Metal oxidation reaction

The passive state of a metal can, under certain circumstances, be prone to localized instabilities. Most investigated is the case of localized dissolution events on oxide-passivated surfaces [51, 106, 107, 108, 109, 110, ill, 112, 113, 114, 115, 116, 117 and 118]. The essence of localized corrosion is that distinct anodic sites on the surface can be identified where the metal oxidation reaction (e.g. Fe —> Fe + 2e ) dominates, surrounded by a cathodic zone where the reduction reaction takes place (e.g. 2Fi + 2e —> Fi2). The result is the fonnation of an active pit in the metal, an example of which is illustrated in figure C2.8.6(a) and (b). 

In some metal oxidation reactions (e.g., those involved in Ni-Ni02 electrodes), unwanted 02 is co-evolved. By setting the disk potential to oxidize the Ni and (unintentionally) evolve 02 and the ring current at a potential to reduce Oz, the amount of02 production along with Ni oxidation can be obtained. 

In all cases of electron transport, whether it be hopping, thermal emission, or quantum tunneling, the effect of the electric field in the oxide film is extremely important. In fact, the electric field effect on ion motion is the primary reason the electronic species must be considered at all in most real metal oxidation reactions. This can be understood better when we discuss the coupled-currents approach [10,11] in Sect. 1.15. 

It is interesting to note that the coordinate system utilized throughout this work (and also throughout the majority of the literature on metal oxidation [10]) is not strictly stationary relative to a laboratory coordinate system. The origin x = 0 is always chosen to be at the phase boundary separating the parent metal from the first oxide layer and this phase boundary moves (relative to the center of mass of the parent metal) as layers of the parent metal are consumed in the metal oxidation reaction. Similarly, xt — 0 always occurs at the phase boundary separating layer i — 1 from layer i. 

A second related issue is the asymmetry in the E-i response near Ecelectron transfer reaction that is different from the metal oxidation reaction. Therefore there is no fundamental reason why pa and pc should be equal, and they should be expected to differ. The extent of their difference defines the degree of asymmetry. Asymmetry matters because the extent of the region where Eq. (2) is a good approximation of Eq. (1) then differs for anodic and cathodic polarization (29). The errors in assuming 10 mV linearity using both the tangent to the E-i data at Econ and for +10 or -10 mV potentiostatic polarizations have been defined for different Tafel slopes (30). 

Stoichiometric and catalytic transition-metal oxidation reactions are of great interest, because of their important role in industrial and synthetic processes. The oxidation of alkenes is one of the fundamental reactions in chemistry.1 Most bulk organic products contain functional groups, which are produced in the chemical industry by direct oxidation of the hydrocarbon feedstock. Usually these reactions employ catalysts to improve the yields, to reduce the necessary activation energy and render the reaction more economic. The synthesis of almost every product in chemical industry nowadays employs at least one catalytic step. The oxidation products of alkenes, epoxides and glycols, may be transformed into a variety of functional groups and therefore the selective and catalytic oxidation of alkenes is an industrially important process. 

More recently, higher temperature processes have been considered (at T > 2000 K), such as two-step thermal chemical cycles using metal oxide reactions.2 The first step is solar the endothermic dissociation of the metal oxide to the metal or the lower-valence metal oxide. The second step is non-solar, and is the exothermic hydrolysis of the metal to form FI2 and the corresponding metal oxide. The net reaction is H2O = H2 + 0.5 O2, but since FI2 and O2 are formed in different steps, the need for high-temperature gas separation is thereby eliminated ... 

This reaction, like all metal oxidation reactions, is strongly exothermic. The standard fine energy of this smd many other oxidation reactions are given in Figure 5.4. Written in this form, the distribution coefficient for all metal oxidation reactions is given by... 

Thermal Treatments and Reduction of Adsorbed Species. - After impregnation many catalysts require a thermal treatment and/or a reduction stage to render the adsorbed metal species active. Thermal treatments may take the form of low-temperature drying operations (up to 150 C) simply to remove water, although some decomposition of species such as chloroplatinic acid is known to occur within this temperature range particularly on relatively non-reactive supports such as silica gels [reaction (5)]. Treatments at temperatures between 150 and 500 °C are principally used to decompose the adsorbed species to the metal or metal oxide [reactions (6), (7), and (8)]. 

The Hf-B-C system presents a situation that falls somewhat between the Ti and Zr systems [60]. Although the HfB phase is stable in the Hf-B binary system, it melts at 2100°C, below the melting point of the Hf parent metal (2227°C). During a directed metal oxidation reaction of molten Hf with B4C at just above the melting point of Hf, e.g., 2400°C, the Hf-B-C isothermal ternary cross section (Fig. 18) indicates that the molten metal is... 

Advanced technologies now under consideration for the conversion of methane to more useful materials involve halogenation [1], oxychlorina-tion [2], oxidation, Including oxidative coupling and metal oxide reactions 3], reaction with superacids [4], and various other methods [5],... 

Most metals occur naturally in their oxide or sulfide forms. The process of metal refining converts these ores into pure metals. Thermodynamically, a metal will return spontaneously to its original oxide form. Metal oxidation can occur at high temperatures, by direct reaction with O2, or at a moderate temperature by reaction with water, O2, and/or H+. The latter oxidation, commonly referred to as wet corrosion, has as its basis the combination of electrochemical cathodic reduction and anodic metal oxidation reactions into a corrosion cell. Thus, many corrosion processes are... 

Almost aU metals corrode, but many metals corrode very slowly under normal environmental conditions, due in part to kinetic limitations of the metal dissolution reaction. Thus, the rate of metal corrosion can be anticipated and controlled by developing kinetic rate expressions for metal oxidation reactions. There is a major difference, however, between classical electrochemical metal dissolution kinetics and metal dissolution in a corrosion system, that difference being the occurrence of one or more oxidation and reduction reactions on the same metal. 

To retain emphasis on corrosion processes, Faraday s law will be derived with reference to the generalized metal oxidation reaction, M —> Mm+ + me. In Fig. 4.11, an anodic area, Aa, is shown over which... 

For uniform corrosion, i.e., the entire surface is accessible to the metal oxidation reaction and the environmental reaction(s). Equation 48 becomes... 

IR spectroscopic studies were conducted of the reaction of polyacrylic acid(PAA) and metal oxides (zinc oxide, calcium oxide, cupric oxide, chromium oxide and aluminium oxide). Factors such as the amount of metal oxide, reaction time, solvents, type of metal oxides and temp, were also evaluated to derive the optimum conditions for this reaction. The reactions of chromium oxide and aluminium oxide were far from complete. An extra solvent added to the reaction system could increase the solubility of PAA and metal oxide in the solution to cause complete reaction. The reactivity of the reaction was increased by using a hydrophilic solvent, particularly water and methanol. Furthermore, the reaction rate increased when temp, decreased. The reactivity of the reaction was proportional to the pH value of the metal oxide in the aqueous solution. 16 refs. 

At a given temperature and standard pressure (ambient 1 atm), the oxidation of a metal to its oxide or the reduction of an oxide to its suboxide or the corresponding metal occurs at a well-defined oxygen partial pressure (Po)- At a given temperature, on either side of this unique one of the two coexisting phases must disappear. This is illustrated in the following by considering a hypothetical metal oxidation reaction ... 

If A f/ is negative, heat is given off by the system to the surroundings, as in the exothermic oxidation of iron. The dissolution of metal usually yields more random, less ordered products that is, an increase in entropy is evident. An increase in temperature contributes to a more negative second term, hence, a more negative AG. If AG is negative, the reaction is favored to go spontaneously as written. If it is positive, the reaction will not go as written (although the reverse may). If it is zero, the system is in a state of equilibrium. For the metal oxidation reaction. 

Note that the metal oxidation reaction releases electrons by the reaction MM+ -be. Iron is susceptible to oxidation by reaction (35) or (36), while Ni and stainless steels will corrode by reactions (37) to (40). Another reaction of interest in chlor-alkali operations is oxidation of Fe or Ni in alkaline media ... 

Various theories have been proposed to explain the nature of the irreversible rest potential. Perhaps the oldest one is the oxide theory,which considers the rest potential as corresponding to the equilibrium potential of the metal-metal oxide reaction. This, however, contradicts point (3) above, as well as the fact that this rest potential is established also on electrodes free of bulk oxides such as gold and prereduced Pt. Other theories involve the concept of a mixed potential, which is the result of simultaneous occurrence of two or more continuous (steady state) electrode reactions, i.e., the four-electron cathodic reduction of O2 plus some anodic reaction. The result is a zero overall current and a rest potential in between the equilibrium potentials of these reactions. 

The non-metal oxide reactions with water led Antoine Laurent Lavoisier (1713-1794) to associate the acidity principle with elemental oxygen, but Sir Humphry Davy, in the beginning of the XIX century, was the first to suggest that the acidifying principle was owed to hydrogen — as in the case of hydrochloric acid. 

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