Figure 2.2 Diagrammatic representation of a steel surface, showing anodes and cathodes. This is a simple way of describing the process where iron is removed as charged particles called ions(Fet)and electrons(e") carry current to balance the electric charge Clearly a balancing reaction must occur at the cathode and under ordin ary natural exposure conditions this can be represented as follows: cathodic reaction 2O2+H2 20H In short, hydroxyl ions are produced at the cathode. These two reactions n be combined in a chemical equation: Fe+O2+H2O→2OH-+Fe2+ The ferrous and hydroxyl ions react together to form ferrous hydroxide 2OH-+Fe2t→Fe(OH)2 This is a simple form of rust which is unstable and is eventually oxidised (i.e. reacts with oxygen) to form the familiar reddish brown rust, chemi- cally denoted as FeOoH, or more commonly Fe2O3. H,O. This is the form of rust usually produced in air, natural water and soils. However, under acidic conditions hydrogen is produced at the cathode and the corrosion product may be FeO(magnetite) 2.3 Corrosion terminology Terms frequently used in relation to corrosion are discussed briefly below. Full explanations are available in standard text books 2.3. Potential There is a theoretical e. m f. series of metals(not alloys) called'standard terms for the otentials. These are important in purely electrochemical equilibriun understanding of processes but are of little importance so far C D.A. Bavliss and D. H. Deacon
This is a simple way of describing the process where iron is removed as charged particles called ions (Fe2) and electrons (e) carry current to balance the electric charge. Clearly a balancing reaction must occur at the cathode and under ordinary natural exposure conditions this can be represented as follows: cathodic reaction O2 H2O 2e → 2OH (oxygen) (moisture) (hydroxyl) In short, hydroxyl ions are produced at the cathode. These two reactions can be combined in a chemical equation: FeO2H2O→2OH Fe2 The ferrous and hydroxyl ions react together to form ferrous hydroxide: 2OH Fe2→Fe(OH)2 This is a simple form of rust which is unstable and is eventually oxidised (i.e. reacts with oxygen) to form the familiar reddish brown rust, chemically denoted as FeOOH, or more commonly Fe2O3.H2O. This is the form of rust usually produced in air, natural water and soils. However, under acidic conditions hydrogen is produced at the cathode and the corrosion product may be Fe3O4 (magnetite). 2.3 Corrosion terminology Terms frequently used in relation to corrosion are discussed briefly below. Full explanations are available in standard text books. 2.3.1 Potential There is a theoretical e.m.f. series of metals (not alloys) called ‘standard equilibrium potentials’. These are important in purely electrochemical terms for the understanding of processes but are of little importance so far The corrosion of steel 9 Figure 2.2 Diagrammatic representation of a steel surface, showing anodes and cathodes. © 2002 D. A. Bayliss and D. H. Deacon
as practical corrosion problems are concerned. More useful are potentials experimentally measured using a suitable reference electrode and pub lished in tables such as the Galvanic series in Sea Water. These have some practical value because the extent of differences in potential between different alloys provides an indication of the effect of coupling hem(see Section 2.2) Potentials are also important in determining the operating effectiv of cathodic protection systems(see Chapter 12) 2.3.2 Polarisation The potential difference between the two electrodes in a cell provides the driving force for the current, which determines the extent of corrosion at the anode. However, when the cell is operating, i.e. when current is flowing, the e. m.f. of the cell is different from that theoretically predicted by taking the difference in potentials of the two metallic electrodes. Polar isation occurs at both the anode and the cathode Polarisation, sometimes termed overpotential or overvoltage, can be defined as the difference of the potential of an electrode from its equilib rium or steady-state potential. This can be considered in terms of the energy required to cause a reaction to proceed. An analogy would be the initial energy required to push a car on a level path. Once the car is moving, less energy is required, but if a slope is reached the energy required on the level is not sufficient to push it up the slope, so it tends to slow down and eventually stops Once the cell is operating, changes occur in the cell; ions tend to collect near the anode and reactants tend to surround the cathode. the net result is reduction in the potential difference between the electrodes 2.3.3 Passivity Under certain conditions a corrosion product forms on the surface of a metal, providing a barrier to the environment, i.e. it acts in a similar way to a coating. To achieve passivity the corrosion product must adhere to the urface and be stable both chemically and physically so that it does not dis ate. Such products are sometimes produced near anodic sites and so tend to passivate these areas. Iron becomes passive when immersed in con centrated nitric acid because a thin film of ferric oxide is formed, which, pro- vided it is not disrupted, isolates the iron from the corrosive environment A good example of a passive film is that produced on stainless steels where, because of the chromium content of the alloy(over 12%),a very resistant film, basically Cr2O3, is formed. This not only stops stainless steels from corroding to any extent in air, but also has the capability of rapidly reforming if it is damaged. It is interesting to note that the corrosion per C D.A. Bavliss and D. H. Deacon
as practical corrosion problems are concerned. More useful are potentials experimentally measured using a suitable reference electrode and published in tables such as the ‘Galvanic Series in Sea Water’. These have some practical value because the extent of differences in potential between different alloys provides an indication of the effect of coupling them (see Section 2.2). Potentials are also important in determining the operating effectiveness of cathodic protection systems (see Chapter 12). 2.3.2 Polarisation The potential difference between the two electrodes in a cell provides the ‘driving force’ for the current, which determines the extent of corrosion at the anode. However, when the cell is operating, i.e. when current is flowing, the e.m.f. of the cell is different from that theoretically predicted by taking the difference in potentials of the two metallic electrodes. Polarisation occurs at both the anode and the cathode. Polarisation, sometimes termed overpotential or overvoltage, can be defined as the difference of the potential of an electrode from its equilibrium or steady-state potential. This can be considered in terms of the energy required to cause a reaction to proceed. An analogy would be the initial energy required to push a car on a level path. Once the car is moving, less energy is required, but if a slope is reached the energy required on the level is not sufficient to push it up the slope, so it tends to slow down and eventually stops. Once the cell is operating, changes occur in the cell; ions tend to collect near the anode and reactants tend to surround the cathode. The net result is reduction in the potential difference between the electrodes. 2.3.3 Passivity Under certain conditions a corrosion product forms on the surface of a metal, providing a barrier to the environment, i.e. it acts in a similar way to a coating. To achieve passivity the corrosion product must adhere to the surface and be stable both chemically and physically so that it does not disintegrate. Such products are sometimes produced near anodic sites and so tend to passivate these areas. Iron becomes passive when immersed in concentrated nitric acid because a thin film of ferric oxide is formed, which, provided it is not disrupted, isolates the iron from the corrosive environment. A good example of a passive film is that produced on stainless steels where, because of the chromium content of the alloy (over 12%), a very resistant film, basically Cr2O3, is formed. This not only stops stainless steels from corroding to any extent in air, but also has the capability of rapidly reforming if it is damaged. It is interesting to note that the corrosion per- 10 Steelwork corrosion control © 2002 D. A. Bayliss and D. H. Deacon
formance of stainless steel is determined by its ability to maintain a passive surface film and, if it is broken down and not repaired, the potential of the alloy changes dramatically and is moved towards the anodic end of the 2.4 Corrosion in air Clearly, there is a plentiful supply of oxygen in air, so the presence or otherwise of moisture determines whether corrosion will occur. steel is often visibly moist after rain or when there has been fog or dew. However, the water vapour in the air can also cause steel to rust even though no visible moisture is present. The amount of water vapour in the air is indicated by the relative humidity, and Vernon carried out some experiments which showed the effect of relative humidity on rusting. He showed that in pure air there ence of small concentrations of impurities, such as sulphur dioxide o pre usting could occur above a certain critical humidity, which was about 70%. Below this level, rusting is slight provided moisture from other sources is not present(see Figure 2.3). These experiments showed the importance of two factors that determine the corrosion rate in air, i.e Air polluted with SO, and with Pure ai Figure 2.3 lified diagram showing the effect of relative humidity and pollution on C D.A. Bavliss and D. H. Deacon
formance of stainless steel is determined by its ability to maintain a passive surface film and, if it is broken down and not repaired, the potential of the alloy changes dramatically and is moved towards the anodic end of the galvanic series. 2.4 Corrosion in air Clearly, there is a plentiful supply of oxygen in air, so the presence or otherwise of moisture determines whether corrosion will occur. Steel is often visibly moist after rain or when there has been fog or dew. However, the water vapour in the air can also cause steel to rust even though no visible moisture is present. The amount of water vapour in the air is indicated by the relative humidity, and Vernon carried out some experiments which showed the effect of relative humidity on rusting.1 He showed that in pure air there was little corrosion below 100% relative humidity (r.h.), but in the presence of small concentrations of impurities, such as sulphur dioxide, serious rusting could occur above a certain critical humidity, which was about 70%. Below this level, rusting is slight provided moisture from other sources is not present (see Figure 2.3). These experiments showed the importance of two factors that determine the corrosion rate in air, i.e. The corrosion of steel 11 100 Air polluted with SO2 and solid particles 50 60 70 80 90 Relative humidity (%) Air polluted with SO2 Pure air Corrosion loss Figure 2.3 Simplified diagram showing the effect of relative humidity and pollution on the corrosion of carbon steel. Source: Vernon.1 © 2002 D. A. Bayliss and D. H. Deacon
(i) relative humidity, and (i)pollutants and contaminants The effect of moisture is related to the length of time it is in contact with the steels, so the influence of relative humidity is generally more important than that of precipitation processes such as rain, because the relative humidity may remain above 70% for long periods, particularly in the United Kingdom and other northern European countries. However, in the absence of pollution such as sulphur dioxide(SO2), corrosion is only slight, but a reasonably linear correlation has been shown to exist between the orrosion rate and the amount of SO2 in the air. Although SO2 can dis- solve in moisture to form acids, the effect is not to produce a direct attack on the steel, but rather is the formation of salts such as ferrous sulphate (FeSO4). These compounds, sometimes called corrosion salts, are able by complex reactions, to produce further rusting. Additionally, they are hygroscopic and so can trap further moisture on the steel surface. Such one of the main causes of coating breakdown when paints are applied is salts are of more than academic interest because their presence in rust is rusted surfaces(see Chapter 3). When chlorides are present on the steel urface, typically near the coast where sea salts(sodium chloride)are prevalent, corrosion may occur at relative humidities as low as 40% Generally, the amount of chloride in the air drops off rapidly as the dis tance from the coast increases. The effects of this drop in corrosion are illustrated in Tables 2.1 and 2.2.4 Results from overseas sites showed clearly that in warm, dry, unpolluted inland sites, such as Khartoum in the sudan and delhi in India. the sion rate was negligible compared with that occurring in industri Probably the most interesting result arising from these tests was he surf beach at Lagos, where the corrosion rate was over 0.6mm per year nearly five times that at Sheffield. Conditions where sea spray continu ously reaches the steel surface always lead to severe corrosion. The splash zone on offshore structures is always a critical area for corrosion; see Figure 2.4 Table 2. Effect of sea salts on the corrosion of steel Distance from coast Salt content of air Corrosion rate (mm per year) 400 006 a Based on tests carried out in Nigeria. b Expressed as a percentage of content at 50 yards. C D.A. Bavliss and D. H. Deacon
(i) relative humidity, and (ii) pollutants and contaminants The effect of moisture is related to the length of time it is in contact with the steels, so the influence of relative humidity is generally more important than that of precipitation processes such as rain, because the relative humidity may remain above 70% for long periods, particularly in the United Kingdom and other northern European countries. However, in the absence of pollution such as sulphur dioxide (SO2), corrosion is only slight, but a reasonably linear correlation has been shown to exist between the corrosion rate and the amount of SO2 in the air.2 Although SO2 can dissolve in moisture to form acids, the effect is not to produce a direct attack on the steel, but rather is the formation of salts such as ferrous sulphate (FeSO4). These compounds, sometimes called corrosion salts, are able, by complex reactions, to produce further rusting. Additionally, they are hygroscopic and so can trap further moisture on the steel surface. Such salts are of more than academic interest because their presence in rust is one of the main causes of coating breakdown when paints are applied to rusted surfaces (see Chapter 3). When chlorides are present on the steel surface, typically near the coast where sea salts (sodium chloride) are prevalent, corrosion may occur at relative humidities as low as 40%.3 Generally, the amount of chloride in the air drops off rapidly as the distance from the coast increases. The effects of this drop in corrosion are illustrated in Tables 2.1 and 2.2.4 Results from overseas sites showed clearly that in warm, dry, unpolluted inland sites, such as Khartoum in the Sudan and Delhi in India, the corrosion rate was negligible compared with that occurring in industrial areas. Probably the most interesting result arising from these tests was that for the surf beach at Lagos, where the corrosion rate was over 0.6 mm per year – nearly five times that at Sheffield. Conditions where sea spray continuously reaches the steel surface always lead to severe corrosion. The splash zone on offshore structures is always a critical area for corrosion; see Figure 2.4. 12 Steelwork corrosion control Table 2.1 Effect of sea salts on the corrosion of steela Distance from coast Salt content of airb Corrosion rate (yards) (mm per year) 50 100 0.95 200 27 0.38 400 7 0.06 1300 2 0.04 a Based on tests carried out in Nigeria. b Expressed as a percentage of content at 50 yards. © 2002 D. A. Bayliss and D. H. Deacon
Table 2.2 Corrosivity of environments Cass Annual metal loss ery lwow Rural areas, low pollution, dry Medium Urban and industrial atmospheres Moderate SO2 pollution Industrial and coastal 650-1500g/m2 Industry with high humidity and aggressive atmospher 650-1500g/m2 Marine coastal offshore marine The size, shape and orientation of the steel all influence the corrosion rate to varying extents because they affect the local environment at the steel surface. The orientation of the steel has most influence because it has a marked effect on the ' time-of-wetness' of the surface. In the northern Hemisphere, north-facing steelwork remains moist for longer periods than south-facing steelwork and so tends to corrode more. Again, on horizon- tally exposed steelwork the upper surface may corrode less rapidly than the groundward side because corrosive particles are washed off by rain Figure 2.4 Corrosion in splash zone. C D.A. Bavliss and D. H. Deacon
The size, shape and orientation of the steel all influence the corrosion rate to varying extents because they affect the local environment at the steel surface. The orientation of the steel has most influence because it has a marked effect on the ‘time-of-wetness’ of the surface. In the Northern Hemisphere, north-facing steelwork remains moist for longer periods than south-facing steelwork and so tends to corrode more. Again, on horizontally exposed steelwork the upper surface may corrode less rapidly than the groundward side because corrosive particles are washed off by rain The corrosion of steel 13 Table 2.2 Corrosivity of environments Class Annual metal loss Exterior Very low–low <10–200g/m2 Rural areas, low pollution, dry Medium 200–400g/m2 Urban and industrial atmospheres Moderate SO2 pollution Moderate coastal Cl High 400–650g/m2 Industrial and coastal Very high 650–1500g/m2 Industry with high humidity and industrial aggressive atmosphere Very high 650–1500g/m2 Marine coastal, offshore, marine high salinity Figure 2.4 Corrosion in splash zone. © 2002 D. A. Bayliss and D. H. Deacon