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Bond-dissociation energy

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The bond-dissociation energy (BDE, D0, or DH°) is one measure of the strength of a chemical bond A−B. It can be defined as the standard enthalpy change when A−B is cleaved by homolysis to give fragments A and B, which are usually radical species.[1][2] The enthalpy change is temperature-dependent, and the bond-dissociation energy is often defined to be the enthalpy change of the homolysis at 0 K (absolute zero), although the enthalpy change at 298 K (standard conditions) is also a frequently encountered parameter.[3]

As a typical example, the bond-dissociation energy for one of the C−H bonds in ethane (C2H6) is defined as the standard enthalpy change of the process

CH3CH2−H → CH3CH2• + H•,
DH°298(CH3CH2−H) = ΔH° = 101.1(4) kcal/mol = 423.0 ± 1.7 kJ/mol = 4.40(2) eV (per bond).[4]

To convert a molar BDE to the energy needed to dissociate the bond per molecule, the conversion factor 23.060 kcal/mol (96.485 kJ/mol) for each eV can be used.

A variety of experimental techniques, including spectrometric determination of energy levels, generation of radicals by pyrolysis or photolysis, measurements of chemical kinetics and equilibrium, and various calorimetric and electrochemical methods have been used to measure bond dissociation energy values. Nevertheless, bond dissociation energy measurements are challenging and are subject to considerable error. The majority of currently known values are accurate to within ±1 or 2 kcal/mol (4–10 kJ/mol).[5] Moreover, values measured in the past, especially before the 1970s, can be especially unreliable and have been subject to revisions on the order of 10 kcal/mol (e.g., benzene C–H bonds, from 103 kcal/mol in 1965 to the modern accepted value of 112.9(5) kcal/mol). Even in modern times (between 1990 and 2004), the O−H bond of phenol has been reported to be anywhere from 85.8 to 91.0 kcal/mol.[6] On the other hand, the bond dissociation energy of H2 at 298 K has been measured to high precision and accuracy: DH°298(H−H) = 104.1539(1) kcal/mol or 435.780 kJ/mol.[5]

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The term bond-dissociation energy is similar to the related notion of bond-dissociation enthalpy (or bond enthalpy), which is sometimes used interchangeably. However, some authors make the distinction that the bond-dissociation energy (D0) refers to the enthalpy change at 0 K, while the term bond-dissociation enthalpy is used for the enthalpy change at 298 K (unambiguously denoted DH°298). The former parameter tends to be favored in theoretical and computational work, while the latter is more convenient for thermochemical studies. For typical chemical systems, the numerical difference between the quantities is small, and the distinction can often be ignored. For a hydrocarbon RH, where R is significantly larger than H, for instance, the relationship D0(R−H) ≈ DH°298(R−H) − 1.5 kcal/mol is a good approximation.[7] Some textbooks ignore the temperature dependence,[8] while others have defined the bond-dissociation energy to be the reaction enthalpy of homolysis at 298 K.[9][10][11]

The bond dissociation energy is related to but slightly different from the depth of the associated potential energy well of the bond, De, known as the electronic energy. This is due to the existence of a zero-point energy ε0 for the vibrational ground state, which reduces the amount of energy needed to reach the dissociation limit. Thus, D0 is slightly less than De, and the relationship D0 = De − ε0 holds.[7]

The bond dissociation energy is an enthalpy change of a particular chemical process, namely homolytic bond cleavage, and "bond strength" as measured by the BDE should not be regarded as an intrinsic property of a particular bond type but rather as an energy change that depends on the chemical context. For instance, Blanksby and Ellison cites the example of ketene (H2C=CO), which has a C=C bond dissociation energy of 79 kcal/mol, while ethylene (H2C=CH2) has a bond dissociation energy of 174 kcal/mol. This vast difference is accounted for by the thermodynamic stability of carbon monoxide (CO), formed upon the C=C bond cleavage of ketene.[7] The difference in availability of spin states upon fragmentation further complicates the use of BDE as a measure of bond strength for head-to-head comparisons, and force constants have been suggested as an alternative.[12]

Historically, the vast majority of tabulated bond energy values are bond enthalpies. More recently, however, the free energy analogue of bond-dissociation enthalpy, known as the bond-dissociation free energy (BDFE), has become more prevalent in the chemical literature. The BDFE of a bond A–B can be defined in the same way as the BDE as the standard free energy change (ΔG°) accompanying homolytic dissociation of AB into A and B. However, it is often thought of and computed stepwise as the sum of the free-energy changes of heterolytic bond dissociation (A–B → A+ + :B−), followed by one-electron reduction of A (A+ + e− → A•) and one-electron oxidation of B (:B− → •B + e−).[13] In contrast to the BDE, which is usually defined and measured in the gas phase, the BDFE is often determined in the solution phase with respect to a solvent like DMSO, since the free-energy changes for the aforementioned thermochemical steps can be determined from parameters like acid dissociation constants (pKa) and standard redox potentials (ε°) that are measured in solution.[14]

Bond energy

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Except for diatomic molecules, the bond-dissociation energy differs from the bond energy. While the bond-dissociation energy is the energy of a single chemical bond, the bond energy is the average of all the bond-dissociation energies of the bonds of the same type for a given molecule.[15] For a homoleptic compound EXn, the E–X bond energy is (1/n) multiplied by the enthalpy change of the reaction EXn → E + nX. Average bond energies given in tables are the average values of the bond energies of a collection of species containing "typical" examples of the bond in question.

For example, dissociation of HO−H bond of a water molecule (H2O) requires 118.8 kcal/mol (497.1 kJ/mol). The dissociation of the remaining hydroxyl radical requires 101.8 kcal/mol (425.9 kJ/mol). The bond energy of the covalent O−H bonds in water is said to be 110.3 kcal/mol (461.5 kJ/mol), the average of these values.[16]

In the same way, for removing successive hydrogen atoms from methane the bond-dissociation energies are 105 kcal/mol (439 kJ/mol) for D(CH3−H), 110 kcal/mol (460 kJ/mol) for D(CH2−H), 101 kcal/mol (423 kJ/mol) for D(CH−H) and finally 81 kcal/mol (339 kJ/mol) for D(C−H). The bond energy is, thus, 99 kcal/mol, or 414 kJ/mol (the average of the bond-dissociation energies). None of the individual bond-dissociation energies equals the bond energy of 99 kcal/mol.[17][7]

Strongest bonds and weakest bonds

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According to BDE data, the strongest single bonds are Si−F bonds. The BDE for H3Si−F is 152 kcal/mol, almost 50% stronger than the H3C−F bond (110 kcal/mol). The BDE for F3Si−F is even larger, at 166 kcal/mol. One consequence to these data are that many reactions generate silicon fluorides, such as glass etching, deprotection in organic synthesis, and volcanic emissions.[18] The strength of the bond is attributed to the substantial electronegativity difference between silicon and fluorine, which leads to a substantial contribution from both ionic and covalent bonding to the overall strength of the bond.[19] The C−C single bond of diacetylene (HC≡C−C≡CH) linking two sp-hybridized carbon atoms is also among the strongest, at 160 kcal/mol.[5] The strongest bond for a neutral compound, including multiple bonds, is found in carbon monoxide at 257 kcal/mol. The protonated forms of CO, HCN and N2 are said to have even stronger bonds, although another study argues that the use of BDE as a measure of bond strength in these cases is misleading.[12]

On the other end of the scale, there is no clear boundary between a very weak covalent bond and an intermolecular interaction. Lewis acid–base complexes between transition metal fragments and noble gases are among the weakest of bonds with substantial covalent character, with (CO)5W:Ar having a W–Ar bond dissociation energy of less than 3.0 kcal/mol.[20] Held together entirely by the van der Waals force, helium dimer, He2, has the lowest measured bond dissociation energy of only 0.022 kcal/mol.[21][22]

Homolytic versus heterolytic dissociation

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Bonds can be broken symmetrically or asymmetrically. The former is called homolysis and is the basis of the usual BDEs. Asymmetric scission of a bond is called heterolysis. For molecular hydrogen, the alternatives are:

Symmetric: H2 → 2 H• ΔH° = 104.2 kcal/mol (see table below)
Asymmetric: H2 → H+ + H− ΔH° = 400.4 kcal/mol (gas phase)[23]
Asymmetric: H2 → H+ + H− ΔG° = 34.2 kcal/mol (in water)[24] (pKaaq = 25.1)

In the gas phase, the enthalpy of heterolysis is larger than that of homolysis, due to the need to separate unlike charges. However, this value is lowered substantially in the presence of a solvent.

Representative bond enthalpies

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The data tabulated below shows how bond strengths vary over the periodic table.

Bond Bond Bond-dissociation enthalpy at 298 K Comment
(kcal/mol) (kJ/mol) (eV/bond)
C−C in typical alkane 83–90 347–377 3.60–3.90 Strong, but weaker than C−H bonds
C−F in CH3F 115 481 4.99 Very strong, rationalizes inertness of Teflon
H−H hydrogen 103 431 4.52 Strong, nonpolarizable bond
H−F hydrogen fluoride 136 569 5.90 Very strong
O−H in water 119 497 5.15 Very strong, hydroxyl radical reactive with almost all organics exothermically by H atom abstraction
O−H in methanol 105 440 4.56 Slightly stronger than C−H bonds
O−H in α-tocopherol (an antioxidant) 77 323 3.35 O−H bond strength depends strongly on substituent on O
C-O methanol 92 385 3.99 typical alcohol
C≡O carbon monoxide 257 1077 11.16 Strongest bond in neutral molecule
O=CO carbon dioxide 127 532 5.51 Slightly stronger than C−H bonds, surprisingly low due to stability of C≡O
O=CH2 formaldehyde 179 748 7.75 Much stronger than C−H bonds
O=O oxygen 119 498 5.15 Stronger than single bonds, weaker than many other double bonds
N≡N nitrogen 226 945 9.79 One of the strongest bonds, large activation energy in production of ammonia

There is great interest, especially in organic chemistry, concerning relative strengths of bonds within a given group of compounds, and representative bond dissociation energies for common organic compounds are shown below.[7][17]

Bond Bond Bond-dissociation energy at 298 K Comment
(kcal/mol) (kJ/mol) (eV/bond)
H3C−H Methyl C−H bond 105 439 4.550 One of the strongest aliphatic C−H bonds
C2H5−H Ethyl C−H bond 101 423 4.384 Slightly weaker than H3C−H
(CH3)2CH−H Isopropyl C−H bond 99 414 4.293 Secondary radicals are stabilized
(CH3)3C−H t-Butyl C−H bond 96.5 404 4.187 Tertiary radicals are even more stabilized
(CH3)2NCH2−H C−H bond α to amine 91 381 3.949 Lone-pair bearing heteroatoms weaken C−H bonds
(CH2)3OCH−H C−H bond α to ether 92 385 3.990 Lone-pair bearing heteroatoms weaken C−H bonds. THF tends to form hydroperoxides
CH3C(=O)CH2−H C−H bond α to ketone 96 402 4.163 Conjugating electron-withdrawing groups weaken C−H bonds
CH2CH−H Vinyl C−H bond 111 464 4.809 Vinyl radicals are uncommon
HCC−H Acetylenic C−H bond 133 556 5.763 Acetylenic radicals are very rare
C6H5−H Phenyl C−H bond 113 473 4.902 Comparable to vinyl radical, uncommon
CH2CHCH2−H Allylic C−H bond 89 372 3.856 Such bonds show enhanced reactivity, see drying oil
C6H5CH2−H Benzylic C−H bond 90 377 3.907 Akin to allylic C−H bonds. Such bonds show enhanced reactivity
H3C−CH3 Alkane C−C bond 83–90 347–377 3.60–3.90 Much weaker than C−H bond. Homolytic cleavage occurs when H3C−CH3 thermolysed at >500 °C
H2C=CH2 Alkene C=C bond ~170 ~710 ~7.4 About 2 times stronger than a C−C single bond; however, the π bond (~65 kcal/mol) is weaker than the σ bond
HC≡CH Alkyne C≡C triple bond ~230 ~960 ~10.0 About 2.5 times stronger than a C−C single bond

See also

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References

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  1. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "Bond-dissociation energy". doi:10.1351/goldbook.B00699
  2. ^ The value reported as the bond-dissociation energy (BDE) is generally the enthalpy of the homolytic dissociation of a gas-phase species. For instance, the BDE of diiodine is calculated as twice the heat of formation of iodine radical (25.5 kcal/mol) minus the heat of formation of diiodine gas (14.9 kcal/mol). This gives the accepted BDE of diiodine of 36.1 kcal/mol. (By definition, diiodine in the solid state has a heat of formation of 0.)
  3. ^ The IUPAC Gold Book does not stipulate a temperature for its definition of bond-dissociation energy (ref. 1).
  4. ^ The corresponding BDE at 0 K (D0) is 99.5(5) kcal/mol.
  5. ^ a b c Luo, Y. R. (2007). Comprehensive handbook of chemical bond energies. Boca Raton: CRC Press. ISBN 978-0-8493-7366-4. OCLC 76961295.
  6. ^ Mulder P, Korth HG, Pratt DA, DiLabio GA, Valgimigli L, Pedulli GF, Ingold KU (March 2005). "Critical re-evaluation of the O−H bond dissociation enthalpy in phenol". The Journal of Physical Chemistry A. 109 (11): 2647–55. Bibcode:2005JPCA..109.2647M. doi:10.1021/jp047148f. PMID 16833571.
  7. ^ a b c d e Blanksby SJ, Ellison GB (April 2003). "Bond dissociation energies of organic molecules". Accounts of Chemical Research. 36 (4): 255–63. CiteSeerX 10.1.1.616.3043. doi:10.1021/ar020230d. PMID 12693923.
  8. ^ Anslyn, Eric V.; Dougherty, Dennis A. (2006). Modern physical organic chemistry. Sausalito, CA: University Science. ISBN 978-1-891389-31-3. OCLC 55600610.
  9. ^ Darwent, B. deB. (January 1970). Bond Dissociation Energies in Simple Molecules (PDF). NSRDS-NBS 31. Washington, DC: U.S. National Bureau of Standards. LCCN 70602101.
  10. ^ Streitwieser, Andrew; Heathcock, Clayton H.; Kosower, Edward M. (2017). Introduction to Organic Chemistry. New Delhi: Medtech (Scientific International, reprint of 4th revised edition, 1998, Macmillan). p. 101. ISBN 978-93-85998-89-8.
  11. ^ Carroll, Felix A. (2010). Perspectives on structure and mechanism in organic chemistry (2nd ed.). Hoboken, N.J.: John Wiley. ISBN 978-0-470-27610-5. OCLC 286483846.
  12. ^ a b Kalescky, Robert; Kraka, Elfi; Cremer, Dieter (2013-08-30). "Identification of the Strongest Bonds in Chemistry". The Journal of Physical Chemistry A. 117 (36): 8981–8995. Bibcode:2013JPCA..117.8981K. doi:10.1021/jp406200w. ISSN 1089-5639. PMID 23927609. S2CID 11884042.
  13. ^ Miller DC, Tarantino KT, Knowles RR (June 2016). "Proton-Coupled Electron Transfer in Organic Synthesis: Fundamentals, Applications, and Opportunities". Topics in Current Chemistry. 374 (3): 30. doi:10.1007/s41061-016-0030-6. PMC 5107260. PMID 27573270.
  14. ^ Bordwell, F. G.; Cheng, Jin Pei; Harrelson, John A. (February 1988). "Homolytic bond dissociation energies in solution from equilibrium acidity and electrochemical data". Journal of the American Chemical Society. 110 (4): 1229–1231. doi:10.1021/ja00212a035.
  15. ^ Norman, Richard O. C.; Coxon, James M. (2001). Principles of organic synthesis (3rd ed.). London: Nelson Thornes. p. 7. ISBN 978-0-7487-6162-3. OCLC 48595804.
  16. ^ Lehninger, Albert L.; Nelson, David L.; Cox, Michael M. (2005). Lehninger Principles of Biochemistry (4th ed.). W. H. Freeman. p. 48. ISBN 978-0-7167-4339-2. Retrieved May 20, 2016.
  17. ^ a b Streitwieser A.; Bergman R. G. (19 September 2018). "Table of Bond Dissociation Energies". University of California, Berkeley. Retrieved 13 March 2019.
  18. ^ Lide, David R., ed. (2006). CRC Handbook of Chemistry and Physics (87th ed.). Boca Raton, Florida: CRC Press. ISBN 0-8493-0487-3.
  19. ^ Gillespie, Ronald J. (July 1998). "Covalent and Ionic Molecules: Why Are BeF2 and AlF3 High Melting Point Solids whereas BF3 and SiF4 Are Gases?". Journal of Chemical Education. 75 (7): 923. Bibcode:1998JChEd..75..923G. doi:10.1021/ed075p923. ISSN 0021-9584.
  20. ^ Grills D. C.; George M. W. (2001), "Transition metal-noble gas complexes", Advances in Inorganic Chemistry, Elsevier, pp. 113–150, doi:10.1016/s0898-8838(05)52002-6, ISBN 978-0-12-023652-7.
  21. ^ Cerpa, Erick; Krapp, Andreas; Flores-Moreno, Roberto; Donald, Kelling J.; Merino, Gabriel (2009-02-09). "Influence of Endohedral Confinement on the Electronic Interaction between He atoms: A He2@C20H20 Case Study". Chemistry – A European Journal. 15 (8): 1985–1990. doi:10.1002/chem.200801399. ISSN 0947-6539. PMID 19021178.
  22. ^ Giese, Timothy J.; York, Darrin M. (2004). "High‐level ab initio methods for calculation of potential energy surfaces of van der Waals complexes". International Journal of Quantum Chemistry. 98 (4): 388–408. doi:10.1002/qua.20074. ISSN 0020-7608.
  23. ^ Bartmess, John E.; Scott, Judith A.; McIver, Robert T. (September 1979). "Scale of acidities in the gas phase from methanol to phenol". Journal of the American Chemical Society. 101 (20): 6046–6056. doi:10.1021/ja00514a030.
  24. ^ Connelly, Samantha J.; Wiedner, Eric S.; Appel, Aaron M. (2015-03-17). "Predicting the reactivity of hydride donors in water: thermodynamic constants for hydrogen". Dalton Transactions. 44 (13): 5933–5938. doi:10.1039/C4DT03841J. ISSN 1477-9234. PMID 25697077.
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