Experimental reaction enthalpies

Table 6.25 Comparison of experimental reaction enthalpies at 0 K (kJ/mol) for the addition of methyl radical to alkenes CH2=CXY with those calculated3 with the wavefunction-based electronic structure methods.

The R2 values refer to the correlation with experimental reaction enthalpies. 

Table 8.6 Experimental reaction enthalpy values and data calculated for some conceivable dehydrogenation processes. ...

Table V-33 Experimental reaction enthalpy values for the formation of Ni(II)-thiocyanate complexes.

Table V-32 Experimental equilibrium data for the Ni(ll) - thiocyanate system. .. 233 Table V-33 Experimental reaction enthalpy values for the formation...

Fig. 1.4 Schematic Born-Haber cycle for the formation of solid NaCI the energetic data (kJ/mol) are Na sublimation enthalpy AHsubi = 100.5 x CI2 dissociation enthalpy H iss = 121.4 Na ionization energy I = 495.7 Cl electron affinity A = -360.5 experimental reaction enthalpy AHr = -411.1.

NIST Chemistry WebBoot, NIST Standard Reference Database Number 69 Mallard, W. G. Linstrom, P. J., Eds. National Institute of Standards and Technology Gaithersburg (http // webbook.nist.gov). The NIST Chemistry WebBook is probably the most extensive of all chemical compilations. It supersedes many of NIST databases and it is composed by several chapters, some of which include thermochemical information of a variety of substances. It is regularly updated, either with new values or with new chapters. Not all of these chapters have thermochemical consistency. Eor instance, the Neutral Thermochemical Data quotes the standard enthalpies of formation directly from the original publications. However, as the experimental reaction enthalpies are also provided, the user can easily derive the correct values. 

Abstract The forward and reverse reactions Br + H2. O HBr + OH are important in atmospheric and environmental chemistry. Five stationary points on the potential energy surface for the Br - - H2O HBr -I- OH reaction, including the entrance complex, transition state, and exit complex, have been studied using the CCSD(T) method with correlation-consistent basis sets up to cc-pV5Z-PP. Contrary to the valence isoelectronic F -I- H2O system, the Br -I- H2O reaction is endothermic (by 31.8 kcal/mol after zero-point vibrational, relativistic, and spin-orbit corrections), consistent with the experimental reaction enthalpy. The CCSD(T)/cc-pV5Z-PP method predicts that the reverse reaction HBr -I- HO Br -I- H2O has a complex... 

Now the vented combustion calorimetry (VCC) method shows that the experimental reaction enthalpy increases about asymptotically with increasing velocity. At low impact speed, the hafnium-based formulation outperforms any other composition and already reaches an efficiency of 52%, whereas the other formulations barely deHver a quarter of the reaction enthalpy. The tantalum-based composition is exceptionally poor as only about 3% of the theoretically stored energy is released. The situation changes at higher speeds (1800-2400 ms ) and shows the gravimetric superiority of Al/PTFE over the other formulations (Figure 13.13). 

Since experimental reaction enthalpies are usually referenced to 298 K, calculated reaction energies are also referenced to that temperature. A o is converted into an energy of reaction at 298 K via equation (16). ... 

Experimental reaction enthalpies have been evaluated via equation (19) for 18 reactions in category 1, 17 reactions in... 

Table 15 9 Experimental reaction enthalpies (kJ/mol). The reactions used for the statistical analysis are given in bold...

In the case of the above hydrogen transfer to benzene, the calculated energy changes are given in Table 2. All of the values are in reasonable agreement with the experimental reaction enthalpy, and the G2 value is in very good agreement. 

It should be noted that the experimental activation enthalpy for the Diels-Alder reaction is 33 kcal/mol (estimated from the reverse reaction and the experimental reaction energy i.e. the MP2/6-31G(d) value is 14kcal/mol too low. Similarly, the calculated reaction energy of —47 kcal/mol is in rather poor agreement with the... 

EXAMPLE 6.7 Determining a reaction enthalpy from experimental data... 

There is also no significant influence of statistic thermodynamical calculations on the reaction parameters. That can be seen in the Tables 3 and 4. In Table 4 the calculated reaction enthalpies and free reaction enthalpies are faced with experimental values estimated by means of thermochemical methods. 

Table 4 shows that the calculated graduation of the reaction enthalpies agrees well with that of the experimental values as well as the free reaction enthalpies. 

Using standard enthalpies from Ref. 132) a reaction enthalpy of 92 kJ mol-1 is the result for reaction (24). This value lies near the above estimated limit and not far from the values which have been calculated from experimental heats of formation for the second and third propagation step (102 kJ mol-1). 

The evaluation of the preexponential term in eq. (100) for gas-phase reactions is straightforward. The absolute zero enthalpy, Eq, can be obtained either from semiempirical calculations on products and reactants or by means of ab initio calculations with a subsequent estimation of the correlation energy. In compromising treatments, the experimentally estimated enthalpies can be employed. 

In equilibrium measurements, there is the possibility of determining the reaction enthalpy AH directly from calorimetry and of combining it with logK (i.e., AG°) to get the reaction entropy, AS . This case, advantageous and simple from the statistical point of view, was only mentioned in a previous paper (149). Since that time, this experimental approach has been widely used (59, 62-65, 74-78, 134, 137, 138, 210, 211) hence, a somewhat more detailed mathematical treatment seems appropriate. 

The standard deviation between experimental and calculated heats of reaction are between 0.5 and 1 kcal/mol for those classes of compounds where enough experimental heats of formation are available to allow a full parameterization. For those classes of compounds where insufficient heats of formation are known to allow the determination of all parameters for 1,2- and 1,3-interactions, an estimate can be given for the bond energy terms which are the dominating parameters. Even here, therefore, a reasonable value for the reaction enthalpy is available. 

Experimentally, the enthalpy change associated with the formation reaction... 

There is general agreement that static-bomb combustion calorimetry is inherently unsatisfactory to determine enthalpies of formation of organolead compounds2,3. Unfortunately, as shown in Table 6 only three substances have been studied by the rotating-bomb method. The experimentally measured enthalpies of formation of the remaining compounds in Table 6 were determined by reaction-solution calorimetry and all rely on AH/(PbPh4, c). 

Two sources to obtain this necessary information are the use of data bases and through experimental determinations. Enthalpies of reaction, for example, can be estimated by computer programs such as CHETAH [26, 27] as outlined in Chapter 2. The required cooling capacity for the desired reactor can depend on the reactant addition rate. The effect of the addition rate can be calculated by using models assuming different reaction orders and reaction rates. However, in practice, reactions do not generally follow the optimum route, which makes experimental verification of data and the determination of potential constraints necessary. 

One of the issues of the industrial process design is related to the heat released by this reaction. A temperature rise will decrease the acetic acid yield, not only because the equilibrium constant becomes lower (the reaction is exothermic see section 2.9) but also because it will reduce the enzyme activity. It is therefore important to keep the reaction temperature within a certain range, for instance, by using a heat exchanger. However, to design this device we need to know the reaction enthalpy under the experimental conditions, and this quantity cannot be easily found in the chemical literature. 

Figure 2.1 Thermochemical cycle, showing how to relate the enthalpy of the experimental reaction 2.1 with reaction 2.2, where reactants and products are in their standard states.

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