is known in the pourbiax diagrams as the “corrosion” area. When a stable species is an oxide or solid hydroxide, that area is called the “emergence” area in the pourbaix diagrams, where the metal is protected against metals corrosion by an oxide or hydroxide surface film.
When the stable species does not react with the metal itself, that area is the safe area. Krueger described pourbiax’s pattern as a “feasibility map”.
Pourbiax diagram for iron
The diagram of iron pourbiax is shown in the figure below. This diagram is of particular interest due to the widespread use of iron and its alloys in construction. Iron can be corroded in acidic or neutral solutions in two different oxidation states, Fe2+ and Fe3+.
Coating is provided by Fe2O 4 and Fe2O3 oxide films. Corrosion in alkaline solutions occurs in the form of complex anions , which is analogous to the dissolved ions , that are similar to aluminum and zinc, respectively, in alkaline solutions.
The simplified pourbiax diagram for iron shown in the figure offers three different concepts of corrosion protection for removing iron from the corrosion area.
For example, the pH of six and the electrode potential of -0.4 volts relative to SHE are related to the corrosion region in the form of Fe2+ ions. Here are three ways to control corrosion:
1. If the potential of the electrode changes in a negative direction to less than -0.7 volts relative to SHE, the iron electrode is moved to the safety zone. This process is called cathodic protection.
2.If instead, the potential of the electrode changes in a positive direction to values higher than about 0.0 volts relative to SHE, the iron electrode is pushed to the surface of the surface. This process is called Anodic protection.
3.The third method of protection is to change the pH of the aqueous solution. If the pH rises to about eight or more, the iron electrode will be in the supernatant again.
Limitations of the Pourbiax diagram
1.Equilibrium is assumed (but in practical cases, the real situation may be far from equilibrium)
2. Pourbiax diagrams do not provide any information about actual corrosion rates.
3.Pourbiax diagrams are for single-element metals only and not for alloys (for a solid solution binary alloy, two-component pourbiax diagrams, as will be shown below, can be combined as a baseline estimate. For engineering alloys, Experimental diagrams can be expanded).
4.Depreciation is attributed to all oxides and hydroxides, regardless of their actual protective properties (corrosion may sometimes be caused by the penetration of ions through oxide films, a process that is ignored in the structure of the diagrams).
5.Pourbiax diagrams do not consider local corrosion by chloride ions.
6.Common pourbiax diagrams are dedicated to the temperatures of 25C° (pourbiax diagrams are available for high temperatures).
Applications of Pourbiax diagrams
At present, many applications of corrosion diagrams in corrosion have been considered. The various applications can be listed as follows:
1. Resistance of metals to uniform corrosion in aqueous solutions
2.The basis for determining which metals can be expected to grow in a wide range of pH and potential conditions.
3.Evaluating the possibility of using oxidizing inhibitors.
4.Identifying the internal conditions of a localized corrosion.
M. Pourbaix, “Atlas of Electrochemical Equilibria in Aqueous Solutions”, National Association of Corrosion Engineers, Houston, TX (1974)
Kruger in “Equilibrium Diagrams: Localized Corrosion”, R. P. Frankenthal Kruger, Eds., p. 45, The Electrochemical Society, Pennington, NJ (1984)