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Topic: Estimating the heat of hydrogenation  (Read 3615 times)

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Offline curiouscat

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Estimating the heat of hydrogenation
« on: December 06, 2016, 11:43:45 AM »
I am trying to estimate the Heat of Hydrogenation for this reaction:



I'm getting conflicting values. A patent says 45 kcal/mol while an article gives 85 kcal/mol (with the caveat that it's a simulated value).

In another source I find the heuristic that "Aromatic hydrogenation reactions are highly exothermic. with heats of reaction typically in the range 63-71 kJ/mole of H2"  i.e. 16 kcal/mol

Any idea what seems the reasonable value?








Offline phth

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Re: Estimating the heat of hydrogenation
« Reply #1 on: December 06, 2016, 12:33:04 PM »
Those calculations are bulls&$#, and should be taken with a grain of salt unless they have been verified to be experimentally correct.  They could have done things such as use gas phase assumptions for a liquid phase reaction....What metal is participating in the reaction?  It seems like it would be easier just to measure the actual value using a bomb calorimeter that can operate at whatever pressure you want...http://pubs.acs.org/doi/pdf/10.1021/ed048p548

Offline curiouscat

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Re: Estimating the heat of hydrogenation
« Reply #2 on: December 06, 2016, 12:46:05 PM »
Those calculations are bulls&$#, and should be taken with a grain of salt unless they have been verified to be experimentally correct.  They could have done things such as use gas phase assumptions for a liquid phase reaction....What metal is participating in the reaction?  It seems like it would be easier just to measure the actual value using a bomb calorimeter that can operate at whatever pressure you want...http://pubs.acs.org/doi/pdf/10.1021/ed048p548

Sure. And at some stage I will do an experimental verification.

But for an initial approximate estimate I am reading up the literature.

Pd is the metal. But how does that affect the overall exothermicity?

Offline Corribus

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Re: Estimating the heat of hydrogenation
« Reply #3 on: December 06, 2016, 02:23:20 PM »
Well, you can certainly ballpark it pretty easily. This reaction entails forming an O-H and C-H bond and breaking a C-O and H-H bond, and releasing any angle strain energy involved in the epoxide. Although bond energies vary, I think it's ok to assume that taking average values will give us an estimate within about 10% of the real value. O-H bonds are about 110 kcal/mol and the C-H of a benzylic hydrogen (C6H5C-H) is about 113 kcal/mol, slightly more energetic than is typical for alkane C-H bonds (around 100 kcal/mol). A quick check around the internet puts angle strain values for ethylene oxide and 3-member alkane rings at 13 and 27.5 kcal/mole, respectively. So, summing up all these values gives a back of the envelope estimate of reaction enthalpy to be 188 kcal mole absorbed due to bond breaking, 223 kcal/mol released due to bond formation, and  an additional 13-28 kcal/mole released due to bond strain. Allowing for extra wiggle room, I'd put a rough estimate of reaction enthalpy to be around 45-65 kcal/mol released, depending on what the energy content of the strained epoxide ring is. Seems to be in line with the patent and article you mentioned first.

Of course, a calorimetric measurement or even a good computer calculation will be a better guide.
What men are poets who can speak of Jupiter if he were like a man, but if he is an immense spinning sphere of methane and ammonia must be silent?  - Richard P. Feynman

Offline curiouscat

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Re: Estimating the heat of hydrogenation
« Reply #4 on: December 06, 2016, 10:04:42 PM »
Thanks Corribus!

Just a nitpick:

Is the C-H of a benzylic slightly more, or *less* energetic than a typical alkane C-H?

I ask because Wikipedia gives a value of 90 kcal/mol

C6H5CH2–H   Benzylic C–H bond   90




Offline phth

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Re: Estimating the heat of hydrogenation
« Reply #5 on: December 07, 2016, 12:59:22 AM »
I would not be so sure about using wikipedia's values because they could be homolytic (radical) enthalpies.  I would find a peer-reviewed source, or read the citation, of bond enthalpies to know for sure.

Offline curiouscat

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Re: Estimating the heat of hydrogenation
« Reply #6 on: December 07, 2016, 02:44:04 AM »
I would not be so sure about using wikipedia's values because they could be homolytic (radical) enthalpies.  I would find a peer-reviewed source, or read the citation, of bond enthalpies to know for sure.

Another source: https://labs.chem.ucsb.edu/zakarian/armen/11---bonddissociationenergy.pdf

This gives 85.16 kcal/mol for C6H5CH2--H

Source: T. L. Cottrell, The Strengths of Chemical Bonds, 2d ed., Butterworth, London, 1958; B. deB. Darwent, National
Standard Reference Data Series, NationalBureau of Standards, no. 31, Washington, 1970; S. W. Benson, J. Chem. Educ.
42:502 (1965); and J. A. Kerr, Chem. Rev. 66:465 (1966).

Offline Corribus

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Re: Estimating the heat of hydrogenation
« Reply #7 on: December 07, 2016, 10:54:21 AM »
@Curiouscat

Sorry, I totally misread the table. You are correct re: approximate benzyllic hydrogen bond dissociation enthalpy.
What men are poets who can speak of Jupiter if he were like a man, but if he is an immense spinning sphere of methane and ammonia must be silent?  - Richard P. Feynman

Offline Enthalpy

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Re: Estimating the heat of hydrogenation
« Reply #8 on: December 07, 2016, 11:33:52 AM »
Those calculations are bulls&$# [...]
100% agreed. They usually try to estimate the gas at zero kelvin and zero pascal, and fail even on that.

Bond energies and cycle strain have some sense in them, but they cumulate errors up to being imprecise, and on cage molecules they can't work. Also, cyclopropyl makes much stronger bonds than alkyls do, and here the interaction with a phenyl is hard to estimate.
http://staff.ustc.edu.cn/~luo971/2010-91-CRC-BDEs-Tables.pdf

The best approach would be to compare the heats of formation. If you access Nist's professional tables (I don't) and they have the data for phenylethyl alcohol and phenyloxirane, bingo
http://webbook.nist.gov/cgi/inchi?ID=C60128&Mask=2#Thermo-Condensed
http://webbook.nist.gov/cgi/inchi/InChI%3D1S/C8H8O/c1-2-4-7(5-3-1)8-6-9-8/h1-5%2C8H%2C6H2

If not, use similar molecules, for instance propylene oxide and propanol, both liquid at +25°C:
http://webbook.nist.gov/cgi/cbook.cgi?ID=C75569&Mask=2#Thermo-Condensed
(note the wide dispersion of measured values, alas)
http://webbook.nist.gov/cgi/inchi?ID=C71238&Mask=2#Thermo-Condensed
-122.6kJ/mol propylene oxide
-302.5kJ/mol propanol
Though, this neglects the cyclopropyl-phenyl interaction.

These last heats of formation would suggest 180kJ/mol=43kcal/mol hydrogenation heat.

As you plan to measure the heat of hydrogenation, please tell the Nist, in case their professional tables know only one ΔHf, so they can deduce the other one. Note carefully if liquid, gaseous, what temperature, estimated precision...

Offline Babcock_Hall

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Re: Estimating the heat of hydrogenation
« Reply #9 on: December 07, 2016, 07:19:44 PM »
From McMurry's 8th edition Organic Chemistry I found 375 kJ/mol (89.6 kcal/mol) for the bond dissociation energy for the benzylic C-H bond.  I am unsure what is meant by hemolytic (radical) enthalpies.  Are these the same as, or different from, bond dissociation energies?

Offline Enthalpy

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Re: Estimating the heat of hydrogenation
« Reply #10 on: December 09, 2016, 06:28:37 PM »
Bond dissociation energy from Yu Ran Luo:
H−CH2C6H5 375.5 ± 5.0 kJ/mol

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