Decomposition enthalpy

Considering the low energy level of elemental nitrogen, the decomposition enthalpy ofp-ethoxyphenylpentazole (5.4 kcal/mole) indicates the high resonance energy of the pentazole system. 

Equation 1.34 is plotted for a number of hydrides in Fig. 1.25. As can be seen all the data points fit very well in a simple straight line whose slope is equal to AS -130 J mol" K [162]. This clearly shows that the entropy term is, indeed, a nearly constant value for all the solid state hydrogen systems. Figure 1.25 also shows that a low desorption temperature at 1 atm of pressure (more or less an operating pressure of a PEMFC) can only be achieved with hydrides having the forma-tion/decomposition enthalpies not larger than 50 kJ moF. For example, hydrides that desorb at room temperature such as LaNi and TiFe have AH 30 and 33.3 kJ mol", respectively [163]. However, too small an enthalpy term would require at 1 atm to be much below 0°C. From this point of view the enthalpy term is one of the most important factors characterizing any hydride. 

From the Van t Hoff plot the decomposition enthalpy change AH) and entropy change (AS) for the first and second step were calculated to be AH j = 93.9 kJ/mol-Hj and AS, = 116.2 J/mol-H K, and AH = 102.2 kJ/mol-H and = 125.9 1/ mol-HjK. Taking into account the very high desorption temperature of this compound, much higher than that of a catalyzed MgH, and the fact that MgH is used in its synthesis by ball milling, the compound is not at all competitive to MgH. ... 

Measurement of the decomposition enthalpy demonstrates the stronger binding in the cluster compound relative to the bulk, in agreement with the shorter gold-gold distance observed by EXAFS. 

Admittedly, we are not surprised that the decomposition enthalpies of styrene and 2-vinylnaphthalene peroxides are close to each other, as are those of the isomeric isopropenyl species. There is a rather rehable constant enthalpy of formation difference between phenyl and naphthyl derivatives, and as a corollary, a near-equaUty of the enthalpies of formation of 1- and 2-naphthyl derivatives, cf. the combined calculational and calorimetric studies of M. V. Roux, M. Temprado, R. Notario, S. P. Verevkin, V. N. Emel yanenko, D. E. DeMasters and J. F. Liebman, Mol. Phys., 102, 1909 (2004) and references cited therein. It is perhaps more surprising that the a-methyl group on the unsaturated moiety (vinyl -> isopropenyl) causes such a small change. 

A range of heterocyclic azides were studied by differential scanning calorimetry, the enthalpies of thermal decomposition were lower than might have been expected, from 0.3—1 kJ/g. If ortho substituents onto which the azide could cyclise were present, decomposition enthalpies were, as one would expect, lower still [5], as well as the individually indexed compounds ... 

Tritylation of some compounds of low thermal stability increases neither the kinetic nor thermodynamic stability, measured decomposition enthalpies per gram remaining high. Possibly the energy theoretically represented by the three benzene rings in the trityl group makes itself apparent. 

Taking the concentration into account, the decomposition enthalpy should be 2500kj - mob1. Such a high decomposition enthalpy cannot be found in the table. Thus, the severity of a eventual decomposition reaction is assessed to be low. 

A major difference with desired reactions is that the stoichiometry is often unknown, that is, the decomposition products are unknown. The reason is that decomposition reactions are often affected by the triggering conditions and thus often run along different reaction paths. This is a major difference compared to a total combustion, for example. The consequence is that the decomposition enthalpy cannot be predicted using standard enthalpies of formation AHjj taken from, for example, tables or estimated by group increment methods, such as Benson groups [3, 4] ... 

Table 11.1 Typical values of decomposition enthalpies for different functional groups.

The respective decomposition enthalpies are dependent on the amount of hydrogen substituted by fluorine in the range 25-65 kj (mol H2) compared to 78 kj (mol H2) or, after taking zero-point energy corrections into account, 13-59 kJ (mol H2) compared to 61 kJ (mol H2) for the pure LiBH4. 

The chemical structure of the polymer is diagramed in Scheme 9. The average molecular weight is about 71,000 g/mole, corresponding to about 190 monomer units (n in Scheme 9). It should be noted that the decomposition of the monomer is exothermic with a decomposition enthalpy of —624 J/g [122]. The absorption coefficient at 248 nm is -66,000 cm-1, which is quite high nevertheless, it is not at the maximum (typically at wavelengths in the 310-350 nm region). 

Increase of the decomposition enthalpy with T (for reactions with 8.2... 

Table 16.11 The initial data and results of determination of the decomposition enthalpy of Pb304 by the third-law method...

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