CEM 850, Fall 2004

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CEM 850, Fall 2004

• Some notes on Thermochemistry, Bond Strengths,
  and Strain energies
• Ned Jackson

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Alkane ?Hf values (kcal/mol)
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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

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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.

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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

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Group Equivalents (contd)

• Analogous equivalents for alkenes and aromatics
  can be similarly derived, with value for CH3 held
  at -10.08 kcal/mol no matter what its attached
  to.
• This method defines a strainless ideal for
  hydrocarbons of arbitrary formula, and allows the
  definition of strain.

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Strain Energies

• Cycloalkanes (CH2)n

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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!

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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

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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
• H2CCHCH2-H 88.1
• PhCH2-H 89.6

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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

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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.

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Some Heats of Formation
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All energies in kcal/mol
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Bond Dissn Energies (BDEs)
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The strength of a p bond

• Breaking ethylenes p bond doesnt lead to two
  well-defined fragments. How can we define a
  separate bond strength for it?
• Cis-trans isomerization of HDCCHD?
• Hydrogenation energies?
• Spectroscopic measurements?
• Others (full disassembly of molecule)?

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Ethylene isomerization

• Heat cis or trans DHCCHD 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?

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Hydrogenation Strategy

• H2CCH2 H2 gt H-H2C-CH2-H12.5 0 gt
  -20.0 ?Hrxn -32.5 kcal/mol
• Broken C-C p bond, H-H_at_104.2 kcal/mol Formed
  Two ethane C-H bonds _at_100.5 kcal/mol each
• BDE(p) 201. -32.5 -104.2 64.3
  kcal/mol !Looks good!

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Spectroscopic approaches?

• pgtp Excited state has no p bonding, but lmax
  171 nm 167 kcal/mol!? Pretty far from 66!
• ?IE (ethylene - ethyl) (Electrons energy-drop
  from non- to p-bonding
• 10.51 - 8.12 eV 55 kcal/mol per e gt 110
  kcal/mol!?

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Energetics of Full Disassembly of Ethylene

• Try to make a prediction
• C-C s BDE is 90 kcal/mol
• the p bond is 65 kcal/mol
• Predict 155 kcal/mol ?H for C2H4 gt 2CH2
• ?Hf(ethylene) 12.5 ?Hf(CH2) 92.3 184.6-12.5
  172.1, almost 20 kcal/mol too large--whats
  going on?
• C-H bond strengths increase from C2H4 and CH2

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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 Es help?

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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) 65huh?

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The propanediyl disaster

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

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Why dont we expect a barrier

• General radical dimerization barrierless
• Conceptual reason theres 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.

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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)

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The Methane Activation Problem

• Methane combustion is very exothermic
• CH4 2O2 --gt CO2 2H2O
• ?Hcomb -17.9 0 --gt -94.1 2(-57.8) -191.8
  kcal/mol (plenty exothermic)
• Its a great fuel, butit isnt liquid
• BP(CH4) -162 C 111 K
• \ Not practical for automotive use
• (similar issues surround H2)

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Partial oxidation to liquify CH4?

• Oxidation to methanol would be exothermic
• CH4 1/2O2 --gt CH3OH
• ?H -17.9 0 --gt -48.0 -30.1 kcal/mol
• Energy from CH3OH combustion?
• CH3OH O2 --gt CO2 2H2O
• ?Hcomb -48.0 --gt -209.7 -161.7 kcal/mol

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Hydrocarbon vs. Methanol FuelsEnergy Densities

• Typical hydrocarbon (CH2)n
• Mass 14 g/mol
• ?Hcomb -5 --gt -94.1(-57.8) -146.9 kcal/mol
• 10.5 kcal/molgram
• Methanol CH3OH
• Mass 30 g/mol
• ?Hcomb -161.7 kcal/mol
• 5.4 kcal/molgram

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Challenge CH4 --gt CH3OH
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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.

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Radical selectivities

• Isobutane halogenation
• Bond strengths matter!

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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.

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How to obtain reaction barrier i.e. ?Hact values?

• Kineticsfor a later discussion
• Measure reaction rates as a function of T
• Extract rate constants for various T values
• Arrhenius or Eyring plots to obtain ?Eact and/or
  ?Hact ?Sact

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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?)?