Reaction thermodynamics with soluble oxidants

The results of simultaneous weighted least-squares regression of the data and some of the unfitted but derived quantities are shown in Tables la, lb, and 2. Table la displays elemental arsenic, its simple oxides, and the reactions for arsenic oxidizing to arsenic trioxides. Table lb introduces the hydrolysis species for As(III) and As(V) in solution, the hydrolysis reactions, and the solubility reactions for the simple oxides. Single species are shown at the top of each table with the reactions underneath. The following discussion describes some of the mineral occurrences for these substances, describes their relative stabilities from field observations, and considers the implications of the evaluated thermodynamic data in terms of these occurrences. 

A central issue in the attempt to establish a reliable database is the requirement of critically evaluated thermodynamic data for several key species. One such pivotal element is aluminum, which has an extensive literature of solubility and thermochemical data from which to choose, for each of the aqueous species or complexes. The aluminum species are fundamental to the calculation of solubility and reaction state with respect to many silicates and aluminum oxides and hydroxides and are principal components in numerous surface chemical reactions in the environment. Two key chapters in this volume address this fundamental problem Apps and Neil give a critical evaluation of the data for the aluminum system and Hem and Roberson present the kinetic mechanisms for hydrolysis of aluminum species. 

Hydrocarbons and most other organic compounds in the atmosphere are thermodynamically unstable toward oxidation and tend to be oxidized through a series of steps. The oxidation process terminates with formation of CO2, solid organic particulate matter that settles from the atmosphere, or water-soluble products (for example, acids, aldehydes), which are removed by rain. Inorganic species such as ozone or nitric acid are byproducts of these reactions. 

The data given in Tables 1.9 and 1.10 have been based on the assumption that metal cations are the sole species formed, but at higher pH values oxides, hydrated oxides or hydroxides may be formed, and the relevant half reactions will be of the form shown in equations 2(a) and 2(b) (Table 1.7). In these circumstances the a + will be governed by the solubility product of the solid compound and the pH of the solution. At higher pH values the solid compound may become unstable with respect to metal anions (equations 3(a) and 3(b), Table 1.7), and metals like aluminium, zinc, tin and lead, which form amphoteric oxides, corrode in alkaline solutions. It is evident, therefore, that the equilibrium between a metal and an aqueous solution is far more complex than that illustrated in Tables 1.9 and 1.10. Nevertheless, as will be discussed subsequently, a similar thermodynamic approach is possible. 

To examine the situation with alloys in a little more detail, the Cu-Ni alloys will first be considered. Here the mutual solubility of the two oxides NiO and CU2O can probably be neglected, and these are the only two possible oxidation products. Assume for simplicity that the alloy is thermodynamically ideal, and let and Xn be the mole fractions in the alloy. Consider the reactions... 

Equilibria involving reductive dissolution reactions add to the complexity of mineral solubility phenomena in just the way that pE-pH diagrams are more complicated than ordinary predominance diagrams, like that in Fig. 3.7. The electron activity or pE value becomes one of the master variables whose influence on dissolution reactions must be evaluated in tandem with other intensive master variables, like pH or p(H4Si04). Moreover, the status of microbial catalysis under the suboxic conditions that facilitate changes in the oxidation states of transition metals has to be considered in formulating a thermodynamic description of reductive dissolution. This consideration is connected closely to the existence of labile organic matter and, in some cases, to the availability of photons.26... 

Besides the effect of the electrode materials discussed above, each nonaqueous solution has its own inherent electrochemical stability which relates to the possible oxidation and reduction processes of the solvent,the salts, and contaminants that may be unavoidably present in polar aprotic solutions. These may include trace water, oxygen, CO, C02 protic precursor of the solvent, peroxides, etc. All of these substances, even in trace amounts, may influence the stability of these systems and, hence, their electrochemical windows. Possible electroreactions of a variety of solvents, salts, and additives are described and discussed in detail in Chapter 3. However, these reactions may depend very strongly on the cation of the electrolyte. The type of cation present determines both the thermodynamics and kinetics of the reduction processes in polar aprotic systems [59], In addition, the solubility product of solvent/salt anion/contaminant reduction products that are anions or anion radicals, with the cation, determine the possibility of surface film formation, electrode passivation, etc. For instance, as discussed in Chapter 4, the reduction of solvents such as ethers, esters, and alkyl carbonates differs considerably in Li or in tetraalkyl ammonium salt solutions [6], In the presence of the former cation, the above solvents are reduced to insoluble Li salts that passivate the electrodes due to the formation of stable surface layers. However, when the cation is TBA, all the reduction products of the above solvents are soluble. 

A thermodynamic model of dissolution is presented in this chapter, which relates the solubility product constant to the thermodynamic potentials and measurable parameters, such as temperature and pressure of the solution. The resulting relations allow us to develop conditions in which CBPCs are likely to form by reactions of various oxides (or minerals) with phosphate solutions. Thus, the model predicts formation of CBPCs. 

Mn (II) is the most important soluble form, generally stable in the pH range 6-9 in natural waters (Morgan and Stumm, 1965). Since Mn (III) is thermodynamically unstable in a desert environment lacking abundant humic acids, Mn (IV) is the primary insoluble oxyhydroxide found in varnish, with average valencies of about +3.8 to +3.9 (Potter and Rossman, 1979a McKeown and Post, 2001). The general Mn (II) oxidation reaction is ... 

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