Degradation of Materials (Corrosion) 9.1.Corrosion Mechanisms Despite the numerous useful properties of iron and steel and the cultural changes that came along with the introduction of iron, it has to be kept in mind that iron and steel are plagued by a grave detriment.This detriment is rusting,also often referred to by the generic and more accommodating names corrosion or environmental interaction.Specifically,rusting destroys goods valued at approximately 5%of the gross national product in in- dustrialized countries.Billions of dollars have to be spent annu- ally to replace or repair corrosion-related damages or to prevent corrosion.(About $250 billion per year in the United States alone.)Moreover,corrosion can weaken the strength of struc- tures made from iron and changes their appearance.It is of lit- tle consolation that many other materials such as glass and poly- mers likewise undergo some form of deterioration.Rusting transforms iron or ferrous alloys into ceramic compounds (e.g., iron into iron oxide or hydrated iron oxide),as we shall eluci- date momentarily.Actually,corrosion is a slow form of burning. In short,rusting is a prime destructive mechanism that affects a society which places its trust and investments into iron and steel. Many industrialists do not see these facts as just stated.For one,they consider the advantage of rusting to be that it creates a "repetitive market"for the goods which have been destroyed. This yields jobs for otherwise unemployed people and dividends for investors.Further,the usage of iron and steel is often an eco- nomic factor.Occasionally,replacing relatively inexpensive goods made of steel may be less costly than building equipment out of more durable materials that may cost several times more
9 Despite the numerous useful properties of iron and steel and the cultural changes that came along with the introduction of iron, it has to be kept in mind that iron and steel are plagued by a grave detriment. This detriment is rusting, also often referred to by the generic and more accommodating names corrosion or environmental interaction. Specifically, rusting destroys goods valued at approximately 5% of the gross national product in industrialized countries. Billions of dollars have to be spent annually to replace or repair corrosion-related damages or to prevent corrosion. (About $250 billion per year in the United States alone.) Moreover, corrosion can weaken the strength of structures made from iron and changes their appearance. It is of little consolation that many other materials such as glass and polymers likewise undergo some form of deterioration. Rusting transforms iron or ferrous alloys into ceramic compounds (e.g., iron into iron oxide or hydrated iron oxide), as we shall elucidate momentarily. Actually, corrosion is a slow form of burning. In short, rusting is a prime destructive mechanism that affects a society which places its trust and investments into iron and steel. Many industrialists do not see these facts as just stated. For one, they consider the advantage of rusting to be that it creates a “repetitive market” for the goods which have been destroyed. This yields jobs for otherwise unemployed people and dividends for investors. Further, the usage of iron and steel is often an economic factor. Occasionally, replacing relatively inexpensive goods made of steel may be less costly than building equipment out of more durable materials that may cost several times more. Degradation of Materials (Corrosion) 9.1 • Corrosion Mechanisms
156 9.Degradation of Materials (Corrosion) Third,it is often emphasized that many goods fail by other modes, e.g.,by fatigue or wear,before major corrosion takes place.Fourth, corrosion may return industrial products back to the environment after,of course,a long time period has elapsed.Finally,some peo- ple consider the color of rusting steel quite appealing and adver- tise it as "weathering steel."Actually,the exteriors of a number of commercial buildings have been covered with"weathering steel." Such a surface is,however,not appropriate for all climate zones because the rust,washed down by the rain,may cause some stain- ing of the surrounding grounds.Regardless,corrosion needs to be understood to stay in control of degradation processes.This shall be attempted in the present chapter. Oxidation There are several types of environmental interactions which ma- terials may undergo.Among them is oxidation,that is,the for- mation of a nonmetallic surface film (or scale)which occurs where a metal is exposed to air.Essentially,most metals and al- loys experience some form of superficial oxidation in various de- grees.Often,these surface films are not necessarily disabling, that is,they create protective layers which shield the underlying material from further attack.A chromium oxide film that forms on stainless steel,as discussed in Chapter 8,or a thin,aluminum oxide film that protects bulk aluminum are examples of this.In some cases,an oxide layer is even most welcome,as in insulat- ing SiO2 films which readily grow on silicon wafers and thus pro- vide a basis for microminiaturization in the electronics industry. P-B Ratio Iron oxide(such as FeO,or Fe2O3,etc.)films initially form a pro- tective layer on iron.However,the specific volumes of the oxides are larger than that of metallic iron.This leads,as the thickness of the oxide layer increases,to high compressive stresses in the film and eventually to a flaking from the bulk.As a result,fresh metal is newly exposed to the environment and the oxidation cy- cle starts again.The ratio between oxide volume and metal vol- ume (both per metal atom)is called the Pilling-Bedworth(P-B) ratio.Its values range from less than 1,as for MgO on Mg (lead- ing to a porous film due to tensile stresses in the film),to more than 2,as in the just-mentioned case of iron oxides on iron.A continuous,nonporous,protective,and adherent film is en- countered when the volumes of oxide and metal are about the same,that is,if the P-B ratio is between 1 and 2,as for Al,Cr, and Ti oxides. Moreover,the difference in thermal expansion coefficients be- tween oxide and metal may lead to cracks in the oxide films when considerable thermal cyclings are imposed on the material.As a
Third, it is often emphasized that many goods fail by other modes, e.g., by fatigue or wear, before major corrosion takes place. Fourth, corrosion may return industrial products back to the environment after, of course, a long time period has elapsed. Finally, some people consider the color of rusting steel quite appealing and advertise it as “weathering steel.” Actually, the exteriors of a number of commercial buildings have been covered with “weathering steel.” Such a surface is, however, not appropriate for all climate zones because the rust, washed down by the rain, may cause some staining of the surrounding grounds. Regardless, corrosion needs to be understood to stay in control of degradation processes. This shall be attempted in the present chapter. There are several types of environmental interactions which materials may undergo. Among them is oxidation, that is, the formation of a nonmetallic surface film (or scale) which occurs where a metal is exposed to air. Essentially, most metals and alloys experience some form of superficial oxidation in various degrees. Often, these surface films are not necessarily disabling, that is, they create protective layers which shield the underlying material from further attack. A chromium oxide film that forms on stainless steel, as discussed in Chapter 8, or a thin, aluminum oxide film that protects bulk aluminum are examples of this. In some cases, an oxide layer is even most welcome, as in insulating SiO2 films which readily grow on silicon wafers and thus provide a basis for microminiaturization in the electronics industry. Iron oxide (such as FeO, or Fe2O3, etc.) films initially form a protective layer on iron. However, the specific volumes of the oxides are larger than that of metallic iron. This leads, as the thickness of the oxide layer increases, to high compressive stresses in the film and eventually to a flaking from the bulk. As a result, fresh metal is newly exposed to the environment and the oxidation cycle starts again. The ratio between oxide volume and metal volume (both per metal atom) is called the Pilling–Bedworth (P–B) ratio. Its values range from less than 1, as for MgO on Mg (leading to a porous film due to tensile stresses in the film), to more than 2, as in the just-mentioned case of iron oxides on iron. A continuous, nonporous, protective, and adherent film is encountered when the volumes of oxide and metal are about the same, that is, if the P–B ratio is between 1 and 2, as for Al, Cr, and Ti oxides. Moreover, the difference in thermal expansion coefficients between oxide and metal may lead to cracks in the oxide films when considerable thermal cyclings are imposed on the material. As a Oxidation P–B Ratio 156 9 • Degradation of Materials (Corrosion)
9.1.Corrosion Mechanisms 157 rule,oxides have a smaller expansion coefficient than the re- spective metals. Free Energy The tendency toward oxidation in gaseous(e.g.,oxygen-contain- of Formation ing)environments is different for various metals.Specifically,the oxidation is driven by the free energy of formation,which de- pends on the reaction temperature as well as on the species it- self.As an example,the free energy of formation for aluminum oxide or titanium oxide has a large negative value that leads to a relatively stable oxide;see Figure 9.1.In contrast to this,cop- per or nickel oxides have a small driving force toward oxidation and therefore form relatively unstable oxides. Rate of The rate of oxidation also should be considered.It depends on Oxidation the kind of film that is forming.For example,a porous film(see above)allows a continuous flow of oxygen to the metal surface which,in turn,leads to a linear oxidation rate with time.In con- trast,the most protective films are known to grow much slower, that is,in general,logarithmically with time.Somewhere in be- tween are the growth rates for nonporous oxide layers,as in iron or copper,where a parabolic time-dependence has been found. Leaching Oxidation is only one form of environmental interaction which materials undergo.Some chemical elements are simply dissolved by aqueous solutions.This is called leaching.The well-known lead contamination of the drinking water in old Rome by their leaden water pipes may serve as an example.Lead contamination in drink- ing water,however,has not been eliminated completely in mod- ern days,as indicated by "consumer warning"labels which are packed along with faucets.The reason for this is lead-containing solder joints(for connecting copper pipes)or the use of leaded brass for faucets to facilitate better machining to final shape.(Be- Unstable 4Cu+02→2Cu20 Free energy oxides 2Fe+O2→2Fe0 FIGURE 9.1.Schematic representation of formation Si+02→Si02 of the free energy of formation for A1+O2→A03 selected metals as a function of tem- perature.A small negative free en- ergy of formation means a small dri- ving force toward oxidation and an Stable oxides unstable oxide.(Richardson-Elling- ham diagram.)
rule, oxides have a smaller expansion coefficient than the respective metals. The tendency toward oxidation in gaseous (e.g., oxygen-containing) environments is different for various metals. Specifically, the oxidation is driven by the free energy of formation, which depends on the reaction temperature as well as on the species itself. As an example, the free energy of formation for aluminum oxide or titanium oxide has a large negative value that leads to a relatively stable oxide; see Figure 9.1. In contrast to this, copper or nickel oxides have a small driving force toward oxidation and therefore form relatively unstable oxides. The rate of oxidation also should be considered. It depends on the kind of film that is forming. For example, a porous film (see above) allows a continuous flow of oxygen to the metal surface which, in turn, leads to a linear oxidation rate with time. In contrast, the most protective films are known to grow much slower, that is, in general, logarithmically with time. Somewhere in between are the growth rates for nonporous oxide layers, as in iron or copper, where a parabolic time-dependence has been found. Oxidation is only one form of environmental interaction which materials undergo. Some chemical elements are simply dissolved by aqueous solutions. This is called leaching. The well-known lead contamination of the drinking water in old Rome by their leaden water pipes may serve as an example. Lead contamination in drinking water, however, has not been eliminated completely in modern days, as indicated by “consumer warning” labels which are packed along with faucets. The reason for this is lead-containing solder joints (for connecting copper pipes) or the use of leaded brass for faucets to facilitate better machining to final shape. (BeFree Energy of Formation Rate of Oxidation Leaching 9.1 • Corrosion Mechanisms 157 0 Unstable oxides Stable oxides Free energy of formation 4Cu + O2 2Cu2O Al + O2 Al2O3 2Fe + O2 2FeO Si + O2 SiO2 T 4 3 2 3 FIGURE 9.1. Schematic representation of the free energy of formation for selected metals as a function of temperature. A small negative free energy of formation means a small driving force toward oxidation and an unstable oxide. (Richardson–Ellingham diagram.)
158 9.Degradation of Materials (Corrosion) cause of governmental regulations,manufacturers are now switch- ing to bismuth-copper alloys.)"Soft"water is generally more cor- rosive (dissolves more trace elements)than unsoftened water.[To avoid lead contamination of drinking water,it is often recom- mended to run the water for a few seconds before use,particu- larly if it was standing for some time (e.g.,overnight)in the line or faucets.Moreover,it is advisable to avoid drinking or cooking with water drawn from the hot water side of the tap altogether. Dealloying Selective leaching or dealloying are terms used when one al- loy constituent is preferentially dissolved by a solution.Selective leaching of zinc from brass (called dezincification),or selective loss of graphite in buried gray cast iron gas lines (possibly lead- ing to porosity and explosions)may serve as examples. Corrosion is often more fully described by electrochemical reactions in which free electrons are interpreted to be trans- ferred from one chemical species (or from one part of the same species)to another.Specifically,the deterioration of metals and alloys is interpreted to be caused by an interplay between oxi- dation and reduction processes.This shall be explained in a few examples.During the oxidation process,electrons are transferred from,say iron,to another part of iron (or a different metal)ac- cording to the reaction equation: Fe→Fe2++2e- (9.1) or generally for a metal,M: M→Mm++ne-, (9.2) where n is the valency of the metal ion or the number of elec- trons transferred,and e-represents an electron.The site at which oxidation takes place is defined to be the anode.During a re- duction process,the free electrons which may have been gener- ated during oxidation are transferred to another portion of the sample and there become a part of a different chemical species according to: 2H++2e-→H2 (9.3) or in the case of a metal,M: Mn++e-→M. (9.4) The site where reduction takes place is called the cathode.In other words,oxidation and reduction can be considered as mir- ror-imaged processes.As an example,the dissolution of iron in an acid solution(e.g.,HCl)is represented by the above equations (9.1)and (9.3),or,in summary: Fe+2H+→Fe2++H2. (9.5)
cause of governmental regulations, manufacturers are now switching to bismuth–copper alloys.) “Soft” water is generally more corrosive (dissolves more trace elements) than unsoftened water. [To avoid lead contamination of drinking water, it is often recommended to run the water for a few seconds before use, particularly if it was standing for some time (e.g., overnight) in the line or faucets. Moreover, it is advisable to avoid drinking or cooking with water drawn from the hot water side of the tap altogether.] Selective leaching or dealloying are terms used when one alloy constituent is preferentially dissolved by a solution. Selective leaching of zinc from brass (called dezincification), or selective loss of graphite in buried gray cast iron gas lines (possibly leading to porosity and explosions) may serve as examples. Corrosion is often more fully described by electrochemical reactions in which free electrons are interpreted to be transferred from one chemical species (or from one part of the same species) to another. Specifically, the deterioration of metals and alloys is interpreted to be caused by an interplay between oxidation and reduction processes. This shall be explained in a few examples. During the oxidation process, electrons are transferred from, say iron, to another part of iron (or a different metal) according to the reaction equation: Fe Fe2 2e (9.1) or generally for a metal, M: M Mn ne, (9.2) where n is the valency of the metal ion or the number of electrons transferred, and e represents an electron. The site at which oxidation takes place is defined to be the anode. During a reduction process, the free electrons which may have been generated during oxidation are transferred to another portion of the sample and there become a part of a different chemical species according to: 2H 2e H2 (9.3) or in the case of a metal, M: Mn ne M. (9.4) The site where reduction takes place is called the cathode. In other words, oxidation and reduction can be considered as mirror-imaged processes. As an example, the dissolution of iron in an acid solution (e.g., HCl) is represented by the above equations (9.1) and (9.3), or, in summary: Fe 2H Fe2 H2. (9.5) Dealloying 158 9 • Degradation of Materials (Corrosion)
9.2.Electrochemical Corrosion 159 Fe2+ H'H Acid e Fe FIGURE 9.2.Schematic representation of the dissolution (corrosion) of iron in an acid solu- Anode Cathode tion. The iron ions thus created transfer into the solution(or may,in other cases,react to form an insoluble compound).The process just described is depicted schematically in Figure 9.2. Rust forms readily when iron is exposed to damp air(relative humidity >60%)or oxygen-containing water.Rusting under these conditions occurs in two stages in which iron is,for ex- ample,at first oxidized to Fe2+and then to Fe3+according to the following reaction equations: Fe+02+H20→Fe(OH)2 (9.6) and 2Fe(OH)2+02+H20→2Fe(OH)3. (9.7) Fe (OH)3,that is,hydrated ferric oxide,is insoluble in water and relatively inert,that is,cathodic.As outlined above,Fe(OH)3 is not the only form of "rust."Indeed,other species,such as Feo, Fe2O3,FeOOH,and Fe304,generally qualify for the same name. In these cases,different reaction equations than those shown above apply. 9.2.Electrochemical Corrosion Electrochemical corrosion is often studied by making use of two electrochemical half-cells in which each metal is immersed in a one-molar (1-M)solution of its ion.(A 1-M solution contains 1 mole of the species in 1 dm3 of distilled water.One mole is the atomic mass in grams of the species.)The two half-cells are sep- arated by a semipermeable membrane which prevents interdiffu- sion of the solutions but allows unhindered electron transfer.Fig- ure 9.3 depicts an example of such an electrochemical cell or galvanic couple in which a piece of iron is immersed in a solution
The iron ions thus created transfer into the solution (or may, in other cases, react to form an insoluble compound). The process just described is depicted schematically in Figure 9.2. Rust forms readily when iron is exposed to damp air (relative humidity ! 60%) or oxygen-containing water. Rusting under these conditions occurs in two stages in which iron is, for example, at first oxidized to Fe2 and then to Fe3 according to the following reaction equations: Fe 1 2 O2 H2O Fe(OH)2 (9.6) and 2Fe(OH)2 1 2 O2 H2O 2Fe(OH)3. (9.7) Fe (OH)3, that is, hydrated ferric oxide, is insoluble in water and relatively inert, that is, cathodic. As outlined above, Fe(OH)3 is not the only form of “rust.” Indeed, other species, such as FeO, Fe2O3, FeOOH, and Fe3O4, generally qualify for the same name. In these cases, different reaction equations than those shown above apply. Electrochemical corrosion is often studied by making use of two electrochemical half-cells in which each metal is immersed in a one-molar (1-M) solution of its ion. (A 1-M solution contains 1 mole of the species in 1 dm3 of distilled water. One mole is the atomic mass in grams of the species.) The two half-cells are separated by a semipermeable membrane which prevents interdiffusion of the solutions but allows unhindered electron transfer. Figure 9.3 depicts an example of such an electrochemical cell or galvanic couple in which a piece of iron is immersed in a solution 9.2 • Electrochemical Corrosion 159 Fe2+ Acid Anode Cathode Fe e – e – H+ H+ H2 FIGURE 9.2. Schematic representation of the dissolution (corrosion) of iron in an acid solution. 9.2 • Electrochemical Corrosion
160 9.Degradation of Materials (Corrosion) shrinks grows Fe Cu FIGURE 9.3.Schematic representation of Fe2+solution two electrochemical half-cells in which Cu2+solution two metal electrodes are immersed in Anode Cathode 1-M solutions of their ions (Galvanic couple). Membrane that contains 1-M Fe2+ions.The other part consists of an analo- gous copper half-cell.The two metal electrodes are electrically con- nected through a voltmeter for reasons that will become obvious momentarily.After some time one observes that the iron bar has slightly decreased in size,that is,some of the solid Fe atoms have transferred as Fe2+into the solution.On the other hand,some copper ions have been electroplated onto the copper electrode, which,as a consequence,has increased somewhat in size.The per- tinent reaction equations are thus as follows: Fe→Fe2++2e- (9.8) Cu2++2e-→Cu. (9.9) During the entire process,a voltage (called electromotive force, or emf)of 0.78 V is measured.One concludes from this experi- mental result that copper is more noble than iron.Indeed,a table can be created which ranks all metals (or their electrode reac- tions)in decreasing order of inertness to corrosion along with their emfs compared to a reference for which the hydrogen re- action has been arbitrarily chosen;see Table 9.1.(The standard hydrogen reference cell consists of an inert platinum electrode that is "immersed"in 1 atm of flowing hydrogen gas at 25C.) The electromotive forces,Eo,listed in Table 9.1 are valid for 1-M solutions and for room temperature(25C)only.For other conditions,the potential of a system needs to be adjusted by mak- ing use of the Nernst equation: +In C+00392 og CV). (9.10) nF
that contains 1-M Fe2 ions. The other part consists of an analogous copper half-cell. The two metal electrodes are electrically connected through a voltmeter for reasons that will become obvious momentarily. After some time one observes that the iron bar has slightly decreased in size, that is, some of the solid Fe atoms have transferred as Fe2 into the solution. On the other hand, some copper ions have been electroplated onto the copper electrode, which, as a consequence, has increased somewhat in size. The pertinent reaction equations are thus as follows: Fe Fe2 2e (9.8) Cu2 2e Cu. (9.9) During the entire process, a voltage (called electromotive force, or emf) of 0.78 V is measured. One concludes from this experimental result that copper is more noble than iron. Indeed, a table can be created which ranks all metals (or their electrode reactions) in decreasing order of inertness to corrosion along with their emf’s compared to a reference for which the hydrogen reaction has been arbitrarily chosen; see Table 9.1. (The standard hydrogen reference cell consists of an inert platinum electrode that is “immersed” in 1 atm of flowing hydrogen gas at 25°C.) The electromotive forces, E0, listed in Table 9.1 are valid for 1-M solutions and for room temperature (25°C) only. For other conditions, the potential of a system needs to be adjusted by making use of the Nernst equation: E E0 R nF T ln Cion E0 0.0 n 592 log Cion[V], (9.10) 160 9 • Degradation of Materials (Corrosion) Fe Cu shrinks grows Fe2+ solution Cu2+ solution Anode Cathode Membrane e– e– FIGURE 9.3. Schematic representation of two electrochemical half-cells in which two metal electrodes are immersed in 1-M solutions of their ions (Galvanic couple).
9.2.Electrochemical Corrosion 161 TABLE 9.1.Standard emf series for selected elements in 1-M solution at 25C Electrode potential,or emf Electrode reaction Eo(V) Au3++3e-→Au +1.42 Cathodic Pt2++2e-→Pt +1.2 Ag++1e-→Ag +0.80 Cu2++2e-→Cu +0.34 2H++2e-→H2 0.000 Sn2++2e-→Sn -0.14 Ni2++2e-→Ni -0.25 Fe2++2e-→Fe -0.44 Cr3++3e-→Cr -0.74 Zn2++2e-→Zn -0.76 A13++3e-→Al -1.66 Mg2++2e-→Mg -2.36 Anodic (Oxidation) where R and F are gas and Faraday constants,respectively (see Appendix ID),T is the absolute temperature [set equal to 298.2 K or 25C in the right side of Eq.(9.10)].n is again the number of electrons transferred,and Cion is the molar ion concentration.As expected,one obtains E Eo for a 1-M solution.Equation(9.10) shows that a tenfold increase in the ion concentration leads to a rise of the half-cell potential of 59.2 mV for a single electron re- action. Galvanic It is customary to rank metals and alloys in a simpler way,that Series is,in terms of their relative reactivities to each other in a given environment.Table 9.2 presents such a list,called the galvanic series.It includes some commercial metals and alloys that have been exposed to sea water.No emf's are usually included in a galvanic series. Galvanic The galvanic series of electrochemical corrosion explains a num- Corrosion ber of important mechanisms.To start with,when two different metals (such as an iron and a copper pipe)are electrically con- nected and are exposed to an electrolyte (e.g.,ordinary water), the less noble metal (here,the iron)corrodes near the junction (Figure 9.4).This is an example of a mechanism called galvanic corrosion.It can be essentially prevented by inserting an insu- lated coupling(plastic)between the two unlike pipes.(This tech- nique does not prevent galvanic corrosion in the case when,for example,copper is dissolved "upstream"and then plated on the steel pipe "downstream".)Galvanic corrosion also occurs in a
where R and F are gas and Faraday constants, respectively (see Appendix II), T is the absolute temperature [set equal to 298.2 K or 25°C in the right side of Eq. (9.10)], n is again the number of electrons transferred, and Cion is the molar ion concentration. As expected, one obtains E E0 for a 1-M solution. Equation (9.10) shows that a tenfold increase in the ion concentration leads to a rise of the half-cell potential of 59.2 mV for a single electron reaction. It is customary to rank metals and alloys in a simpler way, that is, in terms of their relative reactivities to each other in a given environment. Table 9.2 presents such a list, called the galvanic series. It includes some commercial metals and alloys that have been exposed to sea water. No emf’s are usually included in a galvanic series. The galvanic series of electrochemical corrosion explains a number of important mechanisms. To start with, when two different metals (such as an iron and a copper pipe) are electrically connected and are exposed to an electrolyte (e.g., ordinary water), the less noble metal (here, the iron) corrodes near the junction (Figure 9.4). This is an example of a mechanism called galvanic corrosion. It can be essentially prevented by inserting an insulated coupling (plastic) between the two unlike pipes. (This technique does not prevent galvanic corrosion in the case when, for example, copper is dissolved “upstream” and then plated on the steel pipe “downstream”.) Galvanic corrosion also occurs in a Galvanic Series Galvanic Corrosion 9.2 • Electrochemical Corrosion 161 TABLE 9.1. Standard emf series for selected elements in 1-M solution at 25°C Electrode potential, or emf Electrode reaction E0 (V) Au3 3e Au 1.42 Cathodic Pt2 2e Pt 1.2 Ag 1e Ag 0.80 Cu2 2e Cu 0.34 2H 2e H2 0.000 Sn2 2e Sn 0.14 Ni2 2e Ni 0.25 Fe2 2e Fe 0.44 Cr3 3e Cr 0.74 Zn2 2e Zn 0.76 Al3 3e Al 1.66 Mg2 2e Mg 2.36 Anodic (Oxidation)
162 9.Degradation of Materials (Corrosion) TABLE 9.2.Galvanic series of some commercial metals and alloys exposed to sea watera Platinum Cathodic Gold Graphite Titanium Stainless steel(passive) Copper-nickel alloys Bronze Copper Tin Sn-50%Pb solder Stainless steel (active) Cast iron Iron and steel Aluminum alloys Aluminum Galvanized steel Zinc Magnesium Anodic "Note that each metal or alloy is anodic with respect to those above it. marine environment.Examples are steel rivets that have been (improperly)used to join aluminum sheets,or steel ball bearings of sliding doors that sit on an aluminum track.The list of "gal- vanic corrosion sins"observed in daily life is almost endless and, of course,is not restricted to iron and steel. Galvanic corrosion is also encountered due to concentration dif- ferences of ions or dissolved gases in an electrolyte.The respec- tive mechanism involved has been termed a concentration cell.The deterioration often occurs at the site at which the solution is most diluted.For example,the oxygen concentration is lower immedi- ately under a drop of water sitting on an iron plate.In contrast, the oxygen concentration at the periphery of the water drop is much higher due to its increased exposure to the atmosphere.As a consequence,an iron plate corrodes under the center of the wa- ter drop and forms what is called a corrosion pit;see Figure 9.5(a). Similar observations are made immediately below the waterline Anode Cathode (corrosion) FIGURE 9.4.Galvanic cor- rosion of two dissimilar H2O metals (water pipes)(see Table 9.1)
marine environment. Examples are steel rivets that have been (improperly) used to join aluminum sheets, or steel ball bearings of sliding doors that sit on an aluminum track. The list of “galvanic corrosion sins” observed in daily life is almost endless and, of course, is not restricted to iron and steel. Galvanic corrosion is also encountered due to concentration differences of ions or dissolved gases in an electrolyte. The respective mechanism involved has been termed a concentration cell. The deterioration often occurs at the site at which the solution is most diluted. For example, the oxygen concentration is lower immediately under a drop of water sitting on an iron plate. In contrast, the oxygen concentration at the periphery of the water drop is much higher due to its increased exposure to the atmosphere. As a consequence, an iron plate corrodes under the center of the water drop and forms what is called a corrosion pit; see Figure 9.5(a). Similar observations are made immediately below the waterline 162 9 • Degradation of Materials (Corrosion) TABLE 9.2. Galvanic series of some commercial metals and alloys exposed to sea watera Platinum Cathodic Gold Graphite Titanium Stainless steel (passive) Copper–nickel alloys Bronze Copper Tin Sn–50% Pb solder Stainless steel (active) Cast iron Iron and steel Aluminum alloys Aluminum Galvanized steel Zinc Magnesium Anodic aNote that each metal or alloy is anodic with respect to those above it. Anode (corrosion) Cathode H Fe 2O e– e– Cu FIGURE 9.4. Galvanic corrosion of two dissimilar metals (water pipes) (see Table 9.1)
9.2.Electrochemical Corrosion 163 High oxygen concentration Water Water (cathode) droplet drop e e- Low oxygen Iron plate concentration (anode); corrosion pit (a) (b) FiGURE 9.5.Schematic representation of concentration cells due to an oxy- gen concentration gradient(a)under a water drop and (b)in a crevice. of ships,which is,for more than one reason,the preferred site for deterioration (waterline corrosion).Finally,crevices between two metal plates which have been mechanically joined may trap wa- ter that deprives the affected area from atmospheric oxygen and thus leads to crevice corrosion there;Figure 9.5(b).Crevice corro- sion can be prevented by,for example,closing the ends of the crevice with a weld using a material that is similar to the plates. Cathodic On the positive side,galvanic corrosion can be put to good use Protection for preventing deterioration of buried steel pipes.For this,the metal to be protected needs to be supplied with electrons that force the pipe to become cathodic.Specifically,a less noble metal (e.g.,Mg,Al,Zn;see Table 9.2)is wired to a steel pipe,while both are imbedded in a suitable back-fill of moist soil,Figure 9.6(a). The magnesium (or Mg-Zn alloy)is then slowly consumed and eventually needs to be replaced. Understandably,the less noble metal is called the sacrificial an- ode and the mechanism just described is known by the name cathodic protection.Cathodic protection is also employed for ships,tanks,and hot-water heaters,to name a few examples.An- other method to provide the metal to be protected with electrons is by connecting it to a direct-current power source (e.g.,a solar photovoltaic cell)as depicted in Figure 9.6(b).If unprotected,the corrosion of buried steel pipes would be caused either by long line currents that flow through the pipes from areas of one type of soil to another one,or by short line currents from the bottom to the top of the pipe,again caused by different soil species or by dif- ferential aeration of the soil. A further useful application of galvanic corrosion is made in al- kaline batteries in which the emf,generated between two dissimi- lar metals,is utilized until the less noble metal has been consumed
of ships, which is, for more than one reason, the preferred site for deterioration (waterline corrosion). Finally, crevices between two metal plates which have been mechanically joined may trap water that deprives the affected area from atmospheric oxygen and thus leads to crevice corrosion there; Figure 9.5(b). Crevice corrosion can be prevented by, for example, closing the ends of the crevice with a weld using a material that is similar to the plates. On the positive side, galvanic corrosion can be put to good use for preventing deterioration of buried steel pipes. For this, the metal to be protected needs to be supplied with electrons that force the pipe to become cathodic. Specifically, a less noble metal (e.g., Mg, Al, Zn; see Table 9.2) is wired to a steel pipe, while both are imbedded in a suitable back-fill of moist soil, Figure 9.6(a). The magnesium (or Mg–Zn alloy) is then slowly consumed and eventually needs to be replaced. Understandably, the less noble metal is called the sacrificial anode and the mechanism just described is known by the name cathodic protection. Cathodic protection is also employed for ships, tanks, and hot-water heaters, to name a few examples. Another method to provide the metal to be protected with electrons is by connecting it to a direct-current power source (e.g., a solar photovoltaic cell) as depicted in Figure 9.6(b). If unprotected, the corrosion of buried steel pipes would be caused either by long line currents that flow through the pipes from areas of one type of soil to another one, or by short line currents from the bottom to the top of the pipe, again caused by different soil species or by differential aeration of the soil. A further useful application of galvanic corrosion is made in alkaline batteries in which the emf, generated between two dissimilar metals, is utilized until the less noble metal has been consumed. FIGURE 9.5. Schematic representation of concentration cells due to an oxygen concentration gradient (a) under a water drop and (b) in a crevice. 9.2 • Electrochemical Corrosion 163 Iron plate Water droplet Water drop High oxygen concentration (cathode) Low oxygen concentration (anode); corrosion pit e – e – (a) (b) Cathodic Protection
164 9.Degradation of Materials(Corrosion) Mg2+ Mg e Mg2+ Anode Cathode Soil (a) Solar panel e Scrap iron FIGURE 9.6.(a)Sacrificial anode which is wired to a steel pipeline to provide cathodic protection.(b) Soil A photovoltaic device supplies electrons to a steel pipe for cathodic protection.The connecting wires need to be electrically insulated! (b) Faraday The question then arises,how much metal is electrochemically Equation removed by corrosion (or deposited)in a given time interval,t, while applying a certain current,I?M.Faraday (see Chapter 10) empirically arrived at the following equation: w=9M=1·M n·F n·F3 (9.11) where W is the liberated mass,g is the electric charge,M is the atomic mass of the metal,F is the Faraday constant (see Ap- pendix II),and n is the valence of the metal ion.We shall return to this topic in Chapter 10
The question then arises, how much metal is electrochemically removed by corrosion (or deposited) in a given time interval, t, while applying a certain current, I? M. Faraday (see Chapter 10) empirically arrived at the following equation: W q n M F I n t F M, (9.11) where W is the liberated mass, q is the electric charge, M is the atomic mass of the metal, F is the Faraday constant (see Appendix II), and n is the valence of the metal ion. We shall return to this topic in Chapter 10. FIGURE 9.6. (a) Sacrificial anode which is wired to a steel pipeline to provide cathodic protection. (b) A photovoltaic device supplies electrons to a steel pipe for cathodic protection. The connecting wires need to be electrically insulated! 164 9 • Degradation of Materials (Corrosion) e – e – (a) Mg + Mg2+ Mg2+ – Anode Cathode Soil e – e – (b) + – Scrap iron Soil Solar panel Faraday Equation