Nitric oxide formation thermodynamics

Thermodynamic calculation results are shown in Table 4.1. For reaction (5), the main parameters are the following free energy variation 5165 kJ, equilibrium constant at 600 °C 3.4 10-3 and the reagent conversion to reaction products is negligibly low. Much less favorable is the equilibrium state in the reaction (6). Therefore, both reactions are not practically executed. Reaction (6) described in the monograph by Zeldovich el al. [39] and in the article by Anbar [40] runs at a temperature above 1273 K with nitric oxide formation by the mechanism, which includes elementary stages with atomic oxygen participation. However, atomic... 

Thus far we have studied thermodynamics and kinetics imder the assumption that the systems of interest are in equilibrium. However, some natural systems have reaction rates so slow that they exist for long periods under non-equilibrium conditions. The formation of nitric oxide serves as an interesting example. 

From the thermodynamic data of Appendix C, show that the product of the reaction of ammonia gas with oxygen would be nitrogen, rather than nitric oxide, under standard conditions and in the absence of kinetic control by, for example, specific catalysis of NO formation by platinum. (Assume the other product to be water vapor.)... 

According to these data the heat of activation for the decomposition of nitric oxide, to which reaction the factor k refers, is A = 82 10 kcal/mole.10 It should be especially noted that there is no systematic divergence between the data on the formation and on the decomposition of nitric oxide this fact justifies the assumption that the rate of decomposition is directly proportional to the square of the nitric oxide concentration.11 The investigation covered the temperature range from 2000°K to 2900°K in which the rate varies by a factor of 300. As appears from Fig. 13, except for the scattering due to the inevitable errors of the experiments and computations, the points actually do fall on a straight line in the coordinates lg kr, 1/Tm, i.e., the Arrhenius temperature dependence of the reaction rate holds. The thermodynamic relation between the rates of the direct and reverse reactions permits determining the heat of activation A for the formation of nitric... 

Recent investigations have shed light on peculiarities of the NOS action mechanism the role of the H4B cofactor and CaM, and cooperativity in kinetic and thermodynamic properties of different components of the nitric oxide synthesis system. Stop flow experiments with eNOS (Abu-Soud et al., 2000) showed that calmodulin binding caused an increase in NADH-dependent flavin reduction from 0.13 to 86 s 1 at 10 °C. Under such conditions, in the presence of Arg, heme is reduced very slowly (0.005 s 1). Heme complex formation requires a relatively high concentration ofNO (>50 nM) and inhibits the entire process NADH oxidation and citrulline synthesis decreases 3-fold and Km increases 3-fold. NOS reactions were monitored at subzero temperatures in the presence of 50% ethylene glycol as an anti-freeze solvent (Bee et al., 1998). 

The formation of the nitrogen oxides, nitric oxide (NO) and nitrogen dioxide (NO2) from N2 and O2, provides an example of the interplay between thermodynamics and kinetics. The calculation of the equilibrium concentrations of NO and NO2 is straightforward if we assume that these are the only compounds formed and that the initial pressures of N2 and O2 are known. 

In view of the fact that pharmacological effects of nitroprusside, [Fe(CN)5NO], a widely recognized hypotensive agent (61—65), have been attributed to the release of nitric oxide from its reduced form, i.e., [Fe(CN)5NO], the kinetic and thermodynamic properties of both nitrosyl complexes of pentacyanoferrate-(II) and -(III) have attracted considerable attention in the past two decades (66,67). In this context, the formation of [Fe(CN)5NO] and [Fe(CN)5NO] in the direct reactions of [Fe(CN)5(H20)] and [Fe(CN)5(Fl20)] with nitric oxide, respectively, was subjected to detailed kinetic and mechanistic investigations (68-70). As presented below, the results of these studies allowed to draw valuable conclusions concerning the validity of the mechanism of NO release from nitroprusside postulated in the Hterature. 

Industrial fertilizer synthesis starts from ammonia synthesis, and ammonia is then easily oxidized in a separate reactor to nitric oxide over PtRh wire gauze catalyst. Formation of nitric acid requires further oxidation of nitric oxide to nitrogen dioxide (NO2) and absorption of the nitrogen dioxide in water. Overall, three different chemical process plants are used for the synthesis of nitric acid. The ammonia synthesis reaction takes place in a high-tem-perature, high-pressure reactor that requires recycling of products due to the thermodynamic limitations of chanical conversion. The ammonia oxidation reaction is very fast and takes place over a very small reactor length. Finally, nitric acid synthesis takes place in absorption columns. 

Nitric acid is one of the three major acids of the modem chemical industiy and has been known as a corrosive solvent for metals since alchemical times in the thirteenth centuiy. " " It is now invariably made by the catalytic oxidation of ammonia under conditions which promote the formation of NO rather than the thermodynamically more favoured products N2 or N2O (p. 423). The NO is then further oxidized to NO2 and the gases absorbed in water to yield a concentrated aqueous solution of the acid. The vast scale of production requires the optimization of all the reaction conditions and present-day operations are based on the intricate interaction of fundamental thermodynamics, modem catalyst technology, advanced reactor design, and chemical engineering aspects of process control (see Panel). Production in the USA alone now exceeds 7 million tonnes annually, of which the greater part is used to produce nitrates for fertilizers, explosives and other purposes (see Panel). 

The Af-HjO diagrams present the equilibria at various pHs and potentials between the metal, metal ions and solid oxides and hydroxides for systems in which the only reactants are metal, water, and hydrogen and hydroxyl ions a situation that is extremely unlikely to prevail in real solutions that usually contain a variety of electrolytes and non-electrolytes. Thus a solution of pH 1 may be prepared from either hydrochloric, sulphuric, nitric or perchloric acids, and in each case a different anion will be introduced into the solution with the consequent possibility of the formation of species other than those predicted in the Af-HjO system. In general, anions that form soluble complexes will tend to extend the zones of corrosion, whereas anions that form insoluble compounds will tend to extend the zone of passivity. However, provided the relevant thermodynamic data are aveiil-able, the effect of these anions can be incorporated into the diagram, and diagrams of the type Af-HjO-A" are available in Cebelcor reports and in the published literature. 

When some metals or alloys are immersed in an oxidizing solution, corrosion does not occm, although it is favored thermodynamically. For example, iron will dissolve in dilute, but not in concentrated nitric acid. Corrosion is prevented because of the formation of a protective oxide film on the metal surface. This loss in chemical reactivity of a metal is known as passivation. Metals which readily undergo passivation in damp air include Fe, Al, Cr, Ni, Ti, and Pt. 

The thermodynamics of nitric acid production based on ammonia can be characterized as follows N2 and NO (and also N2O) are thermodynamically favored products of NH3 oxidation. The undesirable formation of N2 is thermodynamically more favored than the desired NO formation. The selectivity of the catalyst is therefore important in suppressing N2 formation. NO may decompose to N2 and O2, but this can be avoided by quenching the product gas of... 

The standard potential is shifted by - 2.28 V to - 0.60 V so that gold may be oxidized easily by oxygen and will therefore dissolve in cyanide solution with oxygen access. A similar situation holds when gold is dissolved in a 1 3 mixture of hydrochloric and nitric acid. In this case Au " " is com-plexed by Cl" leading to a very small Au " concentration in solution. With A i,=2.2x 10" for the AuClJ complex, the standard potential of the Au/Au electrode is shifted from 1.42 V to 0.994 V. Apparently, the formation of a thermodynamically stable AUCI4 complex leads to the dissolution of gold, which does not occur in pure nitric acid. 

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