The kinetics of corrosion reactions

It is important to realize that metal corrosion is a non-reversible and nonequilibrium situation. The metal takes up a mixed-potential which depends on the thermodynamics and kinetics of both metal dissolution and the balancing cathodic process. Considering a simple corrosion reaction  

The reaction continuously leads to the production of and H2 and also consumes i.e. the process is non-equilibrium. 

No charge build-up occurs the rate at which electrons are produced in the metal dissolution reaction is balanced by their rate of consumption by the hydrogen evolution reaction. 

In practice, the metal takes up an intermediate value during corrosion, the corrosion potential Ecorr (Fig- corr is the corrosion current. corr icoRR appear at the point of intersection of linear polarization curves for metal dissolution and hydrogen evolution on a semilogarithmic plot (Fig. 10.8). In order to make the diagram applicable to localized attack (when the anode and cathode sizes are disparate) it is necessary to strictly plot log current rather than log current density. 

The i-E curves used in this chapter are related to the Evans diagram, beloved by corrosion chemists, but follow the conventions used throughout the rest of this book. While they are greatly simplified polarization diagrams, their usefulness lies in their ability to predict and rationalize the kinetics of corrosion processes. The figures show, in a diagnostic fashion, the following features  

The partial current/corr is called the corrosion current and the rate of corrosion estimated from this current and Faraday s law should be the same as that from weight loss experiments where the rate of metal dissolution is determined by weighing a sample of metal before and after standing in the electrolyte for a period. It should be noted that the corrosion potential is a, so-called, mixed potential since it is determined by the kinetics of two electron transfer couples, M/M and H2O/H2 (cf. the equilibrium potential for a single couple. Chapter 1). 

To understand more quantitatively the dependence of the corrosion potential and current on the kinetic parameters of the two electrode processes (9.3) and (9.5), the I-E data of Fig. 9.2 should be replotted in a semi-logarithmic form, i.e. as a Tafel plot (Fig. 9.3). A linear Tafel line will only be observed in the potential range where a single electrode process predominates. However, these Tafel lines can be extrapolated to the corrosion potential and if necessary (as in the figure) also to the calculated equilibrium potentials for the H2O/H2 and M/M couples where the exchange currents are I and respectively. The equations for the two 

It can be seen that the corrosion current and potential depend on both the equilibrium potentials for the hydrogen evolution reaction and metal dissolution calculated from the Nernst equation, and the kinetic parameters, the exchange currents and the Tafel slopes. Table 9.1 shows the corrosion currents calculated for some typical values of these parameters it is also important to note that even a 

In most practical situations, the choice of metal is such that the reversible cell potential is seldom much in excess of 0.5 V. Even so, because of the large variations in exchange currents and mechanisms for the two electrode processes which are possible, the corrosion rate for different metals and environments can still vary by many orders of magnitude. 

So far, the quantitative discussion of corrosion has only dealt with deoxygenated solutions. In the real world, this is a very unusual environment oxygen is almost always present. How does the presence of oxygen change our discussion of corrosion In fact, two situations must be considered. 

---

As the nature of the electrified interface dominates the kinetics of corrosive reactions, it is most desirable to measure, e.g., the drop in electrical potential across the interface, even where the interface is buried beneath a polymer layer and is therefore not accessible for conventional electrochemical techniques. The scanning Kelvin probe (SKP), which measures in principle the Volta potential difference (or contact potential difference) between the sample and a sensing probe (which may consist of a sharp wire composed of a conducting, stable phase such as graphite or gold) by the vibrating condenser method, is the only technique which allows the measurement of such data and therefore aU modern models which deal with electrochemical de-adhesion reactions are based on such techniques [1-8]. Recently, it has been apphed mainly for the measurement of electrode potentials at polymer/metal interfaces, especially polymer-coated metals such as iron, zinc, and aluminum alloys [9-15]. The principal features of a scanning Kelvin probe for corrosion studies are shown in Fig. 31.1. 

Thermogravimetry is the most commonly used method for studying the kinetics of corrosion reactions at high temperatures. The sample is placed in an oven with a controlled atmosphere, where one monitors the variation of its mass by means of a microbalance (Figure 9.3). The precision of the technique is on the order of micrograms. If all corrosion products remain on the surface of the sample, the measured mass increase as a function of time permits to identify the rate law that governs the reaction kinetics. 

Temperature and time also affects the corrosion reaction. It is generally known that the rate of a chemical reaction increases with increasing temperature. As a rule of thumb the reaction rate doubles for every 10° increase in temperature. The kinetics of corrosion reactions can develop differently. Fig. 1-8. A constant, linear rate of reaction is observed for acidic corrosion where no protective layer is formed (curve a). Decreasing rates of corrosion are often ob-... 

The kinetics of corrosion reactions 499 Several obvious points may be made ... 

The order of nobiKty observed in actual practice may differ from that predicted thermodynamically. The reasons are that some metals become covered with a passivatingfEm of reaction products which protects the metal from further attack. The dissolution reaction may be strongly irreversible so that a potential barrier must be overcome. In this case, corrosion may be inhibited even though it remains energetically favorable. The kinetics of corrosion reactions are not determined by the thermodynamics alone. 

The rate (or kinetics) and form of a corrosion reaction will be affected by a variety of factors associated with the metal and the metal surface (which can range from a planar outer surface to the surface within pits or fine cracks), and the environment. Thus heterogeneities in a metal (see Section 1.3) may have a marked effect on the kinetics of a reaction without affecting the thermodynamics of the system there is no reason to believe that a perfect single crystal of pure zinc completely free from lattic defects (a hypothetical concept) would not corrode when immersed in hydrochloric acid, but it would probably corrode at a significantly slower rate than polycrystalline pure zinc, although there is no thermodynamic difference between these two forms of zinc. Furthermore, although heavy metal impurities in zinc will affect the rate of reaction they cannot alter the final position of equilibrium. 

Electroless metal deposition at trace levels in the solution is an important factor affecting silicon wafer cleaning. The deposition rate of most metals at trace levels depends mainly on the metal concentration and some may also depend on the interaction with other species as well. For copper the deposition rate at trace levels in HF solutions is different for n and p types. It depends on illumination for p-Si but not for n-Si. It is also different in HF and BHF solutions. In a HF solution the deposition process is controlled by both the supply of minority carriers and the kinetics of cathodic reactions. Thus, a high deposition rate occurs on p-Si only when both and illumination are present. In the BHF solution, the corrosion process is limited by the supply of electrons for p-Si whereas for n-Si it is limited by the dissolution of silicon because the reaction rate is indepaidmt of concentration and illumination. The amount of copper deposition does not correlate with the corrosion current density, which may be attributed to the chemical reactions associated with hydrogen reduction. More information on trace metal deposition can be found in Chapters 2 and 7. 

In this chapter we will examine oxidation-reduction stoichiometry, equilibria, and the graphical representation of simple and complex equilibria, and the rate of oxidation-reduction reactions. The applications of redox reactions to natural waters will be presented in the context of a discussion of iron chemistry the subject of corrosion will provide a vehicle for a discussion of the application of electrochemical processes a presentation of chlorine chemistry will include a discussion of the kinetics of redox reactions and the reactions of chlorine with organic matter finally, the application of redox reactions to various measurement methods will be discussed using electrochemical instruments as examples. 

Cl are chemical compounds or their blends able to retard metal corrosion even in small additions in aggressive media. Cl change the kinetics of electrochemical reactions that bring about corrosion so that the corrosion process rate retards significantly. 

Inhibited polymer films are distinguished from barrier-type films by their electrochemical activity. This property arises from the ability of these films to liberate Cl from the polymer matrix and transfer it to the mated metal surface. The Cl strongly affects the kinetics of electrochemical reactions leading to corrosion damage of metals. The inhibiting mechanism of corrosion can be treated in this case as follows. 

The information about the tendenqf for corrosion to occur that can be obtained from thermodynamic calculations is important and useful. However, most of the science and engineering aspects in the field of corrosion focus on knowing and reducing the rate of corrosion. The rate of corrosion is not addressed by thermodynamics it falls instead within the purview of kinetics. So the kinetics of electrochemical reactions in general, and corrosion reactions specifically, are at the heart of the subject of corrosion. This chapter will introduce electrochemical kinetics at a simple level, with sufficient detail to develop the concept of mixed potential theory. The interested reader is referred to other volumes of this encyclopedia and to textbooks in corrosion [1-9] for a more detailed description. The kinetic underpinnings of some of the electrochemical techniques for determination of corrosion rate will also be presented. The influence of transport on the rates of electrochemical reactions will be discussed in the next chapter (see Chapter 1.4). 

The kinetics of electrochemical reactions described in the last chapter are only strictly applicable to the situation in which the surface concentrations of all species are identical to the bulk solution concentrations. Mass transport considerations often play an important role in both cathodic and anodic corrosion reactions. More details in the area of mass transport can be found in several textbooks[l-3]. 

The complete determination of position and shape of overvoltage curves requires knowledge about the quantities io, it, be, and ba, of which the first one is given by Nemst s equation or the Pourbaix diagram. Through such knowledge, the best tool we have for description of the kinetics of electrode reactions is available. With a reservation for passivation effects (which we shall return to) we can with this background obtain a quantitative description of any corrosion process, in principle as shown in Section 4.9. 

In reality, this equation is only valid in the absence of junction potentials (Section 2.4). In addition, it is often difficult, even impossible, to experimentally establish equilibrium conditions because corrosion reactions can occur, or because the kinetics of electrode reactions may be slow. 

The Schikorr reaction results in the formation of a compact layer of magnetite whose growth governs the kinetics of corrosion. Below 300 °C, the rate of film growth is very slow and the magnetite film effectively protects against further corrosion (Section 8.1). 

The kinetics of this reaction are in the focus of research groups worldwide. The studies of Kinoshita in PAFC-based Systems showed that the carbon corrosion of the most commonly used Vulcan XC-72 catalyst support material starts at potentials higher than 0.8 V and temperatures higher than 65° C. When the carbon material is at room temperature, the oxidation starts at potentials higher than 1.0 V. So the corrosion of... 

Electrochemical reaction kinetics is essential in determining the rate of corrosion of a metal M exposed to a corrosive medium (electrolyte). On the other hand, thermodynamics predicts the possibility of corrosion, but it does not provide information on how slow or fast corrosion occurs. The kinetics of a reaction on a electrode surface depends on the electrode potential. Thus, a reaction rate strongly depends on the rate of electron flow to or from a metal-electrolyte interface. If the electrochemical system (electrode and electrolyte) is at equilibrium, then the net rate of reaction is zero. In comparison, reaction rates are governed by chemical kinetics, while corrosion rates are primarily governed by electrochemical kinetics. 

---