Engineering Materials Science and Engineering 7 Corrosion and Control PDF

E N G G 4 1 2 : 7 M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G Chapter Corrosion and Control Intended...

E N G G 4 1 2 : 7 M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G Chapter Corrosion and Control Intended Learning Outcomes After studying this chapter, you should be able to do the following: 1. Distinguish between oxidation and reduction electrochemical reactions. 2. Describe the following: galvanic couple, standard half-cell, and standard hydrogen electrode. 3. Compute the cell potential and write the spontaneous electrochemical reaction direction for two pure metals that are electrically connected and also submerged in solutions of their respective ions. 4. Determine metal oxidation rate, given the reaction current density. 5. Name and briefly describe the two different types of polarization and specify the conditions under which each is rate controlling. 6. For each of the eight forms of corrosion and hydrogen embrittlement, describe the nature of the deteriorative process and then note the proposed mechanism. 7. List five measures that are commonly used to prevent corrosion. 8. Explain why ceramic materials are, in general, very resistant to corrosion. 9. For polymeric materials, discuss (a) two degradation processes that occur when they are exposed to liquid solvents and (b) the causes and consequences of molecular chain bond rupture. The deterioration of each of these material types is discussed in this chapter with special regard to mechanism, resistance to attack by various environments, and measures to prevent or reduce degradation. To one degree or another, most materials experience some type of interaction with a large number of diverse environments. Often, such interactions impair a material’s usefulness as a result of the deterioration of its mechanical properties (e.g., ductility and strength), other physical properties, or appearance. Occasionally, to the chagrin of a design engineer, the degradation behavior of a material for some application is ignored, with adverse consequences. Deteriorative mechanisms are different for the three material types. In metals, corrosion there is actual material loss either by dissolution (corrosion) or by the formation of nonmetallic scale or film (oxidation). Ceramic materials are relatively resistant to deterioration, which usually occurs at elevated temperatures or in rather extreme environments; the process is frequently also called corrosion. For polymers, mechanisms and consequences differ from those for metals and ceramics, and the term degradation is most frequently used. Polymers may dissolve when exposed to a liquid solvent, or they may absorb the solvent and swell; also, electromagnetic radiation (primarily ultraviolet) and heat may cause alterations in their molecular structures. Module No. 7 – Corrosion and Control 1 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G CORROSION OF METALS Corrosion is defined as the destructive and unintentional attack on a metal; it is electrochemical and ordinarily begins at the surface. The problem of metallic corrosion is significant; in economic terms, it has been estimated that approximately 5% of an industrialized nation’s income is spent on corrosion prevention and the maintenance or replacement of products lost or contaminated as a result of corrosion reactions. The consequences of corrosion are all too common. Familiar examples include the rusting of automotive body panels and radiator and exhaust components. ELECTROCHEMICAL CONSIDERATIONS For metallic materials, the corrosion process is normally electrochemical, that is, a chemical reaction in which there is transfer of electrons from one chemical species to another. Metal atoms characteristically lose or give up electrons in what is called an oxidation reaction. For example, a hypothetical metal M that has a valence of n (or n valence electrons) may experience oxidation according to the reaction M Mn+ + ne- (7.1) In which M becomes an n positively charged ion and in the process loses its n valence electrons; e is used to symbolize an electron. Examples in which metals oxidize are Fe Fe2+ + 2e- (7.2) Al Al3+ + 3e- (7.3) The site at which oxidation takes place is called the anode; oxidation is sometimes called an anodic reaction. The electrons generated from each metal atom that is oxidized must be transferred to and become a part of another chemical species in what is termed a reduction reaction. For example, some metals undergo corrosion in acid solutions, which have a high concentration of hydrogen (H+) ions; the H+ ions are reduced as follows: 2H+ + 2e- H2 (7.4) And hydrogen gas (H2) is evolved. Other reduction reactions are possible, depending on the nature of the solution to which the metal is exposed. For an acid solution having dissolved oxygen, reduction according to O2 + 4H+ + 4e- 2H2O (7.5) Will probably occur. For a neutral or basic aqueous solution in which oxygen is also dissolved, O2 + 2H2O + 4e- 4(OH-) (7.6) Any metal ions present in the solution may also be reduced; for ions that can exist in more than one valence state (multivalent ions), reduction may occur by Mn+ + e- M(n-1)+ (7.7a) in which the metal ion decreases its valence state by accepting an electron. A metal may be totally reduced from an ionic to a neutral metallic state according to Mn+ + ne- M (7.7b) The location at which reduction occurs is called the cathode. It is possible for two or more of the preceding reduction reactions to occur simultaneously. An overall electrochemical reaction must consist of at least one oxidation and one reduction reaction and will be their sum; often the Module No. 7 – Corrosion and Control 2 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G individual oxidation and reduction reactions are termed half-reactions. There can be no net electrical charge accumulation from the electrons and ions—that is, the total rate of oxidation must equal the total rate of reduction, or all electrons generated through oxidation must be consumed by reduction. For example, consider zinc metal immersed in an acid solution containing H ions. At some regions on the metal surface, zinc will experience oxidation or corrosion as illustrated in Figure 17.1, according to the reaction Zn Zn2+ + 2e- (7.8) Because zinc is a metal and therefore a good electrical conductor, these electrons may be transferred to an adjacent region at which the H ions are reduced according to 2H+ + 2e- H2 (gas) (7.9) If no other oxidation or reduction reactions occur, the total electrochemical reaction is just the sum of reactions 17.8 and 17.9, or Zn Zn2+ + 2e- + 2H + 2e- H2 (gas) Zn + 2H+ h Zn2+ + H2 (gas) (7.10) Another example is the oxidation or rusting of iron in water, which contains dissolved oxygen. This process occurs in two steps; in the first, Fe is oxidized to Fe2+ [as Fe(OH)2], (7.11) Fe + 2O2 + H2O Fe2 + + 2OH- Fe(OH)2 and, in the second stage, to Fe3+ [as Fe(OH)3] according to (7.12) 1 2 Fe(OH)2 + O2 + H2O 2Fe(OH)3 2 The compound Fe(OH)3 is the all-too-familiar rust. As a consequence of oxidation, the metal ions may either go into the corroding solution as ions (reaction 17.8) or form an insoluble compound with nonmetallic elements as in reaction 7.12. Figure 7.1 The electrochemical reactions associated with the corrosion of zinc in an acid solution. (From M. G. Fontana, Corrosion Engineering, and 3rd edition. Copyright 1986 by McGraw-Hill Book Company. Reproduced with permission.) Module No. 7 – Corrosion and Control 3 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G Figure 7.2 an electrochemical cell consisting of iron and copper electrodes, each of which is immersed in a 1 M solution of its ion. Iron corrodes while copper electrodeposits. ELECTRODE POTENTIALS Not all metallic materials oxidize to form ions with the same degree of ease. Consider the electrochemical cell shown in Figure 7.2. On the left-hand side is a piece of pure iron immersed in a solution containing Fe 2 ions of 1 M concentration.1 The other side of the cell consists of a pure copper electrode in a 1 M solution of Cu 2 ions. The cell halves are separated by a membrane, which limits the mixing of the two solutions. If the iron and copper electrodes are connected electrically, reduction will occur for copper at the expense of the oxidation of iron, as follows: Cu2+ + Fe Cu + Fe2+ Or Cu2+ ions will deposit (electrodeposit) as metallic copper on the copper electrode, whereas iron dissolves (corrodes) on the other side of the cell and goes into solution as Fe 2+ ions. Thus, the two half-cell reactions are represented by the relations Fe Fe2+ + 2e- (7.14a) Cu2+ + 2e- Cu (7.14b) When a current passes through the external circuit, electrons generated from the oxidation of iron flow to the copper cell in order that Cu2+ be reduced. In addition, there will be some net ion motion from each cell to the other across the membrane. This is called a galvanic couple-two metals electrically connected in a liquid electrolyte in which one metal becomes an anode and corrodes while the other acts as a cathode. An electric potential or voltage exists between the two cell halves, and its magnitude can be determined if a voltmeter is connected in the external circuit. A potential of 0.780 V results for a copper-iron galvanic cell when the temperature is 25°C (77°F). Concentration of liquid solutions is often expressed in terms of molarity, M, the number of moles of solute per liter (1000 cm 3) of solution. Module No. 7 – Corrosion and Control 4 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G (Fig. 7.3) (Fig. 7.4) Figure (7.3) an electrochemical cell consisting of iron and zinc electrodes, each of which is immersed in a 1 M solution of its ion. The iron electrodeposits while the zinc corrodes. Figure (7.4) the standard hydrogen reference half-cell. Now consider another galvanic couple consisting of the same iron half-cell connected to a metal zinc electrode that is immersed in a 1 M solution of Zn2 ions (Figure 7.3). In this case, the zinc is the anode and corrodes, whereas the Fe becomes the cathode. The electrochemical reaction is thus Fe2+ + Zn Fe + Zn2+ (7.15) The potential associated with this cell reaction is 0.323 V. Thus, various electrode pairs have different voltages; the magnitude of such a voltage may be thought of as representing the driving force for the electrochemical oxidation–reduction reaction. Consequently, metallic materials may be rated as to their tendency to experience oxidation when coupled to other metals in solutions of their respective ions. A half-cell similar to those just described [i.e., a pure metal electrode immersed in a 1 M solution of its ions and at 25⁰C (77⁰F)] is termed a standard half-cell. THE STANDARD EMF SERIES These measured cell voltages represent only differences in electrical potential, and thus it is convenient to establish a reference point, or reference cell, to which other cell halves may be compared. This reference cell, arbitrarily chosen, is the standard hydrogen electrode (Figure 7.4). It consists of an inert platinum electrode in a 1 M solution of H ions saturated with hydrogen gas that is bubbled through the solution at a pressure of 1 atm and a temperature of 25⁰C (77⁰F). The platinum itself does not take part in the electrochemical reaction; it acts only as a surface on which hydrogen atoms may be oxidized or hydrogen ions may be reduced. The electromotive force (emf) series (Table 7.1) is generated by coupling to the standard hydrogen electrode, standard half-cells for various metals and ranking them according to measured voltage. Table 7.1 shows the corrosion tendencies for the several metals; those at the top (i.e., gold and platinum) are noble, or chemically inert. As one moves down the table, the metals become increasingly more Module No. 7 – Corrosion and Control 5 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G active—that is, more susceptible to oxidation. Sodium and potassium have the highest reactivities. Table 7.1 THE STANDARD EMF SERIES The voltages in Table 7.1 are for the half-reactions as reduction reactions, with the electrons on the left-hand side of the chemical equation; for oxidation, the direction of the reaction is reversed and the sign of the voltage changed. Consider the generalized reactions involving the oxidation of metal M 1 and the reduction of metal M2 as 𝑀1 𝑀1𝑛+ + ne- −𝑉10 (7.16A) 𝑀2𝑛+ + ne- M2 +𝑉20 (7.16B) Where the V 0s are the standard potentials as taken from the standard emf series. Because metal M1 is oxidized, the sign of V0 is opposite to that as it appears in Table 7.1. Addition of Equations 7.16a and 7.16b yields 𝑀1 + 𝑀2𝑁+ 𝑀1𝑁+ +𝑀1 (7.17) And the overall cell potential AV0 is 𝑉0 = 𝑉20 = 𝑉10 (7.18) For this reaction to occur spontaneously, ∆V0 must be positive; if it is negative, the spontaneous cell direction is the reverse of Equation 7.17. When standard half-cells are coupled together, the metal that lies lower in Table 7.1 experiences oxidation (i.e., corrosion), whereas the higher one is reduced. Module No. 7 – Corrosion and Control 6 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G INFLUENCE OF CONCENTRATION AND TEMPERATURE ON CELL POTENTIAL The emf series applies to highly idealized electrochemical cells (i.e., pure metals in 1 M solutions of their ions at 25⁰C). Altering temperature or solution concentration or using alloy electrodes instead of pure metals changes the cell potential, and, in some cases, the spontaneous reaction direction may be reversed. Consider again the electrochemical reaction described by Equation 7.17. If M1 and M2 electrodes are pure metals, the cell potential depends on the absolute temperature T and the molar ion concentrations [Mn+ 1 ] and [Mn+ 2 ] according to the Nernst equation, (7.19) where R is the gas constant, n is the number of electrons participating in either of the half-cell reactions, and F is the Faraday constant, 96,500 C/mol—the magnitude of charge per mole (6.022 x 1023) of electrons. At 25⁰C (about room temperature), (7.20) to give ΔV in volts. Again, for reaction spontaneity, ∆V must be positive. As expected, for 1 M concentrations of both ion types (that is, [Mn+ 1 ] = [Mn+ 2 ] = 1), Equation 7.19 simplifies to Equation 7.18. EXAMPLE PROBLEM 7.1 Determination of Electrochemical Cell Characteristics One-half of an electrochemical cell consists of a pure nickel electrode in a solution of Ni 2+ ions; the other half is a cadmium electrode immersed in a Cd 2+ solution. a. If the cell is a standard one, write the spontaneous overall reaction and calculate the volt- age that is generated. b. Compute the cell potential at 25°C if the Cd2+ and Ni2+ concentrations are 0.5 and 10-3 M, respectively. Is the spontaneous reaction direction still the same as for the standard cell? Solution (a) The cadmium electrode is oxidized and nickel is reduced because cadmium is lower in the emf series; thus, the spontaneous reactions are yields 𝐶𝑑 → 𝐶𝑑 2+ + 2𝑒 − (7.21) 𝑦𝑖𝑒𝑙𝑑𝑠 𝑁𝑖 2+ 2𝑒 − → 𝑁𝑖 𝑦𝑖𝑒𝑙𝑑𝑠 𝑁𝑖 2+ + 𝐶𝑑 → 𝐶𝑑 + 𝑁𝑖 2+ Module No. 7 – Corrosion and Control 7 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G From Table 7.1, the half-cell potentials for cadmium and nickel are, respectively, 0.403 and 0.250 V. Therefore, from Equation 17.18, ∆V = V0 Ni - V0 Cd = -0.250 V - (-0.403 V) = +0.153 V (b) For this portion of the problem, Equation 17.20 must be used, because the half-cell solution concentrations are no longer 1 M. At this point, it is necessary to make a calculated guess as to which metal species is oxidized (or reduced). This choice will either be affirmed or refuted on the basis of the sign of ∆V at the conclusion of the computation. For the sake of argument, let us assume that in contrast to part (a), nickel is oxidized and cadmium reduced according to 𝑦𝑖𝑒𝑙𝑑𝑠 𝐶𝑑 +2 + 𝑁𝑖 → 𝐶𝑑 + 𝑁𝑖 +2 (17.22) 0 0 𝑅𝑇 [𝑁𝑖 2+ ] Thus, ∆𝑉 = (𝑉𝐶𝑑 − 𝑉𝑁𝑖 )− 𝑙𝑛 2+ 𝑛ℱ [𝐶𝑑 ] 0.0592 10−3 = −0.403𝑉 − (−0.250𝑉) − log( ) 2 0.50 = −0.073𝑉 Because ΔV is negative, the spontaneous reaction direction is the opposite to that of Equation 17.22. That is, cadmium is oxidized and nickel is reduced. THE GALVANIC SERIES Even though Table 7.1 was generated under highly idealized conditions and has limited utility, it nevertheless indicates the relative reactivities of the metals. A more realistic galvanic series and practical ranking is provided by the galvanic series, Table 7.2. This represents the relative reactivities of a number of metals and commercial alloys in seawater. The alloys near the top are cathodic and unreactive, whereas those at the bottom are most anodic; no voltages are provided. Comparison of the standard emf and the galvanic series reveals a high degree of correspondence between the relative positions of the pure base metals. Most metals and alloys are subject to oxidation or corrosion to one degree or another in a wide variety of environments that is, they are more stable in an ionic state than as metals. In thermodynamic terms, there is a net decrease in free energy in going from metallic to oxidized states. Consequently, essentially all metals occur in nature as compounds for example, oxides, hydroxides, carbonates, silicates, sulfides, and sulfates. Two exceptions are the noble metals gold and platinum, for which oxidation in most environments is not favorable, and, therefore, they may exist in nature in the metallic state. Module No. 7 – Corrosion and Control 8 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G CORROSION RATES The half-cell potentials listed in Table 7.1 are thermodynamic parameters that relate to systems at equilibrium. For example, for the discussions pertaining to Figures 7.2 and 7.3, it was tacitly assumed that there was no current flow through the external circuit. Real corroding systems are not at equilibrium; there is a flow of electrons from anode to cathode (corresponding to the short- circuiting of the electrochemical cells in Figures 7.2 and 7.3), which means that the half-cell potential parameters (Table 7.1) cannot be applied. TABLE 7.2 THE GALVANIC SERIES Furthermore, these half-cell potentials represent the magnitude of a driving force, or the tendency for the occurrence of the particular half-cell reaction. However, although these potentials may be used to determine spontaneous reaction directions, they provide no information on corrosion rates. That is, even though a ΔV potential computed for a specific corrosion situation using Equation 7.20 is a relatively large positive number, the reaction may occur at only an insignificantly slow rate. From an engineering perspective, we are interested in predicting the rates at which systems corrode; this requires the use of other parameters, as discussed next. The corrosion rate, or the rate of material removal as a consequence of the chemical action, is an important corrosion parameter. This may be expressed as the corrosion penetration rate (CPR), or the thickness loss of material per unit of time. The formula for this calculation is 𝐾𝑊 𝐶𝑃𝑅 = 𝜌𝐴𝑡 (7.23) Module No. 7 – Corrosion and Control 9 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G where W is the weight loss after exposure time t; 𝜌 and A represent the density and exposed specimen area, respectively, and K is a constant, its magnitude depending on the system of units used. The CPR is conveniently expressed in terms of either mils per year (mpy) or millimeters per year (mm/yr). In the first case, K = 534 to give CPR in mpy (where 1 mil = 0.001 in.), and W,𝜌, A, and t are specified in units of milligrams, grams per cubic centimeter, square inches, and hours, respectively. In the second case, K=87.6 for mm/yr, and units for the other parameters are the same as for mils per year, except that A is given in square centimeters. For most applications, a corrosion penetration rate less than about 20 mpy (0.50 mm/yr) is acceptable. In as much as there is an electric current associated with electrochemical corrosion reactions, we can also express corrosion rate in terms of this current, or, more specifically, current density— that is, the current per unit surface area of material corroding which is designated i. The rate r, in units of mol/m2-s, is determined using the expression: 𝑖 𝑟= (7.24) 𝑛ℱ Where, again, n is the number of electrons associated with the ionization of each metal Atom and F is 96,500 C/mol. PREDICTION OF CORROSION RATES POLARIZATION Consider the standard Zn/H2 electrochemical cell shown in Figure 7.5, which has been short-circuited such that oxidation of zinc and reduction of hydrogen occurs at their respective electrode surfaces. The potentials of the two electrodes are not at the values determined from Table 7.1 because the system is now a non-equilibrium one. The displacement of each electrode potential from its equilibrium value is termed polarization, and the magnitude of this displacement is the overvoltage, normally represented by the symbol h. Overvoltage is expressed in terms of plus or minus volts (or millivolts) relative to the equilibrium potential. For example, suppose that the zinc electrode in Figure 17.5 has a potential of 0.621 V after it has been connected to the platinum electrode. The equilibrium potential is 0.763 V (Table 17.1), and, therefore, There are two types of polarization activation and concentration. We now discuss their mechanisms because they control the rate of electrochemical reactions. Activation Polarization All electrochemical reactions consist of a sequence of steps that occur in series at the interface between the metal electrode and the electrolyte solution. Activation polarization refers to the condition in which the reaction rate is controlled by the one step in the series that occurs at the slowest rate. The term activation is applied to this type of polarization because an activation energy barrier is associated with this slowest, rate limiting step. Module No. 7 – Corrosion and Control 10 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G Figure 7.5 Electrochemical cell consisting of standard zinc and hydrogen electrodes that has been short- circuited. Figure 7.6 Schematic representation of possible steps in the hydrogen reduction reaction, the rate of which is controlled by activation polarization. (From M. G. Fontana, Corrosion Engineering, and 3rd edition. Copyright © 1986 by McGraw-Hill Book Company. Reproduced with permission.) Figure 7.7 For a hydrogen electrode, plot of activation polarization overvoltage versus logarithm of current density for both oxidation and reduction reactions. (Adapted from M. G. Fontana, Corrosion Engineering, 3rd edition. Copyright © 1986 by McGraw-Hill Book Company. Reproduced with permission.) Module No. 7 – Corrosion and Control 11 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G CONCENTRATION POLARIZATION Concentration polarization exists when the reaction rate is limited by diffusion in the solution. For example, consider again the hydrogen evolution reduction reaction. When the reaction rate is low and/or the concentration of H+ is high, there is always an adequate supply of hydrogen ions available in the solution at the region near the electrode interface (Figure 7.8a). However, at high rates and/or low H concentrations, a depletion zone may be formed in the vicinity of the interface because the H ions are not replenished at a rate sufficient to keep up with the reaction (Figure 7.8b). Thus, diffusion of H to the interface is rate controlling, and the system is said to be concentration polarized. Figure 7.8 For hydrogen reduction, schematic representations of the H + distribution in the vicinity of the cathode for (a) low reaction rates and/or high concentrations and (b) high reaction rates and/or low concentrations, where a depletion zone is formed that gives rise to concentration polarization.(Adapted from M. G. Fontana, Corrosion Engineering, 3rd edition. Copyright © 1986 by McGraw-Hill Book Company. Reproduced with permission. EXAMPLE PROBLEM 7.2 Rate of Oxidation Computation Zinc experiences corrosion in an acid solution according to the reaction Zn + 2H+ Zn2+ + H2 The rates of both oxidation and reduction half-reactions are controlled by activation polarization. (a) Compute the rate of oxidation of Zn (in mol/cm 2-s), given the following activation polarization data: For Zn For Hydrogen V(Zn/Zn2+) = - 0.763 V (H+/H2) = 0 V V i0 = 10-7 A/cm2 i0 = 10-10 A/cm2 b = +0.09 b = -0.08 (b) Compute the value of the corrosion potential. Module No. 7 – Corrosion and Control 12 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G Solution (a) To compute the rate of oxidation for Zn, it is first necessary to establish relationships in the form of Equation 7.25 for the potential of both oxidation and reduction reactions. Next, we set these two expressions equal to one another, and then we solve for the value of i that is the corrosion current density, iC. Finally, the corrosion rate may be calculated using Equation 7.24. The two potential expressions are as follows: For hydrogen reduction, 𝑖 𝑉𝐻 = 𝑉(𝐻 + /𝐻2) + 𝛽𝐻 log( ) 𝑖0𝐻 and for Zn oxidation, 𝑖 𝑉𝑍𝑛 = 𝑉(𝑍𝑛/𝑍𝑛2+) + 𝛽𝑍𝑛 log( ) 𝑖0𝑍𝑛 Now, setting VH = VZn leads to Solving for log i (i.e., log iC) leads to 𝑖 𝑖 𝑉𝐻 = 𝑉(𝐻 +/𝐻2) + 𝛽𝐻 log( ) = 𝑉𝑍𝑛 = 𝑉(𝑍𝑛/𝑍𝑛2+) + 𝛽𝑍𝑛 log( ) 𝑖0𝐻 𝑖0𝑍𝑛 1 𝑖 𝑖 𝑙𝑜𝑔𝑖𝑐 = ( ) [𝑉 𝐻 + − 𝑉 𝑍𝑛 − 𝛽𝐻 log( ) − 𝛽𝑍𝑛 log( )] 𝛽𝑍𝑛 − 𝛽𝐻 (𝐻 ) ( 𝑍𝑛2+ ) 𝑖0𝐻 𝑖0𝑍𝑛 2 1 =[ ] [0 − (−0.08)(𝑙𝑜𝑔10−10 )] + (0.09)(𝑙𝑜𝑔10−7 )] 0.09 − (−0.08) =-3.924 Or ic= 10-3.924= 1.19x10-4 A/cm2 From Equation 7.24, 1.9𝑥10−4 𝐶/𝑠 ∗ 𝑐𝑚2 = = 6.17𝑥10−10 𝑚𝑜𝑙/𝑐𝑚2. 𝑠 2 ∗ 96500 𝐶/𝑚𝑜𝑙 𝑖 𝑐 𝑛Ϝ (b) Now it becomes necessary to compute the value of the corrosion potential VC. This is possible by using either of the preceding equations for VH or VZn and substituting for i the value determined previously for iC. Thus, using the VH expression yields 𝑖𝐶 𝑉𝑐 = 𝑉(𝐻 +/H) + 𝛽𝐻 log( ) 𝑖0𝐻 1.19𝑥10−4 𝐴 = 0 + (−0.08𝑉) log ( 𝑐𝑚2 ) = −0.486𝑉 10−10 𝐴 𝑐𝑚2 Module No. 7 – Corrosion and Control 13 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G PASSIVITY Under particular environmental conditions, some normally active metals and alloys, lose passivity their chemical reactivity and become extremely inert. This phenomenon, termed passivity, is displayed by chromium, iron, nickel, titanium, and many of their alloys. It is believed that this passive behavior results from the formation of a highly adherent and very thin oxide film on the metal surface, which serves as a protective barrier to further corrosion. Stainless steels are highly resistant to corrosion in a rather wide variety of atmospheres as a result of passivation. They contain at least 11% chromium, which as a solid-solution alloying element in iron, minimizes the formation of rust; instead, a protective surface film forms in oxidizing atmospheres. (Stainless steels are susceptible to corrosion in some environments and therefore are not always “stainless.”) Aluminum is highly corrosion resistant in many environments because it also passivates. If damaged, the protective film normally re-forms very rapidly. However, a change in the character of the environment (e.g., alteration in the concentration of the active corrosive species) may cause a passivated material to revert to an active state. Subsequent damage to a preexisting passive film could result in a substantial increase in corrosion rate, by as much as 100,000 times. This passivation phenomenon may be explained in terms of polarization potential- log current density curves discussed in the preceding section. ENVIRONMENTAL EFFECTS The variables in the corrosion environment, which include fluid velocity, temperature, and composition, can have a decided influence on the corrosion properties of the materials that are in contact with it. In most instances, increasing fluid velocity enhances the rate of corrosion due to erosive effects, as discussed later in the chapter. The rates of most chemical reactions rise with increasing temperature; this also holds for most corrosion situations. Increasing the concentration of the corrosive species (e.g., H+ ions in acids) in many situations produces a more rapid rate of corrosion. However, for materials capable of passivation, raising the corrosive content may result in an active-to-passive transition, with a considerable reduction in corrosion. Cold working or plastically deforming ductile metals is used to increase their strength; however, a cold-worked metal is more susceptible to corrosion than the same material in an annealed state. For example, deformation processes are used to shape the head and point of a nail; consequently, these positions are anodic with respect to the shank region. Thus, differential cold working on a structure should be a consideration when a corrosive environment may be encountered during service. FORMS OF CORROSION It is convenient to classify corrosion according to the manner in which it is manifest. Metallic corrosion is sometimes classified into eight forms: uniform, galvanic, crevice, pitting, intergranular, selective leaching, erosion-corrosion, and stress corrosion. The causes and means of prevention of each of these forms are discussed briefly. In addition, we have elected to discuss the topic of hydrogen embrittlement in this section. Hydrogen embrittlement is, in a strict sense, a type of failure rather than a form of corrosion; however, it is often produced by hydrogen that is generated from corrosion reactions. Module No. 7 – Corrosion and Control 14 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G UNIFORM ATTACK Uniform attack is a form of electrochemical corrosion that occurs with equivalent intensity over the entire exposed surface and often leaves behind a scale or deposit. In a microscopic sense, the oxidation and reduction reactions occur randomly over the surface. Familiar examples include general rusting of steel and iron and the tarnishing of silverware. This is probably the most common form of corrosion. It is also the least objectionable because it can be predicted and designed for with relative ease. GALVANIC CORROSION Galvanic corrosion occurs when two metals or alloys having different compositions are electrically coupled while exposed to an electrolyte. This is the type of corrosion or dissolution that was described in Section 7.2. The less noble or more reactive metal in the particular environment experiences corrosion; the more inert metal, the cathode, is protected from corrosion. As examples, steel screws corrode when in contact with brass in a marine environment, and if copper and steel tubing are joined in a domestic water heater, the steel corrodes in the vicinity of the junction. Depending on the nature of the solution, one or more of the reduction reactions, Equations 7.3 through 7.7, occurs at the surface of the cathode material. Figure 7.9 shows galvanic corrosion. Figure 7.9 Photograph showing galvanic corrosion around the inlet of a single-cycle bilge pump that is found on fishing vessels. Corrosion occurred between a magnesium shell that was cast around a steel core. The galvanic series in Table 7.2 indicates the relative reactivities in seawater of a number of metals and alloys. When two alloys are coupled in seawater, the one lower in the series experiences corrosion. Some of the alloys in the table are grouped in brackets. Generally the base metal is the same for these bracketed alloys, and there is little danger of corrosion if alloys within a single bracket are coupled. It is also worth noting from this series that some alloys are listed twice (e.g., nickel and the stainless steels), in both active and passive states. The rate of galvanic attack depends on the relative anode-to-cathode surface areas that are exposed to the electrolyte, and the rate is related directly to the cathode–anode area ratio—that is, for a given cathode area, a smaller anode corrodes more rapidly than a larger one because corrosion rate Module No. 7 – Corrosion and Control 15 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G depends on current density (Equation 7.4)—the current per unit area of corroding surface—and not simply the current. Thus, a high current density results for the anode when its area is small relative to that of the cathode. A number of measures may be taken to reduce the effects of galvanic corrosion significantly including the following: 1. If coupling of dissimilar metals is necessary, choose two that are close together in the galvanic series. 2. Avoid an unfavorable anode-to-cathode surface area ratio; use an anode area as large as possible. 3. Electrically insulate dissimilar metals from each other. 4. Electrically connect a third, anodic metal to the other two; this is a form of cathodic protection. Figure 7.10 On this plate, which was immersed in seawater, crevice corrosion has occurred at the regions that were covered by washers. (Photograph courtesy of LaQue Center for Corrosion Technology, Inc.) CREVICE CORROSION Electrochemical corrosion may also occur as a consequence of concentration differences of ions or dissolved gases in the electrolyte solution and between two regions of the same metal piece. For such a concentration cell, corrosion occurs in the locale that has the lower concentration. A good example of this type of corrosion occurs in crevices and recesses or under deposits of dirt or corrosion products where the solution becomes stagnant and there is localized depletion of dissolved oxygen. Corrosion preferentially occurring at these positions is called crevice corrosion (Figure 7.15). The crevice must be wide enough for the solution to penetrate yet narrow enough for stagnancy; usually the width is several thousandths of an inch. The proposed mechanism for crevice corrosion is illustrated in Figure 7.11. After oxygen has been depleted within the crevice, oxidation of the metal occurs at this Figure 7.11 Schematic illustration of the mechanism of crevice corrosion between two riveted sheets. (From M. G. Fontana, Corrosion Engineering, 3rd edition. Copyright © 1986 by McGraw-Hill Book Company. Reproduced with permission.) Module No. 7 – Corrosion and Control 16 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G PITTING Pitting is another form of much localized corrosion attack in which small pits or holes form. They ordinarily penetrate from the top of a horizontal surface downward in a nearly vertical direction. It is an extremely insidious type of corrosion, often going undetected and with very little material loss until failure occurs. An example of pitting corrosion is shown in Figure 17.12. The mechanism for pitting is probably the same as for crevice corrosion, in that oxidation occurs within the pit itself, with complementary reduction at the surface. It is supposed that gravity causes the pits to grow downward, the solution at the pit tip becoming more concentrated and dense as pit growth progresses. INTERGRANULAR CORROSION As the name suggests, intergranular corrosion occurs preferentially along grain boundaries for some alloys and in specific environments. The net result is that a macroscopic specimen disintegrates along its grain boundaries. This type of corrosion is especially prevalent in some stainless steels. When heated to temperatures between 500C and 800C (950F and 1450F) for sufficiently long time periods, these alloys become sensitized to intergranular attack. It is believed that this heat treatment permits the formation of small precipitate particles of chromium carbide (Cr23C6) by reaction between the chromium and carbon in the stainless steel. These particles form along the grain boundaries, as illustrated in Figure 7.12. Figure 7.13 shows this type of intergranular corrosion. Stainless steels may be protected from intergranular corrosion by the following measures: (1) subjecting the sensitized material to a high-temperature heat treatment in which all the chromium carbide particles are re dissolved, (2) lowering the carbon content below 0.03 wt% C so that carbide formation is minimal, and (3) alloying the stainless steel with another metal such as niobium or titanium, hich has a greater tendency to form carbides than does chromium so that the Cr remains in solid solution. Figure 7.12 The pitting of a 304 stainless steel plate by an acid–chloride solution. (Photograph courtesy of Mars G. Fontana. From M. G. Fontana, Corrosion Engineering, 3rd edition. Copyright © 1986 by McGraw- Hill Book Company. Reproduced with permission.) Figure 7.13 Schematic illustration of chromium carbide particles that have precipitated along grain boundaries in stainless steel, and the attendant zones of chromium depletion. Module No. 7 – Corrosion and Control 17 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G SELECTIVE LEACHING Selective leaching is found in solid solution alloys and occurs when one element or constituent is preferentially removed as a consequence of corrosion processes. The most common example is the dezincification of brass, in which zinc is selectively leached from a copper–zinc brass alloy. The mechanical properties of the alloy are significantly impaired because only a porous mass of copper remains in the region that has been dezincified. In addition, the material changes from yellow to a red or copper color. Selective leaching may also occur with other alloy systems in which aluminum, iron, cobalt, chromium, and other elements are vulnerable to preferential removal. Figure 7.14 Weld decay in a stainless steel. The regions along which the grooves have formed were sensitized as the weld cooled. (From H. H. Uhlig and R. W. Revie, Corrosion and Corrosion Control, 3rd edition, Fig. 2, p. 307. Copyright © 1985 by John Wiley & Sons, Inc. Reprinted by permission of John Wiley & Sons, Inc.) EROSION CORROSION Erosion corrosion arises from the combined action of chemical attack and mechanical abrasion or wear as a consequence of fluid motion. Virtually all metal alloys, to one degree or another, are susceptible to erosion–corrosion. It is especially harmful to alloys that passivate by forming a protective surface film; the abrasive action may erode away the film, leaving exposed a bare metal surface. Erosion corrosion is commonly found in piping, especially at bends, elbows, and abrupt changes in pipe diameter positions where the fluid changes direction or flow suddenly becomes turbulent. Propellers, turbine blades, valves, and pumps are also susceptible to this form of corrosion. Figure 7.15 illustrates the impingement failure of an elbow fitting. One of the best ways to reduce erosion corrosion is to change the design to eliminate fluid turbulence and impingement effects. Other materials may also be used that inherently resist erosion. Furthermore, removal of particulates and bubbles from the solution lessens its ability to erode. STRESS CORROSION Stress corrosion, sometimes termed stress corrosion cracking, results from the combined action of an applied tensile stress and a corrosive environment; both influences are necessary. In fact, some materials that are virtually inert in a particular corrosive medium become susceptible to this form of corrosion when a stress is applied. Small cracks form and then propagate in a direction perpendicular to the stress, with the result that failure may eventually occur. Failure behavior is characteristic of that for a brittle material, even though the metal alloy is intrinsically Module No. 7 – Corrosion and Control 18 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G ductile. Furthermore, cracks may form at relatively low stress levels, significantly below the tensile strength. Most alloys are susceptible to stress corrosion in specific environments, especially at moderate stress levels. For example, most stainless steels stress corrode in solutions containing chloride ions, whereas brasses are especially vulnerable when exposed to ammonia. Figure 7.17 is a photomicrograph showing an example of intergranular stress corrosion cracking in brass. The stress that produces stress corrosion cracking need not be externally applied; it may be a residual one that results from rapid temperature changes and uneven contraction or occur for two-phase alloys in which each phase has a different coefficient of expansion. Also, gaseous and solid corrosion products that are entrapped internally can give rise to internal stresses. Probably the best measure to take to reduce or completely eliminate stress corrosion is to lower the magnitude of the stress. This may be accomplished by reducing the external load or increasing the cross- sectional area perpendicular to the applied stress. Furthermore, an appropriate heat treatment may be used to anneal out any residual thermal stresses. HYDROGEN EMBRITTLEMENT Various metal alloys, specifically some steels, experience a significant reduction in ductility and tensile strength when atomic hydrogen (H) penetrates into the material. Fig. 7.15 Fig. 7.16 Figure 7.15 Impingement failure of an elbow that was part of a steam condensate line. (Photograph courtesy of Mars G. Fontana. From M. G. Fontana, Corrosion Engineering, 3rd edition. Copyright © 1986 by McGrawHill Book Company. Reproduced with permission.) Figure 7.16 A bar of steel bent into a horseshoe shape using a nutand-bolt assembly. While immersed in seawater, stress corrosion cracks formed along the bend at those regions where the tensile stresses are the greatest. (Photograph courtesy of F. L. LaQue. From F. L. LaQue, Marine Corrosion, Causes and Prevention. Copyright © 1975 by John Wiley & Sons, Inc. Reprinted by permission of John Wiley & Sons, Inc.) This phenomenon is aptly referred to as hydrogen embrittlement; the terms hydrogen induced cracking and hydrogen stress cracking are sometimes also used. Strictly speaking, hydrogen embrittlement is a type of failure; in response to applied or residual tensile stresses, brittle fracture occurs catastrophically as cracks grow and rapidly propagate. Hydrogen in its atomic form (H as opposed to the molecular form, H2) diffuses interstitially through the crystal lattice, and Module No. 7 – Corrosion and Control 19 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G concentrations as low as several parts per million can lead to cracking. Furthermore, hydrogen- induced cracks are most often Trans granular, although intergranular fracture is observed for some alloy systems. A number of mechanisms have been proposed to explain hydrogen embrittlement; most are based on the interference of dislocation motion by the dissolved hydrogen. Hydrogen embrittlement is similar to stress corrosion, in that a normally ductile metal experiences brittle fracture when exposed to both a tensile stress and a corrosive atmosphere. However, these two phenomena may be distinguished on the basis of their interactions with applied electric currents. Whereas cathodic protection reduces or causes a cessation of stress corrosion, it may, however, lead to the initiation or enhancement of hydrogen embrittlement. Figure 7.17 Photomicrograph showing intergranular stress corrosion cracking in brass. (From H. H. Uhlig and R. W. Revie, Corrosion and Corrosion Control, 3rd edition, Fig. 5, p. 335. Copyright 1985 by John Wiley & Sons, Inc. Reprinted by permission of John Wiley & Sons, Inc.) For hydrogen embrittlement to occur, some source of hydrogen must be present, as well as the possibility for the formation of its atomic species. Situations in which these conditions are met include the following: pickling3 of steels in sulfuric acid; electroplating; and the presence of hydrogen-bearing atmospheres (including water vapor) at elevated temperatures such as during welding and heat treatments. Also, the presence of what are termed poisons such as sulfur (i.e., H2S) and arsenic compounds accelerates hydrogen embrittlement; these substances retard the formation of molecular hydrogen and thereby increase the residence time of atomic hydrogen on the metal surface. Hydrogen sulfide, probably the most aggressive poison, is found in petroleum fluids, natural gas, oil-well brines, and geothermal fluids. High-strength steels are susceptible to hydrogen embrittlement, and increasing strength tends to enhance the material’s susceptibility. Martensitic steels are especially vulnerable to this type of failure; bainitic, ferritic, and spheroiditic steels are more resilient. Furthermore, FCC alloys (austenitic stainless steels and alloys of copper, aluminum, and nickel) are relatively resistant to hydrogen embrittlement, mainly because of their inherently high ductilities. However, strain hardening these alloys enhances their susceptibility to embrittlement. Techniques commonly used to reduce the likelihood of hydrogen embrittlement include reducing the tensile strength of the alloy via a heat treatment, removing the source of hydrogen, “baking” the alloy at an elevated temperature to drive out any dissolved hydrogen, and substituting a more embrittlement-resistant alloy. Module No. 7 – Corrosion and Control 20 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G CATHODIC PROTECTION One of the most effective means of corrosion prevention is cathodic protection; it can be used for all eight different forms of corrosion as discussed earlier and may, in some situations, completely stop corrosion. Again, oxidation or corrosion of a metal M occurs by the generalized reaction, M Mn+ + ne- Cathodic protection simply involves supplying, from an external source, electrons to the metal to be protected, making it a cathode; the preceding reaction is thus forced in the reverse (or reduction) direction. One cathodic protection technique employs a galvanic couple: the metal to be protected is electrically connected to another metal that is more reactive in the particular environment. The latter experiences oxidation and, upon giving up electrons, protects the first metal from corrosion. The oxidized metal is often called a sacrificial anode, and magnesium and zinc are commonly used because they lie at the anodic end of the galvanic series. This form of galvanic protection for structures buried in the ground is illustrated in Figure 7.18a. The process of galvanizing is simply one in which a layer of zinc is applied to the surface of steel by hot dipping. In the atmosphere and most aqueous environments, zinc is anodic to and will thus cathodically protect the steel if there is any surface damage. Any corrosion of the zinc coating will proceed at an extremely slow rate because the ratio of the anode-to-cathode surface area is quite large. For another method of cathodic protection, the source of electrons is an impressed current from an external dc power source, as represented in Figure 7.18b for an underground tank. The negative terminal of the power source is connected to the structure to be protected. The other terminal is joined to an inert anode (often graphite), which, in this case, is buried in the soil; high-conductivity backfill material provides good electrical contact between the anode and the surrounding soil. A current path exists between the cathode and the anode through the intervening soil, completing the electrical circuit. Cathodic protection is especially useful in preventing corrosion of water heaters, underground tanks and pipes, and marine equipment. Figure 7.18 Cathodic protection of (a) an underground pipeline using a magnesium sacrificial anode and (b) an underground tank using an impressed current. (From M. G. Fontana, Corrosion Engineering, 3rd edition. Copyright © 1986 by McGraw-Hill Book Company. Reproduced with permission.) Module No. 7 – Corrosion and Control 21 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G Figure 7.19 Galvanic protection of steel as provided by a coating of zinc. The cathode and the anode through the intervening soil, completing the electrical circuit. Cathodic protection is especially useful in preventing corrosion of water heaters, underground tanks and pipes, and marine equipment. SCALE TYPES Rate of oxidation (i.e., the rate of film thickness increase) and the tendency of the film to protect the metal from further oxidation are related to the relative volumes of the oxide and metal. The ratio of these volumes, termed the Pilling–Bedworth ratio, may be determined from the following expression: 𝐴0 𝜌𝑀 𝑃 − 𝐵 𝑟𝑎𝑡𝑖𝑜 = 𝐴𝑀 𝜌𝑂 Where AO is the molecular (or formula) weight of the oxide, A M is the atomic weight of the metal, and rO and rM are the oxide and metal densities, respectively. For metals having P–B ratios less than unity, the oxide film tends to be porous and un protective Alternatively, electron holes and vacancies may diffuse instead of electrons and ions. 6 For other than divalent metals, Equation 7.32 becomes 𝐴0 𝜌𝑀 𝑃 − 𝐵 𝑟𝑎𝑡𝑖𝑜 = 𝑎𝐴𝑀 𝜌𝑂 Where a is the coefficient of the metal species for the overall oxidation reaction. Module No. 7 – Corrosion and Control 22 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G TABLE 7.3 PILLING–BEDWORTH RATIOS FOR A NUMBER OF METALS/METAL OXIDES Table 7.3 presents P–B ratios for metals that form protective coatings and for those that do not. Note from these data that protective coatings normally form for metals having P–B ratios between 1 and 2, whereas non protective ones usually result when this ratio is less than 1 or greater than about 2. In addition to the P–B ratio, other factors also influence the oxidation resistance imparted by the film; these include a high degree of adherence between film and metal, comparable coefficients of thermal expansion for metal and oxide, and, for the oxide, a relatively high melting point and good high-temperature plasticity. CORROSION OF CERAMIC MATERIALS Ceramic materials, being compounds between metallic and nonmetallic elements, may be thought of as having already been corroded. Thus, they are exceedingly immune to corrosion by almost all environments, especially at room temperature. Corrosion of ceramic materials generally involves simple chemical dissolution, in contrast to the electrochemical processes found in metals, as described previously. Ceramic materials are frequently used because of their resistance to corrosion. Glass is often used to contain liquids for this reason. Refractory ceramics must not only withstand high temperatures and provide thermal insulation, but also, in many instances, must resist high-temperature attack by molten metals, salts, slags, and glasses. Some of the new technology schemes for converting energy from one form into another that is more useful require relatively high temperatures, corrosive atmospheres, and pressures above the ambient pressure. Ceramic materials are much better suited to withstand most of these environments for reasonable time periods than are metals. Module No. 7 – Corrosion and Control 23 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G TABLE 7.4 RESISTANCE TO DEGRADATION BY VARIOUS ENVIRONMENTS FOR SELECTED PLASTIC MATERIALS TABLE 7.5 RESISTANCE TO DEGRADATION BY VARIOUS ENVIRONMENTS FOR SELECTED ELASTOMERIC MATERIALS Module No. 7 – Corrosion and Control 24 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G CHEMICAL REACTION EFFECTS Oxygen, ozone, and other substances can cause or accelerate chain scission as a result of chemical reaction. This effect is especially prevalent in vulcanized rubbers that have doubly bonded carbon atoms along the backbone molecular chains and that are exposed to ozone (O 3), an atmospheric pollutant. One such scission reaction may be represented by Where the chain is severed at the point of the double bond; R and Rrepresent groups of atoms that are unaffected during the reaction. Typically, if the rubber is in an unstressed state, an oxide film will form on the surface, protecting the bulk material from any further reaction. However, when these materials are subjected to tensile stresses, cracks and crevices form and grow in a direction perpendicular to the stress; eventually, rupture of the material may occur. This is why the sidewalls on rubber bicycle tires develop cracks as they age. Apparently these cracks result from large numbers of ozone-induced scissions. Chemical degradation is a particular problem for polymers used in areas with high levels of air pollutants such as smog and ozone. THERMAL EFFECTS Thermal degradation corresponds to the scission of molecular chains at elevated temperatures; as a consequence, some polymers undergo chemical reactions in which gaseous species are produced. These reactions are evidenced by a weight loss of material; a polymer’s thermal stability is a measure of its resilience to this decomposition. For example, the magnitude of the C—F bond is greater than that of the C—H bond, which in turn is greater than that of the C—Cl bond. The fluorocarbons, having C—F bonds, are among the most thermally resistant polymeric materials and may be used at relatively high temperatures. However, because of the weak C—Cl bond, when poly(vinyl chloride) is heated to 200C for even a few minutes it discolors and gives off large amounts of HCl, which accelerates continued decomposition. Stabilizers such as ZnO can react with the HCl, providing increased thermal stability for poly(vinyl chloride). Some of the most thermally stable polymers are the ladder polymers. For example, the ladder polymer having the structure Module No. 7 – Corrosion and Control 25 E N G G 4 1 2 : M A T E R I A L S S C I E N C E A N D E N G I N E E R I N G is so thermally stable that a woven cloth of this material can be heated directly in an open flame with no degradation. Polymers of this type are used in place of asbestos for high-temperature gloves. WEATHERING Many polymeric materials serve in applications that require exposure to outdoor conditions. Any resultant degradation is termed weathering, which may be a combination of several different processes. Under these conditions, deterioration is primarily a result of oxidation, which is initiated by ultraviolet radiation from the sun. Some polymers, such as nylon and cellulose, are also susceptible to water absorption, which produces a reduction in their hardness and stiffness. Resistance to weathering among the various polymers is quite diverse. The fluorocarbons are virtually inert under these conditions; some materials, however, including poly (vinyl chloride) and polystyrene, are susceptible to weathering. Reference: Materials Science and Engineering: An Introduction, 10th Edition by William D. Callister Jr and David G. Rethwisch Module No. 7 – Corrosion and Control 26