The question posed in this slide is how good are the values from spectroscopy? The heat of atomization of CH4, or how much energy it takes to separate methane into its individual constituent atoms, is 397.5 kcal/mol. Heat of atomization is calculated using values for heat of combustion and heat of formation (see the previous lecture's discussion on the heat of atomization of cyclohexane for an example). Thus, knowing this heat of atomization, one can calculate the average bond energy for the C-H bond in methane by dividing the total atomization energy by how many bonds there are (4) to get 99.4 kcal/mol.
However, do all or even any of the bonds in methane exhibit this "average" bond energy? That is, does it take 99.4 kcal/mol of energy to remove each hydrogen from CH4? Barney Ellison (who may come speak to us in the spring!) did an experiment to calculate the actual bond dissociation energies, or how much energy is needed to separate successive C-H bonds in methane. This was done through spectroscopy, where light energy is used to split the molecule into its constituents one hydrogen at a time.
As the table in Frame 9 shows, none of the bonds in methane exhibit the average bond energy taken from simply dividing the total by how many bonds there are. No individual bond was equivalent to the average. The dissociation energy for CH3-H was 104.99 +/- 0.03 whereas C-H was 80.9 +/- 0.2. This is partly because of changes in hybridization as hydrogens are taken off the molecule. However, if one adds all of the bond dissociation energies of methane up, one gets 397.5, which is the total heat of atomization. Thus, the spectroscopy experiments conducted by Ellison are quite accurate.
Next, we shall look at whether we can use these calculated Average Bond Energy values to calculate heats of atomization for other molecules, and whether these calculations are actually useful.
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