Dataplate

This article explains the metal corrosion process including the basic chemistry of how metal loss occurs. Methods to control corrosion are also listed and explained.

It is convenient to classify corrosion by the forms in which it manifests itself, the basis for this classification being the appearance of the corroded metal. Each form can be identified by mere visual observation. In most cases the naked eye is sufficient, but sometimes magnification is helpful or required. Valuable information for the solution of a corrosion problem can often be obtained through careful observation of the corroded test specimens or failed equipment. Examination before cleaning is particularly desirable. Some of the eight forms of corrosion are unique, but all of them are more or less interrelated. The eight forms are: (1) uniform, or general attack, (2) galvanic, or two-metal corrosion, (3) crevice corrosion, (4) pitting, (5) intergranular corrosion, (6) selective leaching, or parting, (7) erosion corrosion, and (8) stress corrosion. This listing is arbitrary but covers practically all corrosion failures and problems. The forms are not listed in any particular order of importance. Below, the eight forms of corrosion are discussed in terms of their characteristics, mechanisms, and preventive measures. Hydrogen damage, although not a form of corrosion, often occurs indirectly as a result of corrosive attack, and is therefore included in this discussion.

Uniform Attack

Uniform attack is the most common form of corrosion. It is normally characterized by a chemical or electrochemical reaction which proceeds uniformly over the entire exposed surface or over a large area. The metal becomes thinner and eventually fails. For example, a piece of steel or zinc immersed in dilute sulfuric acid will normally dissolve at a uniform rate over its entire surface. A sheet iron roof will show essentially the same degree of rusting over its entire outside surface.

Uniform attack, or general overall corrosion, represents the greatest destruction of metal on a tonnage basis. This form of corrosion, however, is not of too great concern from the technical standpoint, because the life of equipment can be accurately estimated on the basis of comparatively simple tests. Merely immersing specimens in the fluid involved is often sufficient. Uniform attack can be prevented or reduced by (1) proper materials, including coatings, (2) inhibitors, or (3) cathodic protection.

Galvanic or Two-Metal Corrosion

A potential difference usually exists between two dissimilar metals when they are immersed in a corrosive or conductive solution. If these metals are placed in contact (or otherwise electrically connected), this potential difference produces electron flow between them. Corrosion of the less corrosion-resistant metal is usually increased and attack of the more resistant material is decreased, as compared with the behavior of these metals when they are not in contact. The less resistant metal becomes anodic and the more resistant metal cathodic. Usually the- cathode or cathodic metal corrodes very little or not at all in this type of couple. Because of the electric currents and dissimilar metals involved, this form of corrosion is called galvanic, or two-metal, corrosion. It is electrochemical corrosion, but we shall restrict the term galvanic to dissimilar-metal effects for purposes of clarity.

Crevice Corrosion

Intense localized corrosion frequently occurs within crevices and other shielded areas on metal surfaces exposed to corrosives. This type of attack is usually associated with small volumes of stagnant solution caused by holes, gasket surfaces, lap joints, surface deposits, and crevices under bolt and rivet heads. As a result, this form of corrosion is called crevice corrosion or, sometimes, deposit or gasket corrosion. 

Pitting

Pitting is a form of extremely localized attack that results in holes in the metal. These holes may be small or large in diameter, but in most cases they are relatively small. Pits are sometimes isolated or so close together that they look like a rough surface. Generally a pit may be described as a cavity or hole with the surface diameter about the same as or less than the depth.

Pitting is one of the most destructive and insidious forms of corrosion. It causes equipment to fail because of perforation with only a small percent weight loss of the entire structure. It is often difficult to detect pits because of their small size and because the pits are often covered with corrosion products. In addition, it is difficult to measure quantitatively and compare the extent of pitting because of the varying depths and numbers of pits that may occur under identical conditions. Pitting is also difficult to predict by laboratory tests. Sometimes the pits require a long time-several months or a year-to show up in actual service. Pitting is particularly vicious because it is a localized and intense form of corrosion, and failures often occur with extreme suddenness. 

Intergranular Corrosion

Grain boundary effects are of little or no consequence in most applications or uses of metals. If a metal corrodes, uniform attack results since grain boundaries are usually only slightly more reactive than the matrix. However, under certain conditions, grain interfaces are very reactive and intergranular corrosion results. Localized attack at and adjacent to grain boundaries, with relatively little corrosion of the grains, is intergranular corrosion. The alloy disintegrates (grains fall out) and/or loses its strength.

Intergranular corrosion can be caused by impurities at the grain boundaries, enrichment of one of the alloying elements, or depletion of one of these elements in the grain-boundary areas. Small amounts of iron in aluminum, wherein the solubility of iron is low, have been shown to segregate in the grain boundaries and cause intergranular corrosion. It has been shown that based on surface tension considerations the zinc content of a brass is higher at the grain boundaries. Depletion of chromium in the grain-boundary regions results in intergranular corrosion of stainless steels. 

Selective leaching

Selective leaching is the removal of one element from a solid alloy by corrosion processes. The most common example is the selective removal of zinc in brass alloys (dezincification). Similar processes occur in other alloy systems in which aluminum; iron, cobalt, chromium, and other elements are removed. Selective leaching is the general term to describe these processes, and its use precludes the creation of terms such as dealuminumification, decobaltification, etc. Parting is a metallurgical term that is sometimes applied, but selective leaching is preferred.

Erosion Corrosion

Erosion corrosion is the acceleration or increase in rate of deterioration or attack on a metal because of relative movement between a corrosive fluid and the metal surface. Generally, this movement is quite rapid, and mechanical wear effects or abrasion are involved. Metal is removed from the surface as dissolved ions, or it forms solid corrosion products which are mechanically swept from the metal surface. Sometimes, movement of the environment decreases corrosion, particularly when localized attack occurs under stagnant conditions, but this is not erosion corrosion because deterioration is not increased.

Erosion corrosion is characterized in appearance by grooves, gullies, waves, rounded holes, and valleys and usually exhibits a directional pattern. In many cases, failures because of erosion corrosion occur in a relatively short time, and they are unexpected largely because evaluation corrosion tests were run under static conditions or because the erosion effects were not considered. 

Stress-corrosion cracking

Stress-corrosion cracking refers to cracking caused by the simultaneous presence of tensile stress and a specific corrosive medium. Many investigators have classified all cracking failures occurring in corrosive mediums as stress-corrosion cracking, including failures due to hydrogen embrittlement. However, these two types of cracking failures respond differently to environmental variables. To illustrate, cathodic protection is an effective method for preventing stress-corrosion cracking whereas it rapidly accelerates hydrogen-embrittlement effects. Hence, the importance of considering stress-corrosion cracking and hydrogen embrittlement as separate phenomena is obvious. For this reason, the two cracking phenomena are discussed separately in this chapter.

During stress-corrosion cracking, the metal or alloy is virtually unattacked over most of its surface, while fine cracks progress through it. This cracking phenomenon has serious consequences since it can occur at stresses within the range of typical design stress. Exposure to boiling MgCl₂ at 310°F (154°C) is shown to reduce the strength capability to approximately that available at 1200°F.

The two classic cases of stress-corrosion cracking are "season cracking" of brass, and the "caustic embrittlement" of steel. Both of these obsolete terms describe the environmental conditions present which led to stress-corrosion cracking. Season cracking refers to the stress-corrosion cracking failure of brass cartridge cases. During periods of heavy rainfall, especially in the tropics, cracks were observed in the brass cartridge cases at the point where the case was crimped to the bullet. It was later found that the important environmental component in season cracking was ammonia resulting from the decomposition of organic matter.

Many explosions of riveted boilers occurred in early steam-driven locomotives. Examination of these failures showed cracks or brittle failures at the rivet holes. These areas were cold-worked during riveting operations, and analysis of the whitish deposits found in these areas showed caustic, or sodium hydroxide, to be the major component. Hence, brittle fracture in the presence of caustic resulted in the term caustic embrittlement.While stress alone will react in ways well known in mechanical metallurgy (i.e., creep, fatigue, tensile failure) and corrosion alone will react to produce characteristic dissolution reactions; the simultaneous action of both sometimes produces the disastrous results. 

The Corrosion Process

Metal corrosion is a chemical reaction between a metal surface and its environment. Corrosion can occur in a gaseous (dry) environment or a damp (wet) environment. Figure 1 shows the behaviour at the atomic level of both dry and wet environmental corrosion.

Corrosion in a gaseous environment produces a surface layer of converted metal. For example atmospheric corrosion of zinc produces the dull, gray zinc oxide layer seen on galvanised street light poles. Unoxidised zinc coating fresh from the hot dip galvanisers is bright and shiny.

Corrosion in a wet environment attacks the metal by removing the atoms on the metal surface. The metal atoms at the surface lose electrons and become actively charged ions that leave the metal and enter the ‘wet’ electrolyte. The metal ions join with/to oppositely charged ions from another chemical and form a new, stable compound.

Corrosion requires energy. During corrosion the reacting components go from a higher to a lower energy state and release the energy needed for the reaction. In the dry corrosion process of Figure 1 the metal and the oxygen combine to produce the oxide on the surface because the reaction leads to a compound (the oxide) at a lower energy level.

The oxide layer shields the metal from the oxygen and forms a barrier. The oxide will not react with the oxygen in the air or the metal. The barrier makes it difficult for oxygen in the air to contact the metal and it eventually grows so thick that the movement of electrons and ions across it stop. Provided the oxide layer does not crack, or is not removed, the metal is protected from further corrosion. But if the bare metal is exposed to the oxygen, it will again react to form the oxide. In this case the presence of oxygen benefits the metal’s protection. Removal of the oxygen removes the metal’s ability to create its own protective corrosion barrier.

In the wet corrosion process of Figure 1 the electrons from the corroding anode metal move to the connected cathode where they recombine with the atoms of oxygen and water in the electrolyte to make a new hydroxyl ion (OH-). This new negatively charged ion then reacts to make a stable compound with the positively charged metal ions (M++) that originally lost the electrons. In this case, the electrons have a continuous pathway to escape the parent metal and the parent metal, which cannot develop a protective barrier, disassociates or falls apart. Once corrosion starts it continues until the ingredients are all used up.

The electrolyte in wet corrosion can be neutral, acidic or alkaline. For corrosion in near neutral solutions (pH 6 – 8) under oxygenated conditions the predominant cathodic reaction is the oxygen absorption reaction (O2 + 2H2O + 4e- = 4OH-) shown in Figure 1. If instead the bimetallic cell has no oxygen present in the electrolyte the hydrogen evolution reaction (H+ + e- = H followed by H + H = H2 gas) becomes the cathodic process and the anode continues to corrode. This reaction is a much slower reaction (the H+ ion has a very low concentration in solution) than the oxygen absorbing reaction. In acidic solutions (pH 0 - 6) the hydrogen ion concentration is higher and the hydrogen evolution reaction is the predominant one. Corrosion rates become extreme as the pH drops (acid gets stronger).

The Electrical Nature of Corrosion

A flow of electrons means there is an electric current. Wet corrosion produces a corrosion cell. Much like a car battery. The electrons used in creating the corrosion product are continually replaced from the corroding metal. The numbers of electrons available for reacting control the amount of current developed between the two metals. The anode cannot corrode unless there is a cathode. One of them will control the rate of electron flow and thus the corrosion rate.

The intensity (the number of electron and positive ion pairs) is dependent on the potential difference, or voltage, which exists between the metals and the surface area of each metal. Different metal combinations have different voltage potentials between them. Joining two metals with a large potential difference between them produces higher corrosion rates than if the metals were close in electrical potential.

Surface Area Effects

The size of the cathode relative to the anode is important. A large cathode has more surface area through which electrons can flow and so develops an intense electric current with the anode (corroding metal). A small anode connected to it is forced to supply these electrons and will quickly corrode and fall apart. Whereas a large anode connected to a small cathode can provide electrons from any location and will take a long time to show evidence of corrosion.

Where a less noble (base, anodic) metal has to be in contact with a noble metal make sure the less noble metal has at least one hundred times more surface area than the noble metal. Remember – large anode, small cathode – not the opposite.

Differential Aeration Effects

The corrosion reaction requires oxygen and where oxygen is present the metal is cathodic and where oxygen is depleted the metal is anodic and corrodes. The parts of the metal in contact with the highest oxygen concentration become cathodic and are protected, and the areas where oxygen concentration is low will corrode. Steel posts dug into the ground will rust just below the surface because of this effect.

Stagnation Effects

During corrosion, ions build up immediately around the anode and cathode saturating their respective regions. The corrosion rate begins to fall due to the concentration of stagnant ions blocking the creation of more ions in the electrolyte. If the ions are removed or more voltage is provided the corrosion rate again picks up. If you want fast corrosion then agitate the electrolyte and add oxygen.

Specific Types of Corrosion

Corrosion produces physical evidence of its presence. The form it takes depends on the mechanism of the corrosion. Some of the more common forms are explained below.

Pitting Corrosion

A metal can corrode without being in contact with another metal. In this case different areas of the metal take on different electrical potentials. This can occur because of variations in the metal metallurgical properties or because of variations in the surface oxide layer, such as a break, thinning, inclusion like mill scale, contaminant like dirt, etc.

In pitting corrosion the metal at the top of the pit has access to the oxygen in the air and becomes the cathode. At the bottom of the pit oxygen is depleted and the metal becomes the anode. The deeper the pit is the less the oxygen available at the bottom and the corrosion rate increases. Figure 2 shows the mechanism of pitting corrosion.

Crevice Corrosion

A crevice is created whenever two objects are brought together. Unless they are perfectly flat a crevice is present and oxygen cannot easily enter the gap but is plentiful outside. Corrosion starts in the crevice because of differential aeration. Figure 3 is a drawing of crevice corrosion occurring under a layer of seawater.

Stress Corrosion

Metal under tensile stresses can corrode at higher rates than normally expected. The stressed areas have changed electrical potentials to the neighbouring metal and are also more likely to develop microscopic surface cracks. Both situations promote increased corrosion rates.

Bacterial Corrosion

There exist many species of bacteria living in moist environments that release acidic waste products or that can strip out elemental components of a metal. If these bacteria grow on pipe walls and metal surfaces they will cause corrosion. They occur in both oxygenated (aerobic) and oxygen free (anaerobic) conditions.

Galvanic Corrosion

Galvanic corrosion needs to be watched. Dissimilar metals of different potentials joined together by an electrolyte, like process water or rainwater, will cause the more anodic metal to corrode. Running copper water pipe to a galvanised tank will cause the tank to corrode very quickly. Joining copper to steel is nearly as bad. In the galvanic series listed in Table 1 only join metals that are near each other.

Some protection from galvanic corrosion can be achieved if the electrolyte is not present. Without the availability of water molecules the corrosion reaction stops because the electrons cannot find a host to complete the chemical reaction. Where dissimilar metals must be used, for example aluminum fins on the copper coils of a refrigeration chiller condenser, protect them from contact with water. If water must be used in contact with dissimilar metals insure it is deionised and oxygen free.

Controlling Corrosion

Corrosion control involves hindering the natural chemical reactions that occur between the metal and its environment. The common methods used are to:

Modify the environment
Modify the properties of a metal
Install a protective coating over the metals
Impose an electric current to supply electrons
Change to non-metallic materials
 

Modify the Environment

Removing oxygen from the environment prevents completion of the corrosion process by slowing the chemical reaction requiring electrons. If oxygen can be kept away from the protected cathode then the electrons cannot readily flow, so causing the current to drop and corrosion to slow.

Another technique is to use corrosion inhibitors that combine with the corroding metal (anode) or the protected metal (cathode) to form a barrier layer that reduces the flow of ions and electrons across it to very low values and virtually stops the corrosion. If the protective barrier layer is damaged corrosion restarts, so it is necessary to keep an amount of the inhibitor in contact with the metals. This is a common technique used in boilers to protect them from corrosion.

Modify the Properties of the Metal

From the galvanic series we can see that the more noble metals are less likely to corrode. When these metals are metallurgically combined with those from lower in the series, the resulting alloy takes on corrosion resistant properties. The resistance can come from the development of a protective oxide film on the outside surface or because the new alloy has a different voltage potential which acts to make it behave more noble.

Passivation of a metal is a method of changing the potential difference of a metal’s surface. By removing the oxide layer normally present on a metal and exposing the bare metal directly with an acid, the acid reacts with the metal surface to make a new compound with more noble electrical properties. The passivated layer covers the metal, and provided the layer is not broken and the voltage potential remains favourable, it will protect the metal under it from corrosion.

Put Protective Coating Over the Metal

Metals can be protected by covering them in a coating of a different material with better corrosion resistant properties. Metallic and non-metallic coatings are used.

Metallic coatings of less noble metal over more noble metal provide sacrificial protection. Galvanizing is a bonded, protective coat of zinc put over steel. The zinc protects the steel from corrosion in two ways. From the galvanic series it can be seen that zinc will corrode before steel (sacrificial). Secondly, a protective layer of zinc oxide forms on the zinc. If the oxide layer is scratched the zinc is exposed to oxygen and the oxide layer reforms. If part of the zinc coat is lost the rest of the bonded zinc starts to corrode in preference to the steel. As long as the zinc remains in contact with the steel it corrodes sacrificially and protects the steel.

Non-metallic coatings put over a metal can be of two types. They can act as a physical barrier and bar access to the metal surface or they can introduce a very high resistance into the corrosion cell circuit and drastically reduce the flow of electrons. The barrier type coatings protect the metal as long as there are no cracks. If a crack occurs corrosion becomes intense at the metal surface. Resistance type coatings include additives that breakdown in the presence of water and oxygen into inhibiting agents.

Impose an Electric Current

Wet corrosion produces a bi-metallic cell and an electrical current of moving electrons flow from the less noble, anodic metal. If instead the electrons were supplied from another source the less noble metal would not corrode first. This technique is known as cathodic protection. By connecting a more anodic metal into the corrosion circuit than the metal to be protected, the more anodic metal will corrode first and provide an alternate source of electrons. This is why zinc blocks are placed on ship hulls to protect any steel in contact with seawater.

Alternatively a metered electrical current from a power source can be connected to the cathode to supply the electrons. Figure 4 shows cathodic protection installed on a below ground pipeline using sacrificial anodes that require regular replacement as they are corroded away.

Change to Non-Metallic Materials

There are many other materials that can be used instead of metals in situations where metallic corrosion is expected. Provided the physical properties of the non-metal are satisfactory their use may prove a more effective choice.

Practices to Adopt to Reduce Corrosion

• Clean out debris from the bottom of metal sumps exposed to the atmosphere to prevent oxygen depletion under the debris.
  
   
• Repair any damage to painted surfaces quickly to seal the metal off from atmospheric oxygen and moisture.
  
   
• Don’t create galvanic cells by mixing metals. Check the galvanic series when ordering equipment and parts to see if components are compatible. Keep them dry and out of the weather.
  
   
• Check the right type of paint is being used with good corrosion protective properties.
  
   
• Grout under base plates completely and finish the grout at the edges of the lower face. Do not bring it up the side of the base plate and create a crevice.
  
   
• Remove a paint blister immediately, clean the metal thoroughly and repaint to prevent differential aeration.
  
   
• Before bolting or clamping metals parts together insure they both have a completely protected surface. This can be with a sacrificial coating or a barrier coating.
  
   
• Be aware that electric currents can flow for hundreds of meters if parts are joined by an electrolyte. Underground pipes in moist soils need to be protected from bother oxygen depletion and galvanic action from dissimilar metals anywhere along their route.
  
   
• Use a generous radius when bending metals to minimise the creation of internal tensile stresses.
  
   
• Use deionised water when washing down components made of dissimilar metals in contact and dry them off quickly and thoroughly. This is especially important with coils made of dissimilar metals.

Protect underground piping from direct contact with soil and protect the coating from damage. The smallest penetration to the metal will result in rapid pitting corrosion of the pipe wall. If that is not possible install cathodic protection. Beware that some painted coatings still allow moisture and oxygen through to the metal. Check the corrosion protection properties of the coating.