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The introduction of substituents into organic compounds often leads to charge shifts and significant dipole moments. This may have important consequences for structure, energy, and reactivity. A quantity frequently used to describe the polar nature of substituents is the electronegativity, first defined by Pauling (1). The electronegativities of the first- and second-row elements are given in Table 1 (2, 3). They cover a wide range. Fluorine is the most electronegative and sodium and lithium are the least electronegative. Despite its importance, however, the electronegativity of substituents has not received much attention in many organic chemistry textbooks. The purpose of this review is to make organic chemists more aware of charge distributions in molecules and the associated electrostatic effects (4).
Bond Dissociation Energies
The strengths of chemical bonds are usually expressed as bond dissociation energies (BDEs). As a result of recent experimental work, relatively precise values of BDEs for many organic compounds are known (5), and some of them are summarized in Table 2. In the series methylamine, methanol, and methyl fluoride there is a marked increase in the CH3-X BDE, which parallels the increasing difference in electronegativity between the carbon and the substituent. This effect was noted by Pauling many years ago as part of a general trend of increasing strength of covalent bonds as their polarity increased (1). In the case of methyl fluoride, the electronegativity of fluorine leads to a marked polarization in the sense C+-F, and the increased strength of the bond presumably arises from coulombic attraction between the charges on carbon and fluorine. The polarization is much less for the C-N bond in methylamine because of the smaller electronegativity of nitrogen. This leads to the smaller BDE.
The decrease in BDE from ethane to methylamine probably arises from the change in hybridization. In ethane, the C-C bond is formed from two ~sp3 hybrids at carbon, whereas in methylamine, the orbital from nitrogen to carbon has much more p character so as to place the lone pair in an orbital with as much s character as possible. Bond strengths increase with increasing s character (6), leading to the C-C bond being stronger than C-N.
It is interesting that with the second-row substituents, the C-Si bond is stronger than C-CI and much stronger...