[thetada]

Is anyone aware of a value for the mean bond dissociation enthalpy for aromatic carbon carbon bonds, such as the C-C bond in benzene? Data tables tend only to have C-C bonds and C=C bonds. Obviously I could use three of each to calculate the enthalpy of, say, combustion of benzene. I'm just wondering if more accurate data are available.

[Enthalpy]

[...] Obviously I could use three of each to calculate the enthalpy of, say, combustion of benzene [...]

No, you couldn't. You should re-read and re-think what "aromatic" means. Not three double and three single bonds.


You should also understand that after breaking one bond, the other bonds are affected as well, a lot. That's even more true with aromatic and conjugated bonds. My tables don't even give bond energies for that case, because it wouldn't be very useful, but for sure very misleading. Useful to re-think too.

And since you mention it: bond energies shouldn't serve to compute reaction energies, because they're too inaccurate and because this method would subtract big quantities. Use the heat of formation of the molecules instead: it starts from the elements in their normal state, rather than separated atoms.

Tables of bond dissociation energies there - you could for instance observe C / CH / CH₂ / CH₃ / CH₄ and think at it:
www.nist.gov/data/nsrds/NSRDS-NBS31.pdf
http://staff.ustc.edu.cn/~luo971/2010-91-CRC-BDEs-Tables.pdf

[thetada]

Thank you. I understand that a benzene ring doesn't consist of alternating single and double bonds, that's why I've asked for what I've asked for.

I take your point that it's a sub-optimal way to calculate an enthalpy change, but it uses a technique my students are expected to master for their exams. Also the molecule in question, di(2-ethylhexyl) phthalate is too complex to have readily available enthalpy of formation data. I might switch the problem on its head and have them estimate the enthalpy of formation using the enthalpy of combustion value, which is available.

Thanks for your input.

[Enthalpy]

[...] estimate the enthalpy of formation using the enthalpy of combustion value, which is available. [...]

Which is the standard method from which enthalpies of formation are obtained and published.

Using instead the heat of hydrogenolysis should give more precise figures because hydrogenolysis releases less heat than a combustion, but I suspect it would be impractical, or at least less general than a combustion.

When a molecule like di(2-ethylhexyl) phthalate has no tabulated enthalpy of formation, the most accurate way among the many I've tried is to compare with other molecules which are tabulated. If (I didn't check) you have enthalpies of formation for dibutyl phthalate, butane and 2-ethylhexane, all in the same state of aggregation, then you can make formal transformations and you get reasonable estimates.

Though, that last way is less systematic hence more difficult to teach. It also fails on cage molecules - no method exists for them.

The biggest single difficulty, whatever the method, is to guess whether a published enthalpy of formation is measured or false.