Inductive Electronic Effects of Substituents on Acid/Base Strength |
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Inductive Electronic Effects of Substituents on Acid/Base Strength
In the above sections we examined three factors that give rise to differences in relative acidities of X–H bonds within different functional groups. Most biological and drug molecules, however, have more than one functional group (substituents) and these may influence the electronic properties and acid/base properties of each other. The extent to which the properties are altered depends on the identity and placement of the substituents relative to one another, as well as on what types of bonds connect them. In this and the following two sections we examine the three main types of interactions between substituents: (1) inductive electronic effects, (2) combined resonance and inductive electronic effects, and (3) proximity effects.
As noted for resonance electronic effects, any structural feature that withdraws or pulls electron density away from a basic X atom decreases its basicity and increases the acidity of the corresponding X–H bond. Conversely, features that donate or push electron density toward a basic X atom increase its basicity and decrease the acidity of its corresponding X–H acid. While resonance effects result from direct delocalization of electron pairs into a conjugated π system, inductive effects result from induced polarization of σ bonds connecting the basic X atom of interest and a polarized substituent. The substituent is typically a functional group with a polarized bond, that is, a dipole moment. Inductive effects can be electron donating or electron withdrawing, depending on the direction of the dipole relative to the X–H bond (Figure 1.1).
Figure 1.1 Polarized bonds between atoms with different electronegativities may be inductively electron donating or electron withdrawing. In electron donating substituents the negative end of the dipole points toward the acidic X–H group, whereas in electron withdrawing substituents the positive end of the dipole points toward the acidic X–H group.
Recall that the directionality and magnitude of the dipole in a polarized bond is determined by the relative electronegativities of the bonded atoms. Of the atoms found most often in organic molecules, only H is less electronegative than C, while all of the common heteroatoms (X = N, O, S, F, Cl, Br, I) are more electronegative than C (Table 1.1). Thus, H–C bonds are polarized with slight greater electron density on carbon and thus are electron donating. In contrast, the greater electronegativity of the X heteroatoms makes X–C bonds polarized in the opposite direction and electron withdrawing.
Recall that the directionality and magnitude of the dipole in a polarized bond is determined by the relative electronegativities of the bonded atoms. Of the atoms found most often in organic molecules, only H is less electronegative than C, while all of the common heteroatoms (X = N, O, S, F, Cl, Br, I) are more electronegative than C (Table 1.1).
Table 1.1 Electronegativities of Elements in the First Three Rows of the Periodic Table.