Logic vs Misconceptions in Undergraduate Organic Textbooks: Radical Stabilities and Bond Dissociation Energies

• Andreas A. Zavitsas ^[1]
  1. [1] Long Island University

     Long Island University

     Town of Oyster Bay, Estados Unidos

     
• Localización: Journal of chemical education, ISSN 0021-9584, Vol. 78, Nº 3, 2001, pág. 417
• Idioma: inglés
• Texto completo no disponible (Saber más ...)
• Resumen

  • Undergraduate textbooks for organic chemistry relate alkyl radical stability to homolytic bond dissociation energy of the corresponding alkane: BDE[H3C-H] = 105.0 > BDE[CH3CH2-H] = 100.5 > BDE[(CH3)2CH-H] = 99.1 > BDE[(CH3)3C-H] = 95.2 kcal/mol. This practice is misleading, grossly overestimates the stability of the radicals relative to methyl, and leads to values that are not transferable to other molecules. It is based on the flawed argument that, since H· is a common product, any differences in BDE[R-H] are due to the relative stabilities of R·. When the argument is applied to other classes of compounds (R-OH, R-Cl, R-OCH3, R-NH2, and R-F), the BDE order is quite different: the methyl compounds have the weakest bond, and methyl radical appears as the most stable. The focus is placed on the fact that BDE[A-B] is the difference between the starting state (A-B molecule) and the final state (A· and B·). While the stability of the radicals affects BDE, neglect of the effect of the bond dipole on the stability of the starting state leads to the paradox. Defining alkyl radical relative stabilization energies in terms of BDE[R-Me] and use of Pauling's electronegativity equation solves the problem.