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## Über dieses Buch

Corrosion due to water is one of the most significant and complex causes of damage to metallic products. Written from the viewpoint of physical chemistry, this authoritative and established text deals with the aqueous corrosion of metals. Available for the first time in English, Corrosion of Metal addressing engineers, metallurgists, physicists and chemists. This self-contained, valuable reference comprehensively organizes and makes readily accessible the accumulated wealth of fundamental and applied knowledge. The concentration is on the underlying essentials of corrosion and failure, and the material is consistently presented in relation to practical applications to corrosion protection. The first chapters introducing the physicochemical principles are ideal for students. The following chapters provide an overview of the state of research for those familiar with the fundamentals. An exhaustive bibliography and appendices conclude the volume.

## Inhaltsverzeichnis

### 1. Introduction

Abstract
Corrosion of metals is understood to be degradation of metals by chemical surface reactions with aggressive components of the environment. The metals may be structural materials such as steel reinforcments in concrete structures, or steel cables of suspension bridges; or they may be functional materials such as dental alloys, or copper leads for printed circuits. A typical example of corrosion is rusting of iron: Rust, a mixture of oxides and hydroxides of iron, is the product of iron surface atoms reacting with oxygen and water, both present in the surrounding moist atmosphere, or else present in a surrounding aqueous solution. Corrosion thus differs from wear, which is degradation caused by mechanical friction, as for instance abrasion. The primary products of wear are metallic particles, whereas products of corrosion always are non-metallic chemical species either solid or dissolved.
Helmut Kaesche

### 2. Corrosion Reactions and Corrosion Products

Abstract
To outline corrosion reaction chemistry we turn in the first place to acid corrosion, considering iron in dilute hydrochloric acid solution as a typical example. The reaction equation is
$$(\text{Fe})_m + 2(\text{HCl})_l \to (\text{FeCl})_l + (\text{H}_\text{2} )_g ,$$
(2.1)
where the subscripts m, l, and g indicate the phases solid metal, liquid aqueous solution, and gas, containing the species solid pure iron, dissolved hydrochloric acid, dissolved ferrous chloride, and gaseous hydrogen, respectively. When components of different phases participate in a reaction, this indexing is useful, but we shall drop it nevertheless whenever possible without danger of confusion, in order to simplify notation. The arrow ‘→’ indicates that the reaction spontaneously proceeds as written from left to right: metallic iron and dissolved hydrochloric acid are the reactants, dissolved ferrous chloride and gaseous hydrogen the products of the overall reaction.
Helmut Kaesche

### 3. Chemical Thermodynamics of Corrosion

Abstract
In any corrosion process, the solid metal, the contacting aqueous solution and, normally, an adjacent gas phase are parts of a system in which the reaction causes physical and chemical changes. A description of the system in terms of thermodynamics then requires the specification of the boundaries separating the system and its surroundings [1-6]: These may be permeable a) to both matter and energy, where energy is work and/or heat, or b) permeable to energy alone, or c) impermeable to both matter and energy. The system is then said to be a) open, or b) closed, or c)isolated, respectively.
Helmut Kaesche

### 4. The Electrolytic Mechanism of Corrosion

Abstract
In this chapter, we shall verify the supposition that corrosion proceeds by the superposition of kinetically independent electrode half-reactions: anodic metal dissolution and cathodic reduction of an oxidizing species, such as hydrogen ions, or oxygen molecules, or others.
Helmut Kaesche

### 5. The Kinetics of Electrode Reactions

Abstract
In this section, the discussion of the mechanism of hydrogen evolution at metallic surfaces will be used to outline the principles of electrode reaction kinetics. This follows the historic development, since originally hydrogen evolution was considered to be the model example of a simple electrode reaction. The view is not really up-to-date: The theory of electrode kinetics is much more advanced concerning redox reactions such as: Fe2+= Fe3++ e, involving only a tunneling exchange of electrons with an ion undergoing a change in valence and a rearrangement of the hydration shell only. It is for this type of reactions that an advanced theory is available which, however will be disregarded in the following. Hydrogen evolution certainly is an important half-reaction of corrosion processes, and the principles of kinetics can be demonstrated for this example quite well, if only at the expense of a quantitative quantum-mechanical treatment. For a deeper understanding of electrode kinetics the reader is referred to the literature [1-5].
Helmut Kaesche

### 6. Uniform Electrolytic Corrosion

Abstract
Continuing to verify the principles of electrolytic corrosion, we turn from liquid amalgams to solid metals and to solutions with pH up to ca. 7, which is the overall range of ‘acid’ corrosion by superposition of anodic metal dissolution and cathodic hydrogen deposition by reduction of hydrogen ions rather than water. To exclude superposed oxygen reduction, the solution is assumed to be deaerated. The temperature is at 25°C or at some similar ambient value. The solution volume is supposed to be large, so that initial concentrations remain virtually constant during the course of the corrosion reaction, which is: Me + zH+ → Mez+ (z/2)H2. The metal surface is thought to be bare, i.e. free of films such as air-formed oxide films, and free of layers of solid corrosion products such as oxides or hydroxides.
Helmut Kaesche

### 7. Adsorption Inhibitors of Acid Iron Corrosion

Abstract
By definition, corrosion inhibitors are soluble substances which, added to the aggressive solution, reduce the corrosion rate without participating in the overall reaction. As they are not reactants in the overall corrosion reaction they are negative catalysts. In this chapter we shall not deal with inhibitors such as sodium carbonate, or other substances increasing the pH of the solution and favoring spontaneous passivation, or hydrazine as a reagent reducing dissolved oxygen and thus removing the aggressive species in many neutral-solution corrosion processes. Also, we shall disregard inhibitors as components of protective paint systems, such as the classical read lead oxide, Pb3O4, an intensely oxidizing species again aiming at passivating iron and steel surfaces. Instead, the object will presently be the mechanism of corrosion protection by inhibitors of acid corrosion acting through adsorption on the metal surface from a solution containing the inhibitive substance in millimolar or even smaller concentration only. The typical industrial application of adsorption inhibitors is protection of hot-rolled steel against heavy metal loss by corrosion during acid pickling. Another important application is protection of underground gas and oil well piping against corrosion and accompanying embrittlement by attack by underground acid hydrogen sulfide solutions.
Helmut Kaesche

### 8. Corrosion of Homogeneous Alloys

Abstract
Concerning ambient-temperature corrosion of solid alloys, a preliminary note in Chap. 3 on thermodynamics was that bulk/surface equilibria of alloy components cannot normally be expected, because solid-state diffusion usually is too slow to counteract selective dissolution rates of components. The situation is different for liquid amalgams, as they can be mechanically stirred. Also, liquid amalgam surfaces are uniform down to atomic distances. For solid alloys, slow bulk-to-surface diffusion usually will instead leave the surface depleted with respect to ‘fast’ components, and enriched with respect to ‘slow’ components, and the result is an overall decrease in corrosion rate. The crystallografic structure of solid surfaces suggests that the main effect is blocking of kink sites active in metal dissolution, and this supposition will further below be seen to account for what may be described as fine-tuning of concepts, which we do not, however, immediately need to approach.
Helmut Kaesche

### 9. Rusting of Iron and Steel

Abstract
Iron and unalloyed or low-alloy steels are the predominant metallic structural materials governing technology and by far out-weighing all other structural metals. Loss of iron and steel by corrosion, and costs of protecting iron and steel against corrosion by coatings, lacquers, and varnishes therefore have outstanding economic importance. Since corrosion of iron and steel usually is rusting, a closer look at the mechanism of iron reacting with oxygen dissolved in approximately neutral solutions appears appropriate.
Helmut Kaesche

### 10. Passivity

Abstract
When iron is immersed in concentrated nitric acid, it is seen to immediately react with the solution, as evident by a burst of evolution of gaseous nitrous oxides, but the reaction stops almost immediately, and henceforth the metal remains bright and uncorroded indefinitely. It is ‘passivated’ with respect to the thermodynamically strongly favored corrosion reaction, thus showing ‘similar-to-noble metal behaviour’. The explanation — already proposed by Faraday — is that the oxidizing power of HNO3 is sufficient to produce an thin, nonporous oxide film which, while easily soluble, happens to dissolve into the acid at only a negligible rate. The film thickness is below the wavelength of visible light and the film, therefore, transparent.
Helmut Kaesche

### 11. Galvanic Corrosion Cells

Abstract
Two dissimilar metals, A and C, subject to corrosion in the same electrolyte, normally are at different corrosion potentials, (εa)corr and (εc)corr. The electrode system metal C/solution/metal A then is a galvanic cell with the open-circuit voltage
$$\Delta \varepsilon = (\varepsilon _c )_{corr} - (\varepsilon _a )_{corr}$$
(11.1)
We choose indices such that (εc)corr is more positive than (εa)corr so that A is the anode and C the cathode in the cell. When short circuited by an external metallic connection between A and C, a cell current, j, will flow, moving positive charge carriers in the solution, and electrons in the metalllic connection, in the direction from the anode to the cathode. The current passing the anode/solution interface has the positive strength +|j|; a current of the same strength but opposite sign, −|j|, passes the solution/cathode interface. At both electrodes the total current, j, is, in principle, the sum of anodic and cathodic partial currents. Often, circumstances are simpler, when at the anode the cathodic partial current of the reduction of an oxidizing species, and/or when at the cathode the partial current of anodic metal dissolution are non-existent or negligible, but this is not the general case. The early assumption that, instead, the action of galvanic cells is the general principle of corrosion in fact has been a source of much confusion.
Helmut Kaesche

### 12. Pitting Corrosion

Abstract
In this chapter, the subject is pitting corrosion as a dangerous process causing deep localized attack in metal surfaces, including perforations of thin-walled structures. In this context, we do not consider trivial cases of localized corrosion caused by accidental defects in coatings. Instead, the matter of serious practical concern is unexpected pitting of metals passivated by oxide films. Attending to this type of pitting only, two other types of attack also resulting in pitting are disregarded: One is pitting of iron after prolonged exposure to acid solutions containing an adsorption inhibitor [2a]. The other is pitting of Gs [3a], which is interesting, as apparently a space charge layer in the semiconductor surface is now in the role of the oxide film on passive metals. The reason to concentrate on pitting of passive metals is its outstanding practical importance for safety considerations. In contrast to uniform corrosion, whose rate will normally be predictable from vast experience, pitting corrosion tends to proceed at an unexpected and high rate of localize metal dissolution. This major aspect of the matter is, at least in principle, easily understood considering that passive metals usually are at a corrosion potential, £corr which by the presence of the oxide film is shifted far beyond the equilibrium potential of most of the metals alloys components, EMe/Me z+. Whenever in pore of the film bare metal is exposed to the electrolytic solution, the overpotential, nMe = E corr − EMe/Me Z+, available for the acceleration of anodic metal dissolution thus tends to be very high.
Helmut Kaesche

### 13. Intercrystalline and Intracrystalline Corrosion

Abstract
Intercrystalline corrosion (IC) is preferential attack along grain boundaries into the interior of polycrystalline metaUic materials. Intracrystalline corrosion is preferential attack into grains, along subgrain boundaries, or stacking fault planes, or other crystallografic irregularities.
Helmut Kaesche

### 14. Hydrogen Embrittlement

Abstract
An important danger to structural materials subject to tensile load is unexpected brittle, instead of usually predictable ductile, fracture, and a notorious variety is embrittlement by uptake of atomic hydrogen as a solute in the metal lattice. Since ductility of metals is caused by dislocations moving in stress fields, hydrogen embrittlement usually is an effect of interaction of absorbed hydrogen with dislocations. The subject and its implications are vast, and we shall not go into many details, referring the reader to competent reviews [2]. Hydrogen embrittlement is not, however, always caused by hydrogen/dislocation interaction. Atomic hydrogen, if dissolved to high concentration or high thermodynamic activity, also can recombine at internal surfaces to molecular gaseous hydrogen at high pressure, and the combination of the two effects tends to very much complicate the situation [3]. Therefore, the two limiting cases of pure effects: a) high-pressure internal molecular hydrogen and b) dissolved atomic hydrogen, will be briefly introduced presently, postponing the description of details to later sections. For the limiting cases, the main principles are, in fact, well understood and well documented [3-5].
Helmut Kaesche

### 15. Stress Corrosion Cracking

Abstract
Stress corrosion cracking (SCC) is defined as crack nucleation and propagation in metals caused by the synergystic action of tensile stresses, either constant or slowly changing with time, together with crack-tip chemical reactions or other environment-induced crack-tip effects. In principle, dry hydrogen-induced cracking also is SCC, but we shall, in this chapter, assume that the environmental effect is electrolytic crack tip corrosion.
Helmut Kaesche

### 16. Corrosion Fatigue

Abstract
Fatigue, in general, is the loss in strength of structural materials caused by cyclic loads [1]. A possible dangerous outcome is fatigue crackings i.e. the initiation and propagation of quasi-brittle cracks under the influence of loads cycled with an amplitude, ∆σ = σmax − σmin, where σmax is lower than the static ultimate strength, Rm, of smooth samples, or else with an amplitude, ∆K = K maxK min, of the stress intensity factor, where K max is lower than the fracture toughness, K c, of notched samples. Fatigue cracking is a widespread source of failure of structural materials. In contrast to stress corrosion cracking, it is a variant of ‘subcritical’ crack growth which does not primarily require simultaneous racktip corrosion processes. As long as it is not enhanced by an environment, it is not corrosion fatigue and, therefore not a problem of corrosion science but of metal physics. Also, corresponding to the prime importance of the subject, the expert literature collected in monographs, in meeting volumes, and in journals is huge and far beyond the scope of the present text.
Helmut Kaesche

### 17. Appendix

Abstract
Anodic protection is the technique of minimizing corrosion by external anodic polarization of metallic structures into the potential range were the metal or the alloy in question is passive. The method, therefore, implies the change from active e.g. acid corrosion to passivity, and it is useful only if the passive corrosion rate, as indicated by the passive current density, ipass defined in Chap. 10, is negligible. Cases of obvious practical relevance are iron and steel, especially high-alloy chromium-nickel steels in acid solutions. In the range of passivity, and in the absence of oxidizing agents in the solution other than H+, the total electrolytic current density is equal to ipass The anodic protective external current, jprot then is equal to the productA ×i pass of the surface. A, of the structure and the passive current density, and often will be very low, rendering the technique extremely useful. Recalling anodic oxygen evolution at passive iron and transpassivity of chromium steels and stainless steels, it is obvious that a potentiostatic polarizing circuit normally will be required in order to avoid driving the potential to undesirable anodic electrode reactions other than passive metal dissolution. Also, when passivity imphes danger of pitting, or intercrystalline corrosion, or stress corrosion cracking, unskilled application of anodic protection may be a serious error. When for the same, potentially fatal corrosion processes threshold potentials exist, as especially for pitting, potentiostatic sub-threshold protection may still be a safe possibility. At any rate it is easily seen, that anodic protection does require knowledge and careful consideration of electrode kinetics. When apphcable, it is, of course, very useful and economic.
Helmut Kaesche

### Backmatter

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