Sulfur dioxide thermodynamic data

The thermodynamic properties of sulfur trioxide, and of the oxidation reaction of sulfur dioxide are summarized in Tables 3 and 4, respectively. Thermodynamic data from Reference 49 are beheved to be more accurate than those of Reference 48 at temperatures below about 435°C. 

Flame Temperature. The adiabatic flame temperature, or theoretical flame temperature, is the maximum temperature attained by the products when the reaction goes to completion and the heat fiberated during the reaction is used to raise the temperature of the products. Flame temperatures, as a function of the equivalence ratio, are usually calculated from thermodynamic data when a fuel is burned adiabaticaHy with air. To calculate the adiabatic flame temperature (AFT) without dissociation, for lean to stoichiometric mixtures, complete combustion is assumed. This implies that the products of combustion contain only carbon dioxide, water, nitrogen, oxygen, and sulfur dioxide. 

Sulfur dioxide hydrate, thermodynamic data and lattice constants, 8 Sulfur hexafluoride (SFa), hydrate, 22, 47... 

The sulfur dioxide enters the reactor with an initial concentration of 10% by volume, the remainder being air. At the exit of the first bed, the temperature is 620°C. Assume ideal gas behavior, the reactor operates at 1 bar and R = 8.3145 kJ-Kr -lonoD1. Assume air to be 21% 02 and 79% N2. Thermodynamic data at standard conditions at 298.15 K are given in Table 6.186. 

Given in the literature are vapor pressure data for acetaldehyde and its aqueous solutions (1—3) vapor—liquid equilibria data for acetaldehyde—ethylene oxide [75-21-8] (1), acetaldehyde—methanol [67-56-1] (4), sulfur dioxide [7446-09-5]— acetaldehyde—water (5), acetaldehyde—water—methanol (6) the azeotropes of acetaldehyde—butane [106-97-8] and acetaldehyde—ethyl ether (7) solubility data for acetaldehyde—water—methane [74-82-8] (8), acetaldehyde—methane (9) densities and refractive indexes of acetaldehyde for temperatures 0—20°C (2) compressibility and viscosity at high pressure (10) thermodynamic data (11—13) pressure—enthalpy diagram for acetaldehyde (14) specific gravities of acetaldehyde—paraldehyde and acetaldehyde—acetaldol mixtures at 20/20°C vs composition (7) boiling point vs composition of acetaldehyde—water at 101.3 kPa (1 atm) and integral heat of solution of acetaldehyde in water at 11°C (7). 

From the principles of thermodynamics and certain thermodynamic data the maximum extent to which a chemical reaction can proceed may be calculated. For example, at 1 atm pressure and a temperature of 680°C, starting with 1 mole of sulfur dioxide and mole of oxygen, 50% of the sulfur dioxide can be converted to sulfur trioxide. Such thermodynamic calculations result in maximum values for the conversion of a chemical reaction, since they are correct only for equilibrium conditions, conditions such that there is no further tendency for change with respect to time. It follows that the net rate of a chemical reaction must be zero at this equilibrium point. Thus a plot of reaction rate [for example, in units of g moles product/(sec) (unit volume reaction mixture)] vs time would always approach zero as the time approached infinity. Such a situation is depicted in curve A of Fig. 1-1, where the rate approaches zero asymptotically. Of course, for some cases equilibrium may be reached more rapidly, so that the rate becomes almost zero at a finite time, as illustrated by curve B. 

Quantitative analysis of different reaction pathways for the transformation of aquated sulfur dioxide in atmospheric droplet systems has been a major objective of the research conducted in the principal investigator s laboratory for the last four years. Available thermodynamic and kinetic data for the aqueous-phase reactions of SO2 have been incorporated into a dynamic model of the chemistry of urban fog that has been developed by Jacob and Hoffmann (23) and Hoffmann and Calvert (39). The fog and cloud water models developed by them are hybrid kinetic and equilibrium models that consider the major chemical reactions likely to take place in atmospheric water droplets. Model results have verified that... 

The efficient design of processes for removal of sulfur dioxide resulting from coal and oil combustion requires thermodynamic and kinetic data for the various materials that might be used in the processes. Examination of the available thermodynamic data for sulfur compounds indicates serious uncertainties and the present review is a step toward providing the best set of internally consistent values obtainable from the literature and from current experiments. 

The XPS results indicated that there were about 3-5 at. % sulfur and 27-47 at. % oxygen incorporated onto the sulfur dioxide plasma treated LDPE substrate surfaces (Table 1). The sulfur atomic concentration reached a maximum at about 50 A from the sample surface (0 = 30°) right after the plasma treatment (Figure 3). (The uncertainty of the XPS multiplex scan for atomic concentration analysis is believed to be 0.5-1.0 at. %.) The sulfur-containing species diffused into the bulk of the polymer (> 100 A) as shown from the XPS data collected eleven days after the plasma reaction. This phenomenon is due to the mobility of the polymer surface. s After the sulfur dioxide plasma modification, the hydrophilic sulfonyl groups on the LDPE backbone diffuse away from the polymer surface toward the bulk of the material so that a lower surface energy can be attained. Because the air/LDPE interface has a low surface tension, thermodynamic equilibrium favors a hydrophobic surface. As a result, the sulfur atomic concentrations in the top 100 A of the substrates decreased with time as the sulfonyl groups diffused away from this surface layer. 

Simenson (1940), and the thermodynamic data for the system SO2-H2O are given by Plummer (1950). In water sulfur dioxide exists as the dissolved gas, as undissociated sulfurous acid, H2SO3, as the bisulfite ion, HSOj , and as the sulfite ion, SO3—. Sulfurous acid is a rather weak dibasic acid its first ionization constant is 1.7 X 10 , and the second is 5 X 10 . At 25° C. the corresponding pK values are 1.8 and 5.3, respectively. The calculated distribution of the various ionic species in aqueous solution of sulfurous acid in the range of pH of 0 to 8 is shown in Fig. 1. At pH s over 9.5 only SO3 ions exist, in the range of pH 9.5 to 4.5 both SO3 and HSOs occur, and at pH 4.5 and lower, SO3— no longer exists in appreciable concentrations. 

Plummer, A. W. 1950. Thermodynamic data for the system sulfur dioxide water. 

On the basis of thermodynamic constants obtained for hydroxide compounds of iron with different aging time and also of experimental data, the physicochemical character of the diagenetic transformations of iron sediments of various compositions (oxide, silicate, carbonate, sulfide) can be traced. The results obtained are represented graphically in the form of stability diagrams of iron compounds as a function of variations in the main parameters governing the physicochemical character of the environment of diagenesis—pH, Eh, activity of iron and dissolved forms of sulfur and carbon dioxide. 

Early investigatioiis of the photochemistry of a-heteroatom substituted carboxylates 40 (Scheme 19) concentrated on the operation of SET processes. These substances possess two potential electron donor sites. SET from the carboxylate moiety leads to formation of an acyloxy radical 41 while SET from the heteroatom center gives rise to a zwitterionic radical 42. Subsequent loss of carbon dioxide from either of these intermediates yields the same carbon-centered, heteroatom-stabilized radical 43. The critical factors determining which of these pathways is responsible for the SET-promoted decarboxylation reactions of substrates in this family is (1) the relative oxidation potentials of the two potential donor moieties and (2) the relative rates of decarboxylation of intermediates 41 and 42. Redox potential data indicate that electron transfer from sulfur (E(-t) = ca. 0.4 V) and nitrogen (E(-l-) = ca. 0.7 V) centers to acceptors should be thermodynamically more favorable than from carboxylate groups (E(-t) = ca. 1.4 V). 

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