section epub:type=”chapter”> In has been seen that most materials that are unstable in oxygen tend to oxidize. Principally, the main concern is with the loss of material at high temperatures, in dry environments, and it is found that, under these conditions, oxidation is usually controlled by the diffusion of ions or the conduction of electrons through oxide films that are formed on the material surface. Because of the thermally activated nature of the diffusion and reaction processes, it is seen that the rate of oxidation is much greater at high temperature than at low temperature—although even at room temperature, very thin films of oxide do form on all unstable metals. This minute amount of oxidation is important because it helps keep sliding surfaces apart and so influences the coefficient of friction. But the loss of material by oxidation at room temperature under these dry conditions is very slight. This chapter focuses on a different environment for corrosion. Under wet conditions, the situation is different. When mild steel is exposed to oxygen and water at room temperature, it rusts and the loss of metal quickly becomes appreciable. Unless precautions are taken, the life of many things is limited by wet corrosion. In the last two chapters we showed that most materials that are unstable in oxygen tend to oxidize. We were principally concerned with loss of material at high temperatures, in dry environments, and found that, under these conditions, oxidation was usually controlled by the diffusion of ions or the conduction of electrons through oxide films that formed on the material surface. Because of the thermally activated nature of the diffusion and reaction processes we saw that the rate of oxidation was much greater at high temperature than at low, although even at room temperature, very thin films of oxide do form on all unstable metals. This minute amount of oxidation is important: it protects, preventing further attack; it causes tarnishing; it makes joining difficult; and (as we shall see in Chapters 29 and 30) it helps keep sliding surfaces apart, and so influences the coefficient of friction. But the loss of material by oxidation at room temperature under these dry conditions is very slight. Under wet conditions, the situation is different. When mild steel is exposed to oxygen and water at room temperature, it rusts and the loss of metal quickly becomes appreciable. Unless precautions are taken, the life of many things, from cars to bridges, from ships to aircraft, is limited by wet corrosion. The annual bill worldwide for either replacing corroded components, or preventing corrosion (e.g., by painting Scotland’s Forth Bridge), is around US$2.2 trillion—more than 3% of Britain’s GDP. Figure 27.1 shows a graphic example. Why the dramatic effect of water on the rate of loss of material? As an example we shall look at iron, immersed in aerated water (Figure 27.2). Iron atoms pass into solution in the water as Fe++, leaving behind two electrons each (the anodic reaction). These are conducted through the metal to a place where the “oxygen reduction” reaction can take place to consume the electrons (the cathodic reaction). This reaction generates OH− ions which then combine with the Fe++ ions to form a hydrated iron oxide Fe(OH)2 (really FeO·H2O); but instead of forming on the surface where it might give some protection, it often forms as a precipitate in the water itself. The reaction can be summarized by the following, just as in the case of dry oxidation: The formation and solution of Fe++ is analogous to the formation and diffusion of M++ in an oxide film under dry oxidation; and the formation of OH− is closely similar to the reduction of oxygen on the surface of an oxide film. However, the much faster attack found in wet corrosion is due to the following: The result is that the oxidation of iron in aerated water (rusting) goes on at a rate that is millions of times faster than that in dry air. In dry oxidation we quantified the tendency for a material to oxidize in terms of the energy needed, in kJ mol−1 of O2, to manufacture the oxide from the material and oxygen. Because wet oxidation involves electron flow in conductors, which is easier to measure, the tendency of a metal to oxidize in solution is described by using a voltage scale rather than an energy one. Suppose we could separate the cathodic and the anodic regions of a piece of iron, as shown in Figure 27.3. Then at the cathode, oxygen is reduced to OH−, absorbing electrons, and the metal therefore becomes positively charged. The reaction continues until the potential rises to +0.8 V. Then the coulombic attraction between the positive charged metal and the negative charged OH− ion becomes so large that the OH− is pulled back to the surface, and reconverted to H2O and O2; in other words, the reaction stops. At the anode, Fe++ forms, leaving electrons behind in the metal which acquires a negative charge. When its potential falls to –0.6 V, that reaction, too, stops (for the same reason as before). If the anode and cathode are now connected, electrons flow from the one to the other, the potentials converge, and both reactions start up again. The difference in voltage of 1.4 V is the driving potential for the oxidation reaction. The bigger it is, the bigger the tendency to corrode. Figure 27.4 shows a typical Pourbaix diagram (or electrochemical equilibrium diagram)—in this case for copper. The diagrams are maps that show the conditions under which a metal: The axes of the diagram are the electrochemical potential of the metal (in volts) and the pH, which indicates how acidic or alkaline the water is. The three conditions are then represented on the diagram by fields of immunity, corrosion, and passivation. The diagram also shows the line for the oxygen reduction reaction. Note that this slopes, from a potential of +1.23 V when the water is very acid (pH = 0) to +0.4 V when the water is very alkaline (pH = 14). This is because the oxygen reduction reaction produces OH− ions, so increasing the concentration of these ions in the solution (which is what we do when we go from acid to alkaline conditions) shifts the equilibrium position of the reaction. In addition, Figure 27.4 shows that when copper is immersed in water with a pH less than about 6 to 7, it may corrode if the potential of the copper is greater than about 0.2 to 0.3 V. So if there is air in the water, the copper may corrode, because the oxygen reduction line is positioned at a much higher potential. For example, if the pH is 5, the voltage driving force for corrosion is 0.8 V, and because there is no stable oxide film, corrosion may occur. If the pH of the water is between 7 and 12 (from neutral to strongly alkaline), the copper will grow a stable film of copper oxide, which may protect the metal against further corrosion—even though there will still be a driving force of about 0.8 V between the potentials of the copper and the oxygen reduction reaction. It should be emphasized that these diagrams cannot give any indication of the actual rates of corrosion in any situation—they just show where corrosion cannot occur, and where a surface film should form which might prevent or slow down corrosion. Experimental data and field experience in the real environment are essential for corrosion design—even more so because corrosion processes in general, and the barrier properties of surface films in particular, can be very strongly affected by the presence of other ions. These speed up corrosion processes and can attack surface films, which is why seawater—high in chloride ion, Cl−—is such bad news for many metals (see the heavily corroded seawater ballast tank in Figure 27.1). Provided the water is neutral or alkaline (and low in other ions) the rate of corrosion is usually extremely low. This would be expected from the Pourbaix diagram if the oxide film acted as an effective barrier. This is why copper is much used for water pipes. However, if the water is mildly acidic, corrosion “pinholes” can form in the pipe wall (the Pourbaix diagram shows that there is no oxide film below a pH of about 6). Copper is much used for architectural purposes, particularly in the United States and Australia. Figure 27.5 shows a copper rain water pipe on the outside of a building. The thin sheet copper has not corroded significantly, but a protective film has formed on exposure to the elements, which is why the color is darker than freshly polished copper. Older architectural copper usually develops a green color, because the carbon dioxide in the air forms a copper carbonate film. This is just as well, because carbon dioxide also makes the rain water acidic! Copper alloys like brass and bronze are also resistant to corrosion—the fire hydrant shown in Figure 27.6 is as shiny as the day it was made.
Wet Corrosion of Materials
Publisher Summary
27.1 Introduction
27.2 Wet Corrosion
27.3 Voltage Differences as the Driving Force for Wet Oxidation
27.4 Pourbaix (Electrochemical Equilibrium) Diagrams
27.5 Some Examples
Copper
