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CEM 850, Fall 2004

CEM 850, Fall 2004. Some notes on Thermochemistry, Bond Strengths, and Strain energies Ned Jackson. Alkane ∆H f values (kcal/mol). Alkane ∆H f values show Systematic Patterns. Can we estimate ∆H f by summing energy equivalents for transferable molecular “building blocks?” Bond Equivalents

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CEM 850, Fall 2004

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  1. CEM 850, Fall 2004 Some notes on Thermochemistry, Bond Strengths, and Strain energies Ned Jackson

  2. Alkane ∆Hf values (kcal/mol)

  3. Alkane ∆Hf values show Systematic Patterns • Can we estimate ∆Hf by summing energy equivalents for transferable molecular “building blocks?” • Bond Equivalents • Group Equivalents • Fragment transferability in comparisons between compounds implies deep similarity

  4. Bond Equivalents • Estimate ∆Hf values from C-H, C-C bonds: • Ethane ∆Hf = -20.04; 6 C-H, 1 C-C • Propane ∆Hf = -25.02; 8 C-H, 2 C-C • \C-H = -3.765; C-C = +2.55 • Predict ∆Hf(C5H12) = 4C-C + 12C-H = -34.98 • N-pentane -35.08 but isopentane = -40.14 • \Bond equivalents fail for branching.

  5. Group Equivalents • All alkanes can be expressed in terms of four building blocks: (CH3)i(CH2)j(CH)k(C)l (nonspecified bonds implicitly to C) • Enthalpy Equivalents: • CH3 -10.08 • CH2 -4.95 • CH -1.90 • C +0.50

  6. Group Equivalents (cont’d) • Analogous equivalents for alkenes and aromatics can be similarly derived, with value for CH3 held at -10.08 kcal/mol no matter what it’s attached to. • This method defines a “strainless ideal” for hydrocarbons of arbitrary formula, and allows the definition of “strain.”

  7. Strain Energies • Cycloalkanes (CH2)n

  8. Thermochemistry--why care? • Besides simple reaction ∆H and ∆G values, detailed energetics define reaction direction • Combined with bond strengths and kinetics of the reactions of interest, even imperfect energetic ideas put limits on mechanistic possibilities • Lead in to tools for comparing reactions!

  9. Bond Strengths • An X-Y bond, as defined by its atoms X and Y, is not a uniform (thus transferable) molecule building block • The bond equivalent approach did not lead to a reliable method for ∆Hf estimation • A group equivalent approach was required • Some bond strengths allow development of group equivalent ideas for reactions

  10. R-H BDEs worth remembering • H-H 104.2 kcal/mol • CH3-H 105.1 • CH3CH2-H 100.5 • (CH3)2CH-H 99.1 • (CH3)3C-H 95.2 • H2C=CHCH2-H 88.1 • PhCH2-H 89.6

  11. An ordinary C-C s bond • Generic C-C bond strengths in R-R’ • Use group equivalents to estimate ∆Hf values for R-R’, R-H, and R’-H • Get ∆Hf of R•, R’• radicals from R-H, R’-H via C-H BDEs + BDE(H2) = 104.2 kcal/mol • Calculate R-R’ BDE

  12. Cracking of Butane: 1-2 vs 2-3 • ∆Hf(butane) = 2(-10.08 -4.95) = -30.06 est. • ∆Hf(methane) = -17.9 • ∆Hf(ethane) = 2(-10.08) = -20.16 est. • ∆Hf(propane = 2(-10.08) -4.95 = -25.11 est. • ∆Hf(Me•) = -17.9 +105.1 -52.1 = 35.1 est. • ∆Hf(Et•) = -20.16 +100.5 -52.1 = 28.2 est. • ∆Hf(Pr•) = -25.11 +100.5 -52.1 = 23.3 est.

  13. Some Heats of Formation

  14. All energies in kcal/mol

  15. Bond Diss’n Energies (BDEs)

  16. The strength of a π bond • Breaking ethylene’s π bond doesn’t lead to two well-defined fragments. How can we define a separate “bond strength” for it? • Cis-trans isomerization of HDC=CHD? • Hydrogenation energies? • Spectroscopic measurements? • Others (full disassembly of molecule)?

  17. Ethylene isomerization • Heat cis or trans DHC=CHD and measure the rate of isomerization as a function of T. • From kinetic analysis, obtain ∆Hact for c-t isomerization: ~66 kcal/mol. • Problems: at high enough T, lots of other chemistry can happen; some may catalyze isomerization, making barrier appear too low. Or, isomerization might not go via rotation!? How to get a check on this value?

  18. Hydrogenation Strategy • H2C=CH2 + H2 —> H-H2C-CH2-H12.5 + 0 —> -20.0; ∆Hrxn = -32.5 kcal/mol • Broken: C-C π bond, H-H@104.2 kcal/mol; Formed: Two ethane C-H bonds @100.5 kcal/mol each • BDE(π) = 201. -32.5 -104.2 = 64.3 kcal/mol !Looks good!

  19. Spectroscopic approaches? • π—>π* Excited state has no π bonding, but lmax = 171 nm = ~167 kcal/mol!? Pretty far from 66! • ∆IE (ethylene - ethyl) (Electron’s energy-drop from non- to π-bonding = 10.51 - 8.12 eV= 55 kcal/mol per e–=> 110 kcal/mol!?

  20. Energetics of Full Disassembly of Ethylene • Try to make a prediction: • C-C s BDE is ~90 kcal/mol • the π bond is ~65 kcal/mol • Predict ~155 kcal/mol ∆H for C2H4 —> 2CH2 • ∆Hf(ethylene) = 12.5; ∆Hf(CH2) = 92.3; 184.6-12.5 = 172.1, almost 20 kcal/mol “too large”--what’s going on? • C-H bond strengths increase from C2H4 and CH2

  21. Cyclopropane Stereomutation • How strong is a C-C bond in cyclopropane? • Look at isomerization via isotopic labeling • Directly analogous to ethylene cis-trans isomerization • Should go via “real” open-chain biradical•H2C-CH2-CH2• • What about hydrogenation energies? • Can Strain E’s help?

  22. Thermal Stereomutation • Measured ∆Hact for c-t isomerization: • 63.7 kcal/mol (1958); 59.8 kcal/mol (1972) • ∆Hf of cyclopropane = 12.7 kcal/mol • \biradical ∆Hf should be ca. 72.5 kcal/mol • Primary C-H BDE back then was thought to be ca. 97 kcal/mol, instead of 100.5 • Propane = -25 + 2(97-52) = 65…huh?

  23. The propanediyl disaster • Thermochem looked like biradical must rest in a 5-9 kcal/mol well between c,t-isomers

  24. Why don’t we expect a barrier • General radical dimerization barrierless • Conceptual reason: there’s no stabilization to lose as bond formation begins. • “Hammond postulate” and/or Bell-Evans-Polanyi principle--the more exothermic the process, the lower its barrier will be.

  25. Review with current values • We calculated the 2-3 cleavage barrier for butane; cyclopropane should have the same number, lowered by its strain energy, which is released upon ring opening. • So 87.2 -27.5 kcal/mol directly predicts a barrier of 59.7, near the 1972 ∆Hact value. • Just need to revise primary C-H BDE up by 3.5 kcal/mol (x2 = ~the 7.5 kcal/mol error)

  26. The Methane Activation Problem • Methane combustion is very exothermic • CH4 + 2O2 --> CO2 + 2H2O • ∆Hcomb = -17.9 + 0 --> -94.1 + 2(-57.8) = -191.8 kcal/mol (plenty exothermic) • It’s a great fuel, but…it isn’t liquid • BP(CH4) = -162 ˚C = 111 K • \ Not practical for automotive use • (similar issues surround H2)

  27. Partial oxidation to liquify CH4? • Oxidation to methanol would be exothermic • CH4 + 1/2O2 --> CH3OH • ∆H = -17.9 + 0 --> -48.0 = -30.1 kcal/mol • Energy from CH3OH combustion? • CH3OH + O2 --> CO2 + 2H2O • ∆Hcomb = -48.0 --> -209.7 = -161.7 kcal/mol

  28. Hydrocarbon vs. Methanol Fuels:Energy Densities • Typical hydrocarbon “(CH2)n” • Mass = 14 g/mol • ∆Hcomb= -5 --> -94.1+(-57.8) = -146.9 kcal/mol • = 10.5 kcal/mol•gram • Methanol CH3OH • Mass = 30 g/mol • ∆Hcomb = -161.7 kcal/mol • = 5.4 kcal/mol•gram

  29. Challenge: CH4 --> CH3OH

  30. Protection of Methanol? • The C-H bond strengths in methanol are increased from 98.1 to 110 kcal/mol by methanol protonation, becoming stronger than those in methane (105.1 kcal/mol). Is this enough to control selectivity? • The key is the attacking species, presumably either HO• or CH3O• radicals here.

  31. Radical selectivities • Isobutane halogenation • Bond strengths matter!

  32. Radical Reactions: Selectivity vs. Exothermicity • The 103.2 kcal/mol H-Cl bond means that H-abstraction from any simple alkyl R-H is exothermic. • H-Br bond strength is just 87.5 kcal/mol so all H abstractions are endothermic. The relative barriers differ by nearly the whole energy difference between primary and tertiary radicals.

  33. How to obtain reaction barrier i.e. ∆Hact values? • Kinetics…for a later discussion • Measure reaction rates as a function of T • Extract rate constants for various T values • Arrhenius or Eyring plots to obtain ∆Eactand/or ∆Hact + ∆Sact

  34. Reaction Mechanisms • How many particles (intra- vs. intermolecular)? • Activation energies • What parts end up where? • Symmetries of TSs/Intermediates • What bonding changes happen, and when? • Concerted or stepwise? • Ionic or radical? • Catalyzed or direct? • Energy inputs (∆, hn, others?)?

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