Structural Effects on Acidity

STRUCTURAL EFFECTS ON ACIDITY

Now that a method is in hand to compare acid strengths quantitatively and predict the position of acid–base equilibria, a look at Table 3.1 reveals that organic compounds have an enormous range of acidities from very strong acids such as arenesulfonic acids (pK_a = −6.5) and protonated carbonyl compounds (pK_a = −7 to pK_a = −10) to very weak acids such as alkanes (pK_a = 50) and alkenes (pK_a = 45). This huge range of acidity of ∼ 10⁶⁰ is reflective of the huge diversity of structural elements present in organic compounds.

How the structure of an acid influences its pK_a provides a quantitative way to compare the structure of a compound with its reactivity (in this case acidity). Such structure–reactivity correlations are crucial for our understanding of how reactions take place and for being able to predict how a structural change will affect the outcome of a reaction. The ability to predict how a reaction will respond to changes in structure (or other variables) takes us out of the realm of trial and error and into the realm of rational approaches to chemical transformations. Let us examine briefly some of the structural features which are major influences on acidity.

Returning to the dissociation of an acid in water, it is seen that the process has some energy costs and some energy gains. It is this energy balance ( ) which determines the equilibrium constant and hence the strength of an acid. The dissociation of an acid in water has an energy cost from the breaking of a bond to hydrogen and the separation of charges produced by the ionization.

However, there is an energy gain from the formation of the OH bond of the hydronium ion and the solvation of the anion by hydrogen bonding to the solvent and solvation of the hydronium by its hydrogen bonding to the solvent. If a series of different acids is now compared, it becomes clear that a major energetic difference in the dissociation of various acids in water is the stability of the conjugate base and its interaction with the aqueous solvent.

This is because the other energetic factors which influence the equilibria are similar for different acids. Bonds from hydrogen to the first-row elements have similar bond strengths (±5 kcal/mol) so the energy cost of breaking the bond to hydrogen is relatively constant for most acids of first-row elements. This analysis is especially true for acids with the proton bonded to the same element. Moreover, since the solvent is always water, the energy required to separate charges is about the same. Finally, the H–O bond of the hydronium ion is the same, regardless of which acid supplies the proton. Consequently, the principal differences in the ΔG’s of ionization for various acids are due to the differences in stability of their conjugate bases in the reaction mixture.

This analysis suggests that structural features which stabilize the conjugate base (often an anion) will therefore increase the acidity of an acid. While there are exceptions to this general approach (e.g., comparison of the acidities of acids in the second and third rows of the periodic table), it provides a sound basis for predicting what structural factors can increase or decrease the acidity of organic acids.

There are three principal factors that lead to increased stability of anions: (a) the electronegativity of the atom carrying the negative charge, (b) inductive effects which can stabilize negative charge, and (c) resonance effects which delo-calize the negative charge over several atoms and hence stabilize the anion.

Electronegativity

Increased electronegativity of an atom allows it to carry negative charge more readily, and the stability of the anion is increased. It is for this reason that the order of acidity of first-row hydrides is C–H < NH < –OH < FH. Transfer of a proton from these substances to water yields a series of anions whose stabilities are ordered according to the electronegativity of the negatively charged atom.

Such ordering is valid only for elements in the same row in the periodic table. Comparisons between acids in which the proton is lost from elements from dif-ferent rows is not valid because the bond strength to the acidic hydrogen changes greatly from row to row. In the above analysis bond strength is assumed to be relatively constant; thus it cannot be neglected when significant bond strength differences between acids are present.

The effective electronegativity of the atom carrying the charge is also depen-dent on the hybridization of that atom. As the s character of an orbital increases, electrons in that orbital are more stable due to greater attraction to the nucleus. Thus the effective electronegativity of the atom increases. This effect is clearly seen in the relative acidities of hydrocarbons. Removal of a proton from alkanes, alkenes, and alkynes produce conjugate bases with electron pairs in sp³, sp², and sp orbitals, respectively. As the amount of s character increases from 25 to 33 to 50% in this series, the stability of the conjugate base increases and accounts for the marked increase in acidity in the series. Based on these data, it is expected that cyclopropane, which because of ring strain has the hydrogens bonded to carbons which are hybridized at about an sp^2.5 level (29% s character), should have a pK_a between that of an alkane and an alkene. In fact, the pK_a of cyclopropane is 46 as predicted.

The increase in acidity by 25 orders of magnitude between sp³- and sp-hybridized carbon acids is similar to that found for the difference in acid-ity between an ammonium ion (sp³ hybridization) and a protonated nitrile (sp hybridization). It is clear that the hybridization of the orbital they occupy can play a major role in stabilizing electron pairs and thus influencing the effective electronegativity of an atom.

Inductive Effects

The inductive effect is the ability of a substituent or group near the acidic proton to alter the electron distribution at the reaction center by through-bond displace-ment of electrons. The result is that substituents which withdraw electrons from the reaction center by the inductive effect stabilize anions and thus increase the acidity of the conjugate acids of those anions. Conversely, groups which donate electrons make the reaction center more electron rich and thus make the for-mation of the anion at that center more difficult. The conjugate acid is thus a weaker acid.

This is easily demonstrated by considering a group of substituted acetic acids (Table 3.2). Compared to acetic acid (X = H), replacement of a hydrogen by more electron-withdrawing groups [Cl, F, (CH₃)₃N⁺ ] leads to an increase in the acidity. Replacement of hydrogen with an electron-donating t -butyl group decreases the acidity. We can understand these changes in terms of the inductive effect. If we compare the conjugate bases of acetic acid and chloroacetic acids, it is seen that the carbon–chlorine bond has a dipole moment associated with it.

This bond dipole induces smaller dipole moments in adjacent bonds, which in turn induces ever smaller dipole moments in adjacent bonds.

The result of this inductive effect is that the electron density on the carboxylate anion is reduced, the negative charge is distributed over more atoms, and the chloroacetate anion is stabilized relative to acetate. Because the chloroacetate anion is more stable than the acetate ion, its conjugate acid, chloroacetic acid, is a stronger acid than the conjugate acid of the acetate ion, acetic acid (Table 3.1).

As is expected, groups with higher electronegativity (X = F) or electron defi-ciency result in greater inductive electron withdrawal, the anion is more stable, and the acidity is increased. Conversely, a group such as t -butyl is electron donating relative to hydrogen. Its inductive effect serves to increase the electron density on the carboxylate group, destabilize the anion, and thus decrease the acidity of its conjugate acid.

As mentioned, inductive effects operate through bonds by successive bond polarizations. As such, they diminish rapidly with distance so that very little effect results if an inductive effect must be transferred through more than four bonds. As seen in Table 3.3, placement of a chlorine substituent next to the carboxyl group causes a hundredfold increase in acidity. Moving it to the β position reduces the effect significantly, while a γ -chloro substituent causes almost no acidity increase.

Inductive effects serve to alter the electron distributions in molecules, and con-sequently they are very important influences on many types of reactions—not just acidity and basicity. To the extent that electronic changes occur during the con-version of reactants to products, inductive effects can facilitate or impede those electronic changes and thus change the rates of conversion. It is important then to keep them in mind when other examples of reactivity changes are discussed.

Resonance Effects

A final structural effect which influences acidity is the delocalization of electrons via resonance. In terms of acid–base behavior, resonance delocalization can sta-bilize the conjugate base of an acid, thus making the acid a stronger acid. For example, alcohols have pK_a’s of ∼16 whereas carboxylic acids have pK_a’s of ∼5. In each case the acidic proton is lost from oxygen. The bond strength to the proton and the electronegativity of the atom carrying the charge (oxygen) are identical; thus these factors cannot account for the large difference in acidity. On the other hand, the alkoxide ion is a localized anion with the oxygen atom carry-ing a full negative charge while the carboxylate ion is resonance delocalized. In the carboxylate ion the electron pair and negative charge are distributed between both oxygens so that each oxygen carries only a partial negative charge (actually about – 1/2 ) and the anion is greatly stabilized. Thus carboxylic acids are more acidic than alcohols by ≈10¹¹ or so.

The following groups of compounds illustrate the profound effect that reso-nance delocalization has on the stability of anions and hence the acidity of the conjugate acids. To compare the acidities of these acids, the conjugate bases can be ranked according to their resonance stabilization and that ranking of anion stabilization is predictive of the acidity orders.

While resonance stabilization is greatest for those compounds which have more electron density distributed to more electronegative elements (compare 1, 2, and 3), delocalization of charge over any elements results in significant anion stabilization and a corresponding increase in acidity of the conjugate acid of that anion (e.g., 6, 7, 8). Moreover electronegativity effects can be considered in addition to resonance effects where applicable. Both amides 5 and methyl ketones 8 have resonance stabilization, but in amide anions the negative charge is shared between nitrogen and oxygen, while in ketone enolates the negative charge is shared between carbon and oxygen. Due to the greater electronegativity of nitrogen over carbon, the amide anion is more stable and hence amides (pK_a ≈ 18–19) are somewhat more acidic than ketones (pK_a ≈ 20–21)

A particularly strong type of resonance stabilization is found for those com-pounds which form an aromatic ring upon removal of a proton. The enhanced aromatic stability of the conjugate base translates into a large increase in acidity of the acid. Whereas the doubly allylic proton of 1,4-pentadiene is predicted to have a pK_a ≈ 40 due to resonance stabilization of the anion, the doubly allylic proton in cyclopentadiene has a pK_a = 16 because the resulting anion produces an aromatic π system.

Aromaticity also explains why tropolone (pK_a ≈ −5) is slightly more basic than a normal ketone (pK_a ≈ −7). The conjugate acid is stabilized upon proto-nation by the formation of an aromatic tropylium ion.

Inductive and resonance effects described above can significantly alter the electron distributions in molecules and can influence not only acidity but many other reactions as well. A general understanding of these effects will be important in many different transformations we will encounter.