Introduction to the Fundamentals of Corrosion

THE SIGNIFICANT TECHNICAL CHALLENGES and the high cost directly related to corrosion provide strong incentives for engineers and other technical personnel to develop a firm grasp on the fundamental bases of corrosion. Understanding the fundamentals of corrosion is necessary not only for identifying corrosion mechanisms (a significant achievement by itself), but also for preventing corrosion by appropriate corrosion protection means and for predicting the corrosion behavior of metallic materials in service conditions. Understanding the mechanisms of corrosion is the key to the development of a knowledge-based design of corrosion resistant alloys and to the prediction of the long-term behavior of metallic materials in corrosive environments. Two major areas are usually distinguished in the corrosion of metals and alloys. The first area is where the metal or alloy is exposed to a liquid electrolyte, usually water, and thus typically called aqueous corrosion. The second area is where corrosion takes place in a gaseous environment, often called oxidation, high- temperature oxidation, or high-temperature corrosion, and called gaseous corrosion here. These two areas have been (and still are sometimes) referred to as wet corrosion and dry corrosion. This distinction finds its origin (and its justification) in some fundamental differences in the mechanisms, in particular the electrochemical nature of reactions occurring in aqueous solution (or in a nonaqueous electrolyte), as compared to the formation of thick oxide layers in air or other oxidizing atmospheres, at high temperature with fast transport processes by solidstate diffusion through a growing oxide. The separation between the two areas, however, should not be overemphasized, because there are also similarities and analogies, for example: · The initial stages of reaction involve the adsorption of chemical species on the metal surface that can be described by the Gibbs equation for both liquid and gaseous environments. · The nucleation and growth phenomena of oxide layers and other compounds · The use of surface analytical techniques The fundamental aspects of aqueous and gaseous corrosion are addressed in this first Section of the Handbook. Corrosion of metallic materials is generally detrimental and must be prevented, but if it is well understood and controlled, it can also be used in a powerful and constructive manner for electrochemical production of fine patterns on metal as well as on semiconductor surfaces. These constructive purposes also include electrochemical machining (down to the micro- or even the nanoscale), electrochemical and chemicalmechanical polishing, and anodes for batteries and fuel cells. These topics are also addressed in this Section. THE DRIVING FORCE of corrosion is the lowering of energy associated with the oxidation of a metal. Thermodynamics examines and quantifies this driving force. It predicts if reactions can or cannot occur (i.e., if the metal will corrode or be stable). It does not predict at what rate these changes can or will occur: this is the area of kinetics. However, knowing from thermodynamics what reactions are possible is a necessary step in the attempt to understand, predict, and control corrosion. Ions that are present in the solution are charged because of the loss or gain of electrons. The positive charged ions (cations) and negative charged ions (anions) also have an electric field associated with them. The solvent (water) molecules act as small dipoles; therefore, they are also attracted to the charged ions and align themselves in the electric field established by the charge of the ion. Because the electric field is strongest close to the ion, some water molecules reside very close to an ionic species in solution. The attraction is great enough that these water molecules travel with the ion as it moves through the solvent. The tightly bound water molecules are referred to as the primary water sheath of the ion. The electric field is weaker at distances outside the primary water sheath, but it still disturbs the polar water molecules as the ion passes through the solution. The water molecules that are disturbed as the ion passes, but do not move with the ion, are usually referred to as the secondary water sheath. Figure 1 shows a representation of the primary and secondary solvent molecules for a cation in water. Because of their smaller size relative to anions, cations have a stronger electric field close to the ion and more water molecules are associated in their primary water sheath. Anions have few, if any, primary water molecules. A detailed description of the hydration of ions in solution is given in Ref 1 ONE OF THE IMPORTANT FEATURES of the electrified interface between the electrode and the electrolyte in the aqueous corrosion of metals is the existence of a potential difference across the double layer, which leads to the definition of the electrode potential. The electrode potential is one of the most important parameters in both the thermodynamics and the kinetics of corrosion. The fundamentals of electrode potentials are discussed in this article. Examples of the calculations of the potential at equilibrium are given in the article “Potential versus pH (Pourbaix) Diagrams” in this Section of the Volume. electrodes are left unchanged. Hence, the magnitude and the sign of the cell voltage at equilibrium (emf) depend only on the couple of half-reactions involved. From the thermodynamic convention, the free energy change of a spontaneous cell reaction, which liberates energy, is negative. If the emf is the potential of electrode 2 minus the potential of electrode 1 (ΔE = E2 - E1), and if ΔrG designates the free energy of the cell reaction written in the sense of Eq 13, that is, reduction at electrode 2 (Eq 12) and oxidation at electrode 1 (Eq 11), the relation between ΔrG and ΔE is: ΔrG = -nFΔE (Eq 14) where n is the number of electrons exchanged in both half-reactions, and F is the Faraday constant (96,487 coulombs) equal to the charge of 1 mole of electrons. Using Eq 3 and 4, Eq 14 may be rewritten as: (Eq 15) Compared to a chemical equilibrium where, at a given temperature, the ratio is equal to a constant Keq (Eq 6), the electrochemical cell at equilibrium is a system with one more degree of freedom, because either the ratio of activities or the cell emf can impose on it. If the cell reaction occurs under conditions in which the reactants and products are in their standard states, the equation becomes: (Eq 16) where ΔE0 is the standard cell emf. The Hydrogen Potential Scale. The absolute potential of an electrode, or even the potential difference between a metal electrode and the surrounding solution, cannot be determined experimentally. It is only possible to measure the voltage across an electrochemical cell, that is, the difference of potential between two identical wires connected to two electrodes. A potential scale may be defined by measuring all electrode potentials with respect to an electrode of constant potential, called the reference electrode. The reference electrode arbitrarily chosen to establish a universal potential scale is the standard hydrogen electrode (SHE). It consists of a platinized platinum electrode (wire or sheet) immersed in an aqueous solution of unit activity of protons, saturated with hydrogen gas at a fugacity of 1 bar. The half-cell reaction is the equilibrium: H+(aq) + e- H2(g) (Eq 17) The SHE possesses the advantages of achieving its equilibrium potential quickly and reproducibly and maintaining it very stable with time (see comparison with other reference electrodes in the article “Potential Measurements with Reference Electrodes” in this Volume). From the convention, the SHE potential is taken as zero. The potential of any electrode can then be determined with respect to this zero reference and is called the potential of the electrode on the standard hydrogen scale, denoted E(SHE). The Potential Sign Convention (Reduction Convention). Before establishing tables of standard potentials for various electrodes on the standard hydrogen scale, it is necessary to fix the convention for the sign of a standard potential value (E0). The potential of any electrode is expressed with respect to the SHE by building (really or virtually) a cell in which the other electrode is a SHE. Consider a typical half-cell reaction with an oxidationreduction (O/R) couple, where O represents the oxidized species and R the reduced species. The cell is represented as: Pt, H2(g)(f = 1 bar)|H+(aq)(a = 1) || O/R (Eq 18) Depending on the position of the O/R electrode on the hydrogen scale, the spontaneous single electrode or halfcell reaction will proceed in one direction or the other: R oxidation: cR → bO + ne- combined with proton reduction: H+(aq) + e- → H2(g); or, O reduction: bO + ne- → cR combined with hydrogen oxidation: H2(g) → H+(aq) + e-. For example, the coupling of the oxidation/ reduction couple Fe2+/Fe with the H+/H2 couple brings about the spontaneous oxidation of iron (the free energy of the cell reaction is negative, the Fe2+/Fe electrode is negative). The situation is entirely different with a Cu2+/Cu system. If coupled with the H+/H2 couple, the Potential Measurements with Reference Electrodes ELECTRODE POTENTIAL MEASUREMENT is an important aspect of corrosion studies and corrosion prevention. It is included in any determination of the corrosion rate of metals and alloys in various environments and in the control of the potential in cathodic and anodic protection. The potential of an electrode can be determined only by measuring the voltage in an electrochemical cell between this electrode and an electrode of constant potential, called the reference electrode. Many errors and problems can be avoided by careful selection of the best reference electrode for a specific case and by knowledge of the electrochemical principles that control the potential measurements in order to obtain meaningful measurements. A reference electrode, once selected, must be properly used, taking into account the stability of its potential and the problem of ohmic (IR) drop. Many different reference electrodes are available, and others can be designed by the users themselves for particular situations. Each electrode has its characteristic rest potential value, which is used to convert potential values measured with respect to this reference into values expressed with respect to other references. In particular, the conversion of the potentials from or to the hydrogen scale is frequently required for use of potential- pH diagrams, which are discussed in the article “Potential versus pH (Pourbaix) Diagrams” in this Section of the volume. The Three-Electrode Device When a system is at rest and no significant current is flowing, the use of only one other electrode as a reference is sufficient to measure the test (or working) electrode potential versus the reference potential. When a current is flowing spontaneously in a galvanic cell or is imposed on an electrolytic cell, reactions at both electrodes are not at equilibrium, and there is consequently an overpotential on each of them. The potential difference measured between the two electrodes then includes the value of the two overpotentials, and it is not possible to determine the potential of the test electrode. To obtain this value, a third electrode, the auxiliary or counterelectrode, must be used (Fig. 1). In this arrangement, current flows only between the test and the auxiliary electrodes. A high-impedance voltmeter placed between the test and the reference electrodes prevents any significant current flow through the reference electrode, which then shows a negligible overpotential and remains very close to its rest potential. Most reference electrodes can be damaged by current flow. The test electrode potential and its changes under electric current flow can then be measured with respect to a fixed reference potential. The three- electrode system is widely used in laboratory and field potential measurement. Operating Conditions for Reference Electrodes When a reference electrode has been selected for a particular application, its proper use requires caution and specific measurement conditions. When measuring the potential of a polarized test electrode versus a reference electrode, it is important not to polarize or damage the latter by applying a significant current density. Also, the ohmic (IR) drop must be minimized. Very Low Current Density. It is important to use a reference electrode that operates at its known open-circuit potential and thus avoid applying any significant overpotential to it. This is achieved by using a highimpedance voltmeter that has a negligible input current and, for test electrode polarization measurements, by using an auxiliary electrode in a three-electrode system (Fig. 1). The value tolerated for the maximum overpotential on the reference electrode, at the condition that it stays under the limit over which the electrode suffers irreversible damages (like the calomel electrode), is a matter of judgment that depends on the accepted magnitude of error in the particular case under investigation. The use of an electrometer or a high-impedance voltmeter (1012 ohms) fulfills the usual requirements. When a lower impedance instrument is used, an unacceptable overpotential could result if the electrode is too polarizable. The IR Drop and Its Mitigation. The IR drop is an ohmic voltage that results from electric current flow in ionic solutions. Electrolytes have an ohmic resistance; when a current passes through them, an IR voltage can be observed between two distinct points. When the reference electrode is immersed at some distance from a working or test electrode, it is in the electric field somewhere along the current path. An electrolyte resistance exists along the path between the test and the reference electrodes. As current flows through that path, an IR voltage appears in the potential measurement according to: V = VT - VR + IR (Eq 13) where VT is the test potential to be measured, VR is the reference electrode potential, and IR is the ohmic drop. In this case, the liquid junction potential has been neglected. The IR drop constitutes a second unknown value in a single equation. It must be eliminated or minimized. The Luggin capillary is a tube, usually made pipeline accessories of glass, that has been narrowed by elongation at one end. The narrow end is placed as close as possible to the test electrode surface (Fig. 1), and the other end of the tube goes to the reference electrode compartment. The Luggin capillary is filled with cell electrolyte, which provides an electric link between the reference and the test electrode. The use of a high-impedance voltmeter prevents significant current flow into the reference electrode and into the capillary tube between the test electrode and the reference electrode compartment (Fig. 1). This absence of current eliminates the IR drop, and the measurement of VT is then possible. A residual IR drop may, however, exist between the tip of the Luggin capillary and the test electrode. This is usually negligible, however, especially in high-conductivity media. The remote electrode technique can be used only for measurement in an electrolyte with very low resistivity, usually in the laboratory. It is applicable, for example, in a molten salt solution, in which the ohmic resistance R is very small. In such a case, the reference electrode can be placed a few centimeters away from the test electrode, because the IR drop remains negligible. In other electrolytes (for example, in measurements in soils), the ohmic resistance is rather large, and the IR drop cannot be eliminated in this manner. The Current Interruption Technique. In this case, when the current is flowing, the IR drop is included in the measurement. A recording of the potential is shown in Fig. 5. At time t1, the current is interrupted so that I = 0 and IR = 0. At the moment of the interruption, however, the electrode is still polarized, as can be seen at point P in Fig. 5. The progressive capacitance discharge and depolarization of the test electrode take some time. The potential measured at the instant of interruption then represents the test electrode potential corrected for the IR drop. Precise measurements of this potential are obtained with an oscilloscope. Potential Conversion Between Reference Electrodes. Due to the number of different reference electrodes used, each potential measurement must be accompanied by a clear statement of the reference used. It is often needed to express electrode potentials versus a particular reference, regardless of the actual reference used in the measurement. The procedure is illustrated in the following example. The electrode potential of a buried steel pipe is measured with respect to a CuSO4/Cu electrode, and the value is -650 mV for a pH 4 environment. If that value is mistakenly placed in the iron E-pH (Pourbaix) diagram (Fig. 1 in the article “Potential versus pH (Pourbaix) Diagrams” in this Volume), it could be concluded that corrosion will not occur. This conclusion, however, would be incorrect, because the E-pH (Pourbaix) diagrams are always computed with respect to the SHE. It is then necessary to express the measured electrode potential with respect to the SHE before consulting the E-pH (Pourbaix) diagram. The CuSO4/Cu electrode potential is +310 mV versus SHE, so this value must be added to the measured potential: ESHE = -650 + 310 = -340 mV. Calculation and Construction of E-pH Diagrams Potential-pH diagrams are based on thermodynamic calculations. The equilibrium lines that set the limits between the various stability domains are calculated for the various electrochemical or chemical equilibria between the chemical species considered. There are three types of reactions to be considered: · Electrochemical reactions of pure charge (electron) transfer · Electrochemical reactions involving both electron and solvated proton (H+) transfer · Acid-base reactions of pure H+ transfer (no electrons involved) Pure Charge Transfer Reactions. These electrochemical reactions involve only a reduced species on one side and an oxidized species and electrons on the other side. They have no solvated protons (H+) as reacting particles; consequently, they are not influenced by pH. An example of a reaction of this type is the oxidation/ reduction of Ni/Ni2+: Ni Ni2+ + 2e-. From the Nernst equation (Eq 1), the equilibrium potential for the couple Ni2+/Ni can be written Molten Salt Corrosion Thermodynamics MOLTEN SALTS—in contrast to aqueous solutions in which an electrolyte (acid, base, salt) is dissolved in a molecular solvent—are essentially completely ionic. Thus the terms solute and solvent can be defined only in quantitative terms. For example, the terms lose their meaning in a NaCl-AgCl melt where composition can vary continuously from pure NaCl to pure AgCl. This is true even when the electrodes immersed in the melt are reversible only to some of the ions in the melt. For example, in the cell: Ag|AgCl|1NaCl|Cl2 (Eq 1) the chlorine electrode is reversible to Cl- and the silver electrode is reversible to Ag+. When this cell is used to obtain thermodynamic data, it is assumed that the cell is stable; that is, its composition does not change with time. However, when the concentration of AgCl is very low, this will not be the case, since the silver electrode will react spontaneously with the melt: Ag + NaCl = AgCl + Na (Eq 2) Thus the concentration of AgCl will spontaneously increase in the melt, and the electromotive forces (emf) measured for the preceding cell will not be stable but will change in the direction indicating increasing Ag concentration. The point at which this happens depends on the system. Thermodynamics of Cells One major use of electrochemical cells is to obtain thermodynamic data for salts. The basic thermodynamics applicable to galvanic cells for aqueous solutions is discussed elsewhere in this Volume. Only those aspects that are different for molten salts are emphasized in this article. Thus, for the cell given in Eq 1, the cell reactions Methods for Determining Aqueous Corrosion Reaction Rates. CORROSION OF MATERIALS IN AQUEOUS SOLUTIONS is often thermodynamically possible but kinetically limited. Therefore, it is important to determine the rates of corrosion processes. Corrosion rate determination can serve many engineering and scientific purposes. For example, it can be used to: · Screen available materials to find the most resistant material for a given application. · Determine operating conditions where corrosion rates are low versus those where rates are high, by varying conditions. · Determine probable service lifetimes of materials forming components, equipment, and processes. · Evaluate new alloys or treatments or existing alloys in new environments. · Evaluate lots, heats, or treatments of materials to ensure that specified quality is achieved before release, shipment, or acceptance. · Evaluate environmental conditions such as new chemical species, inhibitors, or plant- operation conditions such as temperature excursions. · Determine the most economical means of reducing corrosion through use of inhibitors, pretreatments, coatings, or cathodic protection. · Determine the relative corrosivity of one environment compared to another. · Study corrosion mechanisms. Methods for determination of corrosion rates can be differentiated between those that measure the cumulative results of corrosion over some period of time and those that provide instantaneous rate information. Corrosion rates do not often increase monotonically with environmental conditions but exhibit sharp thresholds that distinguish regions of low corrosion rates from other regions where corrosion rates are dangerously high. It is sometimes of greater interest to define these thresholds than it is to determine the rates in the regions where corrosion rates are high. Examples of the latter are pitting or crevice corrosion where passive films are broken down and local corrosion rates can be extremely high. This article addresses electrochemical methods for instantaneous rate determination and threshold determination as well as nonelectrochemical methods that can determine incremental or cumulative rates of corrosion. Fundamentals of Corrosion in Gases.. ENGINEERING MATERIALS are subject to deterioration when exposed to high-temperature environments. Whether they survive or not in technological applications depends on how fast they react. The rate of corrosion varies widely; some intermetallics (β-NiAl) react extremely slowly. Some metals (Fe) oxidize very rapidly, whereas other metals (Cr, Co) react relatively slowly. From the chemical point of view, the gas- metal reactions represent a broad class of heterogeneous reactions. The composition and structure of the scales produced on metals are a key factor in their behavior in technical applications. Historically, corrosion in gases has been primarily a problem in combustion systems. Thus, the gas-metal reactions are usually referred to as oxidation in its broad chemical sense, whether the reaction is with pure oxygen, water, sulfur dioxide (SO2), or whatever the gas might be. The corrosion product (oxide layer) is termed scale. Being the corrosion product, the protective properties of scale decrease the reaction rate. The concepts and methods developed to understand gas-metal reactions can be used to describe any arbitrary gas solid reaction at high temperature; for example, oxidation of the silicone carbide. The high temperature corrosion is a highly technical challenge; the reason for this is that the efficiency of thermal processes and engines increases with operating temperature. Such high- temperature service is especially damaging to most metals because of the exponential increase of reaction rate with temperature. In most cases, corrosion resistance at high temperatures does not accompany the good mechanical properties of structural materials; therefore, protective coatings must be applied. Electrochemical principles are insufficient to understand the mechanism of oxidation. For gaseous reactions, a basic understanding of the diffusion processes is much more profitable. The first results of a high-temperature corrosion study (not yet defined as corrosion and even diffusion) were published in 1684 by Boyle in Experiments and Considerations about the Porosity of Bodies in Two Essays. In studying reactive diffusion in the Cu-S system, Boyle reported the observation of interaction between copper and sulfur through examination of metallographic cross sections. Electrochemistry and aqueous corrosion principles were developed at the beginning of the 19th century. In 1855 Fick formulated the basic principles of diffusion in solids. The systematic study of high-temperature oxidation began in the 1920s. In 1933 Wagner published his pioneering paper on gas corrosion of metals. The first journal devoted to corrosion in gases, Oxidation of Metals, was published in the 1960s. The following articles introduce the subject of gas corrosion to professional engineers and students. A brief summary of thermodynamic concepts is followed by an explanation of the defect structure of solid oxides and the effect of these defects on the rate of mass transport. Commonly observed kinetics of oxidation are described and related to the observed corrosion mechanisms, as illustrated in Fig. 1 of the next article, “Thermodynamics of Gaseous Corrosion.” In high temperature gaseous corrosion, the oxidant first adsorbs on the metal surface in molecular (physical adsorption) and ionic form (chemical adsorption), and it may also dissolve in metal. Oxide nucleates at favorable sites and most commonly grows laterally, due to surface diffusion, to form a complete thin film (scale). As the scale thickens, it provides a protective barrier to shield the metal from the gas. For scale growth, electrons must move through the film to reach the oxidant atoms adsorbed on the surface, and oxidant ions and/or metal ions must move through the scale barrier. Diffusion of the oxidant into the metal may result in internal oxidation. Growth and thermal stresses in the oxide scale may create microcracks and/or delaminate scale from the underlying metal. Stresses affect the diffusion process and modify the oxidation mechanism and very often cause scale spallation. Improved oxidation resistance can be achieved by developing better alloys, by applying protective coatings, and by altering the composition of the gas phase. Thermodynamics of Gaseous Corrosion METALS can react chemically when exposed to air or to other more aggressive gases. The reaction rate of some metals is so slow that they are virtually unattacked, but for others, the reaction can be violent. As with most chemical processes, elevated-temperature service is more severe because of the exponential increase in reaction rate with temperature. The most common reactant is oxygen in the air; therefore, all gas-metal reactions are usually referred to as oxidation, using the term in its broad chemical sense whether the reaction is with oxygen, water vapor, hydrogen sulfide (H2S), or whatever the gas might be. Throughout this article, the process is called oxidation, and the corrosion product is termed an oxide. Corrosion in gases differs from aqueous corrosion in that electrochemical principles do not help greatly in understanding the mechanism of oxidation. For gaseous reactions, a fundamental knowledge of the diffusion processes involved is much more useful. The principles of high-temperature oxidation began to be understood in the 1920s, whereas electrochemistry and aqueous corrosion principles were developed approximately 100 years earlier. The first journal devoted to corrosion in gases, Oxidation of Metals, began publication in 1970. This article addresses thermodynamic concepts; the commonly observed kinetics of oxidation are described in the article “Kinetics of Gaseous Corrosion Processes” in this Volume. The mechanisms of oxidation are shown schematically in Fig. 1. The gas is first adsorbed on the metal surface as atomic oxygen. Oxide nucleates at favorable sites and most commonly grows laterally to form a complete thin film. As the layer thickens, it provides a protective scale barrier to shield the metal from the gas. For scale growth, electrons must move through the oxide to reach the oxygen atoms adsorbed on the surface, and oxygen ions, metal ions, or both must move through the oxide barrier. Oxygen may also diffuse into the metal. Growth stresses in the scale may create cavities and microcracks in the scale, modifying the oxidation mechanism or even causing the oxide to fail to protect the metal from the gas. Improved oxidation resistance can be achieved by selection of suitable alloys for the given environment and by application of protective coatings.