March's Advanced Organic Chemistry: Reactions, Mechanisms, and Structure, 7th Edition (2013)

Chapter 8. Acids and Bases

8.F. The Effects of Structure on the Strengths of Acids and Bases¹⁷⁵

The structure of a molecule can affect its acidity or basicity in a number of ways. Unfortunately, in most molecules two or more of these effects (as well as solvent effects) are operating, and it is usually very difficult or impossible to say how much each effect contributes to the acid or base strength.¹⁷⁶ Small differences in acidity or basicity between similar molecules are particularly difficult to interpret. It is well to be cautious when attributing them to any particular effect.

1. Field Effects. These effects were discussed in Section 1.I. In general, changes in substituents can have an effect on acidity. As an example of the influence of field effects on acidity, compare the acidity of acetic acid and 2-nitroacetic acid:

The only difference in the structure of these molecules is the substitution of NO₂ for H. Since NO₂ is a strongly electron-withdrawing group, it withdraws electron density from the negatively charged COO⁻ group in the anion of 2-nitroacetic acid (compared with the anion of acetic acid). As the pK_a values indicate, 2-nitroacetic acid is ~1000 times stronger than acetic acid.¹⁷⁷ Any effect that results in electron withdrawal from a negatively charged center (−I effect) is a stabilizing effect because it spreads the charge. Thus, −I groups increase the acidity of uncharged acids (e.g., acetic) because they spread the negative charge of the anion. However, −I groups also increase the acidity of any acid, no matter what the charge. For example, if the acid has a charge of +1 (and its conjugate base is therefore uncharged), a −I group destabilizes the positive center (by increasing and concentrating the positive charge) of the acid, a destabilization that will be relieved when the proton is lost. In general, groups that withdraw electrons by the field effect increase acidity and decrease basicity, while electron-donating groups act in the opposite direction. Another example is the molecule (C₆F₅)₃CH, which has three strongly electron-withdrawing C₆F₅ groups and a pK_a of 16,¹⁷⁸ compared with Ph₃CH, with a pK_a of 31.5 (Table 8.1), an acidity enhancement of ~ 10¹⁵. Table 8.5 shows pK_a values for some acids. An approximate idea of field effects can be obtained from this table. In the case of the chlorobutyric acids, the effect decreases with distance. It must be remembered, however, that field effects are not the sole cause of the acidity differences noted and that in fact solvation effects may be more important in many cases (see Sec. 8.G).¹⁷⁹ The influence of various substituents on the acidity of acetic acid has been calculated,¹⁸⁰ Substituent effects for weak acids (e.g., phenols and benzyl alcohols) have been discussed.¹⁸¹

Field effects are important in benzoic acid derivatives, and the pK_a of the acid will vary with the nature and placement of the “X” group in 6.¹⁸² The pK_a of 3-OMe (6) is 5.55, but 4-OMe (6) is 6.02 in 50% aq methanol,¹⁸³compared with a pK_a of 5.67 when X = H. When X = 4-NO₂, the pK_a is 4.76 and 4-Br is 5.36.¹⁷² The pK_a of 2,6-diphenylbenzoic acid is 6.39.¹⁸⁴

2. Resonance Effects. Resonance that stabilizes a base, but not its conjugate acid, results in the acid having a higher acidity than otherwise expected and vice versa. An example is found in the higher acidity of carboxylic acids¹⁸⁵ compared with primary alcohols.

The RCOO⁻ ion is stabilized by resonance not available to the RCH₂O⁻ ion (or to RCOOH).¹⁸⁶ Note that the RCOO⁻ is stabilized not only by the fact that there are 2 equiv canonical forms, but also by the fact that the negative charge is spread over both oxygen atoms, and is therefore less concentrated than in RCH₂O⁻. The same effect is found in other compounds containing a C=O or CN group. Thus amides (RCONH₂) are more acidic than amines (RCH₂NH₂); esters (RCH₂COOR′) more than ethers (RCH₂CH₂OR′); and ketones (RCH₂COR′) more than alkanes (RCH₂CH₂R′) (Table 8.1). The effect is enhanced when two carbonyl groups are attached to the same carbon (because of additional resonance and spreading of charge); for example, β-keto esters (see 7) are more acidic than simple ketones or carboxylic esters (Table 8.1). Compounds such as (7) are generically referred to as active methylene compounds (X–CH₂–X), where X is an electron-withdrawing groups (a carbonyl, cyano, sulfonyl, etc.).¹⁸⁷ The influence of substituents in the α-position of substituted ethyl acetate derivatives has been studied.¹⁸⁸Extreme examples of this effect are found in the molecules tricyanomethane [(NC)₃CH], with a pK_a of −5 (Table 8.1), and 2-(dicyanomethylene)-1,1,3,3-tetracyanopropene (NC)₂C=C[CH(CN)₂]₂, whose first pK_a is below −8.5 and whose second pK_a is −2.5.

Resonance effects are also important in aromatic amines. m-Nitroaniline is a weaker base than aniline, a fact that can be accounted for by the −I effect of the nitro group. But p-nitroaniline is weaker still, though the −I effect should be less because of the greater distance. This result is obtained by taking the canonical form A into account. Because A contributes to the resonance hybrid,¹⁸⁹ the electron density of the unshared pair is lower in p-nitroaniline than in m-nitroaniline, where a canonical form (e.g, A) is impossible. Note that the pK_a values reported are those of the conjugate acid, the ammonium ion.¹⁹⁰ The basicity is lower in the para compound for two reasons, both caused by the same effect: (1) the unshared pair is less available for attack by a proton, and (2) when the conjugate acid is formed, the resonance stabilization afforded by A is no longer available because the previously unshared pair is now being shared by the proton. The acidity of phenols is affected by substituents in a similar manner.¹⁹¹

In general, resonance effects lead to the same result as field effects. That is, here too, electron-withdrawing groups increase acidity and decrease basicity, and electron-donating groups act in the opposite manner. As a result of both resonance and field effects, charge dispersal leads to greater stability.

3. Periodic Table Correlations. When comparing Brnsted acids and bases that differ in the position of an element in the periodic table:

a. Acidity increases and basicity decreases in going from left to right across a row of the periodic table. Thus acidity increases in the order CH₄ < NH₃ < H₂O < HF, and basicity decreases in the order ⁻CH₃ > ⁻NH₂ > ⁻OH > F⁻. This behavior can be explained by the increase in electronegativity upon going from left to right across the table. It is this effect that is responsible for the great differences in acidity between carboxylic acids, amides, and ketones: RCOOH RCONH₂ RCOCH₃.

b. Acidity increases and basicity decreases in going down a column of the periodic table, despite the decrease in electronegativity. Thus acidity increases in the order HF < HCl < HBr < HI and H₂O < H₂S, and basicity decreases in the order NH₃ > PH₃ > AsH₃. This behavior is related to the size of the species involved. Thus, for example, F⁻, which is much smaller than I⁻, attracts a proton much more readily because its negative charge occupies a smaller volume, and is therefore more concentrated (note that F⁻ is also much harder than I⁻ and is thus more attracted to the hard proton; see Sec. 8.E). This rule does not always hold for positively charged acids. Thus, although the order of acidity for the group 16 hydrides is H₂O < H₂S < H₂Se, the acidity order for the positively charged ions is H₃O⁺ > H₃S⁺ > H₃Se⁺.¹⁹²

Lewis acidity is also affected by periodic table considerations. In comparing acid strengths of Lewis acids of the form MX_n¹⁶¹:

c. Acids that require only one electron pair to complete an outer shell are stronger than those that require two. Thus GaCl₃ is stronger than ZnCl₂. This results from the relatively smaller energy gain in adding an electron pair that does not complete an outer shell and from the buildup of negative charge if two pairs come in.

d. Other things being equal, the acidity of MX_n decreases in going down the periodic table because as the size of the molecule increases, the attraction between the positive nucleus and the incoming electron pair is weaker. Thus BCl₃ is a stronger acid than AlCl₃.¹⁹³

4. Statistical Effects. In a symmetrical diprotic acid, the first dissociation constant is twice as large as expected since there are 2 equiv ionizable hydrogens, while the second constant is only one-half as large as expected because the conjugate base can accept a proton at 2 equiv sites. So K₁/K₂ should be 4, and approximately this value is found for dicarboxylic acids where the two groups are sufficiently far apart in the molecule that they do not influence each other. A similar argument holds for molecules with 2 equiv basic groups.¹⁹⁴

5. Hydrogen Bonding. Internal hydrogen bonding can greatly influence acid or base strength. For example, the pK for o-hydroxybenzoic acid is 2.98, while the value for the para isomer is 4.58. Internal hydrogen bonding between the ⁻OH and COO⁻ groups of the conjugate base of the ortho isomer stabilizes it and results in an increased acidity.

6. Steric Effects. The proton itself is so small that direct steric hindrance is seldom encountered in proton transfers. Steric effects are much more common in Lewis acid–base reactions in which larger acids are used. Spectacular changes in the order of base strength have been demonstrated when the size of the acid was changed. Table 8.6 shows the order of base strength of simple amines when compared against acids of various size.¹⁹⁵ It can be seen that the usual order of basicity of amines (when the proton is the reference acid) can be completely inverted by using a large enough acid. The strain caused by formation of a covalent bond when the two atoms involved each have three large groups is called face strain or F strain.

Steric effects can indirectly affect acidity or basicity by affecting the resonance (Sec. 2.F). For example, o-tert-butylbenzoic acid is ~ 10 times as strong as the para isomer, because the carboxyl group is forced out of the plane by the tert-butyl group. Indeed, virtually all ortho benzoic acids are stronger than the corresponding para isomers, regardless of whether the group on the ring is electron donating or electron withdrawing.

Steric effects can also be caused by other types of strain. 1,8-Bis(diethylamino)-2,7-dimethoxynaphthalene (8) is an extremely strong base for a tertiary amine (pK_a of the conjugate acid = 16.3; cf. N,N-dimethylaniline, pK_a = 5.1), but proton transfers to and from the nitrogen are exceptionally slow; slow enough to be followed by a UV spectrophotometer.¹⁹⁶ Compound 8 is severely strained because the two nitrogen lone pairs are forced to be near each other.¹⁹⁷ Protonation relieves the strain: one lone pair is now connected to a hydrogen, which forms a hydrogen

bond to the other lone pair (shown in 9). The same effects are found in 4,5-bis(dimethylamino)fluorene (10)¹⁹⁸ and 4,5-bis(dimethylamino)phenanthrene (11).¹⁹⁹ Compounds (e.g., 8, 10, and 11), are known as proton sponges.²⁰⁰The basicity of a proton sponge has been calculated as the sum of the proton affinity¹⁵² of an appropriate reference monoamine, the strain released on protonation, and the energy of the intramolecular hydrogen bond formed on protonation.²⁰¹ Another type of proton sponge is quino[7,8-h]quinoline (12).²⁰² Protonation of this compound also gives a stable mono protonated ion similar to 9, but the steric hindrance found in 8, 10, and 11 is absent. Therefore, 12 is a much stronger base than quinoline (13) (pK_a values of the conjugate acids are 12.8 for 12 and 4.9 for 13), but proton transfers are not abnormally slow. A cyclam-like macrocyclic tetramine (15) was prepared by a coupling reaction of bispidine, and was shown to be a new class of proton sponge.²⁰³

Chiral Lewis acids are known. Indeed, an air stable and storable chiral Lewis acid catalyst has been prepared, a chiral zirconium catalyst combined with molecular sieves powder.²⁰⁴ Association of a bulky silicon group with the bis(trifluoromethanesulfonyl)imide (known as triflimide) anion leads to enhancement of the electrophilic character of R₃SiNTf₂. The presence of a chiral substituent derived from (−)-myrtenal on the silicon atom led to a chiral silicon Lewis acid.²⁰⁵

Another type of steric effect is the result of an entropy effect. The compound 2,6-di-tert-butylpyridine is a weaker base than either pyridine or 2,6-dimethylpyridine.²⁰⁶ The reason is that the conjugate acid 14 is less stable than the conjugate acids of nonsterically hindered pyridines. In all cases, the conjugate acids are hydrogen bonded to a water molecule, but in the case of 14 the bulky tert-butyl groups restrict rotations in the water molecule, lowering the entropy.²⁰⁷

The conformation of a molecule can also affect its acidity. The following pK_a values were determined for compounds 16–19.²⁰⁸

Since ketones are stronger acids than carboxylic esters (Table 8.1), it is not surprising that 16 is a stronger acid than 18.²⁰⁹ A comparison of 16 with cyclic diketone 17 shows an increase in acidity of only 2.1 pK units, while a comparison of 18 with cyclic diester 19 shows an increase of 8.6 units. Indeed, 19 (called Meldrum's acid) is an unusually strong acid for a 1,3-diester. In order to account for this very large effect of a ring, MO calculations were carried out for two conformations of methyl acetate and of its enolate ion.²¹⁰ Loss of a proton is easier by ~5 kcal mol⁻¹ (21 kJ mol⁻¹) for the syn than for the anti conformer of the ester. In an acyclic molecule like 18, the preferred conformations are anti, but in Meldrum's acid (19) the conformation on both sides is constrained to be syn.

Facial differences in proton reactivity can lead to enantioselective deprotonation. Enantioselective deprotonation is also achieved by using a chiral base and/or a chiral complexing agent. Enantioselective deprotonation in cyclic ketones,²¹¹ and with heterodimer bases has been studied.²¹² When a Lewis acid coordinates to a base, the resulting complex can have conformational properties that influence reactivity. Coordination of SnCl₄ with aldehydes and esters, for example, leads to a complex where the conformation is determined by interactions of the C=O SnCl₄ unit with substituents attached to the carbonyl.²¹³

7. Hybridization. An s orbital has a lower energy than a p orbital. Therefore, the more s character a hybrid orbital contains, the lower the energy of that orbital. It follows that a carbanion at an sp carbon is more stable than a corresponding carbanion at an sp² carbon. Thus HCC⁻, which has more s character in its unshared pair than CH₂=CH⁻ or CH₃CH₂⁻ (sp vs sp² vs sp³, respectively), is a much weaker base. This explains the relatively high acidity of acetylenes and HCN. Another example is that alcohol and ether oxygen atoms, where the unshared pair is sp³, are more strongly basic than carbonyl oxygen atoms, where the unshared pair is sp² (Table 8.1).

An understanding of the reactivity of bases arises from the study of their structures in solution and in the crystalline state. Due to the importance of dialkylamide bases, there is a significant body of work, led by independent work by Williard and by Collum, that has attempted to understand the structures of these reactive molecules. It is clear that the dialkylamide bases are aggregates. Note that the simplest member of the amide base family, lithium amide (LiNH₂), was shown to be monomeric and unsolvated, as determined using a combination of gas-phase synthesis and millimeter/submillimeter-wave spectroscopy.²¹⁴ Both monomeric LiNH₂ and LiNMe₂ are planar.²¹⁵Lithium diisopropylamide (LiNiPr₂, LDA) was isolated from a THF solution and X-ray crystallography revealed a dimeric structure (20; R = iPr, S = THF) in the solid state.²¹⁶ Lithium diisopropylamide was also shown to be a dimer in solutions of THF²¹⁷ and/or HMPA (see 20, R = iPr and S = THF, HMPA).²¹⁸ In the presence of HMPA, many derivatives of 20 tend to be mixed aggregates.²¹⁹ Extremely hindered LiNR₂ (R = 2-adamantyl) are monomeric under all conditions.²²⁰ In hydrocarbon solvents, lithium tetramethylpiperidide [LTMP, RR′NLi, where RR′ = –CMe₂(CH₂)₃C(Me₂)–] forms cyclic trimers and tetramers, with the tetrameric species predominating.²²¹In THF, lithium hexamethyldisilazide [LHMDS, (Me₃Si)₂NLi] forms a five-coordinate tetrasolvate [(Me₃Si)₂NLi(thf)₄],²²² but in ether there is an equilibrium mixture of monomer and dimer.²²³ A review is available that discusses the solution structures of amide bases LiNR₂.²²⁴ Chiral lithium amide bases are known and they show similar behavior in solution.²²⁵ Chelation effects are common in enantioenriched amide bases, which also form aggregates.²²⁶ The aggregation state of lithium phenylacetonitrile has been studied.²²⁷ Dianion aggregates can be generated, and in the case of the lithiation reaction of N-silyl allylamine, X-ray structure determination showed the presence of three uniquely different aggregates.²²⁸ A mixed aggregate is formed when the lithium enolate of a ketone is mixed with a lithium amide.²²⁹

Similar information is available for other bases. Lithium phenoxide (LiOPh) is a tetramer in THF.²³⁰ Lithium 3,5-dimethylphenoxide is a tetramer in ether, but addition of HMPA leads to dissociation to a monomer.²³¹

Enolate anions are nucleophiles in reactions with alkyl halides (Reaction 10-68), with aldehydes and ketones (Reactions 16-34 and 16-36) and with acid derivatives (Reaction 16-85). Enolate anions are also bases, reacting with water, alcohols and other protic solvents, and even the carbonyl precursor to the enolate anion. Enolate anions exist as aggregates, and the effect of solvent on aggregation and reactivity of lithium enolate anions has been studied.²³² Alkyl substitution has a significant influence on the energetics of enolate anions.²³³

Table 8.4 Some Absolute Hardness Values in Electron Volts^a

Table 8.5 The pK Values for Some Acids^a

a. See Ref. 47.

Table 8.6 Bases Listed in Increasing Order of Base Strength when Compared with Certain Reference Acids.