Ethylene bond energy

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. 

The double bond in ethylene is stronger than the C—C single bond in ethane but It IS not twice as strong Chemists do not agree on exactly how to apportion the total C=C bond energy between its ct and rr components but all agree that the rr bond is weaker than the ct bond... 

The simplest arithmetic ap proach subtracts the C—C (j bond energy of ethane (368 kj/mol 88 kcal/mol) from the C=C bond energy of ethylene (605 kJ/mol 144 5 kcal/mol) This gives a value of 237 kJ/mol (56 5 kcal/mol) for the tt bond energy... 

C06-0100. Use average bond energies (Table 6-2) to compare the combustion energies of ethane, ethylene, and acetylene. Calculate which of these hydrocarbons releases the most energy per gram. 

C06-0106. Use tabulated bond energies in Tabie 6 to estimate the energy change when HCl adds to ethylene (H2 C I CH2) to produce CH3 CH2 Cl and when CI2 adds to ethylene to produce... 

The simple harmonic oscillator picture of a vibrating molecule has important implications. First, knowing the frequency, one can immediately calculate the force constant of the bond. Note from Eq. (11) that k, as coefficient of r, corresponds to the curvature of the interatomic potential and not primarily to its depth, the bond energy. However, as the depth and the curvature of a potential usually change hand in hand, the infrared frequency is often taken as an indicator of the strength of the bond. Second, isotopic substitution can be useful in the assignment of frequencies to bonds in adsorbed species, because frequency shifts due to isotopic substitution (of for example D for H in adsorbed ethylene, or OD for OH in methanol) can be predicted directly. 

It is interesting to note that methane, ethane, and ethylene are all gases hexane, octane, and nonane are all liquids (at room conditions) while low molecular weight PE is a waxy solid. This trend is primarily due to an increase in the mass per molecule and to an increase in the London forces per polymer chain. The London force interaction between methylene units is about 8 kcal/mol. Thus, for methane molecules the attractive forces are 8 kJ/mol for octane it is 64 kJ/mol and for PE with 1000 ethylene (or 2000 methylenes) it is 2000 methylene units X 8 kJ/mol per methylene unit = 16,000 kJ/mol, which is well sufficient to make PE a solid and to break backbone bonds before it boils. (Polymers do not boil because the energy necessary to make a chain volatile is greater than the primary backbone bond energy.)... 

Consider the effect of monomer structure on the enthalpy of polymerization. The AH values for ethylene, propene, and 1-buene are very close to the difference (82-90 kJ mol 1) between the bond energies of the re-bond in an alkene and the a-bond in an alkane. The AH values for the other monomers vary considerably. The variations in AH for differently substituted ethy-lenes arise from any of the following effects ... 

However, direct calculations of accurate bond energies represent a major challenge. Examples are given [13,14] where the ratios of carbon-carbon bond energies, relative to that of ethane, were successfully calculated for ethylene, acetylene, benzene, and... 

Despite the marked differences in both geometric parameters and the SCF Ar values between the molecules involved in this comparison, there are striking regularities the F value calculated for propene is, for all practical reasons, th that of ethylene (which takes care of 3 CH bonds) plus th that of tetramethylethylene (for the CC bond). Capitalizing on this idea, we may well consider transferable bond contributions modeled after Eq. (11.12) and use them to generate new reference bond energies satisfying Eq. (10.36). 

The problems to be solved are best illustrated by a typical example. Take sj, the carbon-carbon bond energy in ethane with Rqc = 1.531 A calculate the energy of a C(sp ) C(sp ) bond like that found in olefins. The latter is for reference charges, which are 35.1 me for the sp carbon (as in ethane) and 7.7 me for the sp carbon (as in ethylene). This transformation is schematically represented in Fig. 11.1. 

The CC and CH bonds of ethane (Example 10.1), and the final selection See = 69.633 and 8ch = 106.806 kcal/mol, are used to get the CC and CH bonds found in unsaturated hydrocarbons by retaining both the contribution of Fkh Eq. (11.12), and the effect of charge variations described by Eq. (10.37). The reference CC double bond of ethylene and the reference CC bonds of benzene, however, roughly estimated along the lines described in Example 10.1, are deduced from the appropriate CH bond energies and the energy of atomization of the corresponding molecule, AE, obtained from experimental data. 

Somewhat analogous reactions would be expected for the reaction of ethylene with 02 ions but the observed reaction rate is lower than for propene, suggesting that the reaction pathway may be controlled by the C—H bond energies. For reactions of propane and 1-butene with 02, oxygenated compounds of the same carbon number as the reactants were produced. The initial step is thought to involve a hydrogen atom abstraction from a secondary carbon atom. 

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