Bond enthalpies calculation

Draw the Lewis structure for the hypothetical molecule N6, consisting of a six-membered ring of nitrogen atoms. Using bond enthalpies, calculate the enthalpy of reaction for the decomposition of N6 to N2(g). Do you expect N6 to be a stable molecule ... 

The bond breaking and making processes can be represented by a Hess s law cycle (see Figure 5.29). This calculated value is slightly different from the real value because the bond enthalpies are averages. In addition, the water that is produced (in the calculation) is not in its standard state, as a liquid. The gaseous state is always used when performing bond enthalpy calculations. 

Calculate the bond enthalpy of the C—C bond in ethane using only the enthalpies of atomization of methane and ethane. Compare this result with the accepted result. 

Once the BEs and SBEs have been decided upon, the normal functioning of the MM program causes each bond to be multiplied by the number of times it appears in the computed molecule to find its contribution to the total bond enthalpy. In ethylene, 26.43 + 4(—4.59) = 8.07kcalmol . In Eile Segment 5-1, this sum is denoted BE. This whole procedure is essentially a conventional bond energy calculation. 

A force field that can produce vibrational spectra has a second advantage in that the Ay// calculations can be put on a much more satisfactory theoretical base by calculating an enthalpy of formation at 0 K as in ab initio procedures and then adding various thermal energies by more r igorous means than simply lumping them in with empirical bond enthalpy contributions to Ay//-. The stronger the theoretical base, the less likely is an unwelcome surprise in the output. 

There is a lively controversy concerning the interpretation of these and other properties, and cogent arguments have been advanced both for the presence of hydride ions H" and for the presence of protons H+ in the d-block and f-block hydride phases.These difficulties emphasize again the problems attending any classification based on presumed bond type, and a phenomenological approach which describes the observed properties is a sounder initial basis for discussion. Thus the predominantly ionic nature of a phase cannot safely be inferred either from crystal structure or from calculated lattice energies since many metallic alloys adopt the NaCl-type or CsCl-type structures (e.g. LaBi, )S-brass) and enthalpy calculations are notoriously insensitive to bond type. 

This limitation explains why, whenever possible, we use enthalpies of formation (AFff) rather than bond enthalpies to calculate the value of AH for a reaction. Calculations involving enthalpies of formation are expected to be accurate within 0.1 kj the use of bond enthalpies can result in an error of 10 kj or more. 

A note on good practice. The use of mean bond enthalpies is hazardous because actual bond enthalpies often differ considerably from mean values. The modem procedure for estimating a reaction enthalpy is to use commercial software to calculate the enthalpies of formation of the reactants and products and then to take the difference, as in Section 6.18. 

Benzene is more stable and less reactive than would be predicted from its Kekule structures. Use the mean bond enthalpies in Table 6.8 to calculate the lowering in molar energy when resonance is allowed between the Kekule structures of benzene. 

Thus, the statishcal criteria for Eqs. (5) and (6) show that the approach provides a reasonably reliable method to calculate H-bond enthalpy and free energy. 

Fig. 29. Bond breakages calculated by EROS for the reaction of hydroxynorbomene 4 with base. Enthalpies are in kcal/mol in parentheses, relative to 4...

Calculate the heats of formation for gaseous XeF2 and XeF4 if the Xe-F bond enthalpy is 133 kj mol-1 and the F-F bond enthalpy is 159 kj mol-1. 

We now turn to the calculation of bond enthalpy terms. In a first approximation, the enthalpy required to break all the chemical bonds in, say, gaseous GeR4, the so-called... 

The calculation of other bond enthalpy terms, such as E (Ge—Ge), E (Ge—O), E (Ge—N) and E(Ge—S), can be made from data in Table 1. However, due to the above-mentioned controversy involving most of the data obtained with static-bomb combustion calorimeters, we refrain from tabulating those terms. 

Despite the previous remarks, organotin compounds are one of the few families of organometallic substances for which the thermochemical data justify the calculation of bond enthalpy terms. Some term values (consistent with the terms recommended by Pilcher and Skinner for hydrocarbons)27 are given in Table 4. A few words on this selection are, however, appropriate. 

Another example illustrating the apparent problems with the enthalpy-of-formation data for some tin compounds concerns the term (Sn—Sn). In order to avoid the complication related to the term (Sn—Cb), the bond enthalpy (Sn—Sn) in Table 4 was calculated from A/ °[(SnMe3)2, g]. As a confirmation, one can then derive the value for the same term from A ff °[(SnEt3)2, g]—they should be similar. Yet, this calculation yields (Sn—Sn) = 223.18 kJmol-1, a value which looks unreasonably high and questions the experimental value of A °[(SnEt3)2, g] in Table 3. 

To calculate the average bond enthalpy of the (C-H) bond in methane, E (C-H), the average of the four bond dissociation enthalpies is calculated ... 

---