Patai S., Rappoport Z. 1996 The chemistry of functional groups. The chemistry of amino, nitroso, nitro and related groups / 0379 422

NHCH2 C(CH3 )3

40.0

The first acidity measurement of aniline was reported by McEwen122. He assigned an ion-pair pKa of 27 to aniline in benzene in relation to MeOH in benzene, which was assumed to have an ion-pair pKa of 16. Acidities of a number of substituted anilines were measured by means of the H method in H2O DMSO mixture123, and by NMR124, calorimetric125 and electrochemical126 methods in liquid ammonia at 31 °C.

The equilibrium acidities of aniline and 26 of its derivatives have recently been measured127 in DMSO solution and they cover a range of 15 pKa units. The pKa values of anilines in DMSO are 10 units higher than the ion-pair pKa values obtained in liquid ammonia. This effect has been attributed to the greater capacity of ammonia to solvate the proton. When the pKa values determined in DMSO were compared with those derived in H2O DMSO or EtOH DMSO, it was evident that the latter were appreciably lower. This behaviour is different from that found for the carbon acids. The H2O DMSO mixture can solvate nitranions better than pure DMSO, in line with the fact that nitranions are better HBA than carbanions128 (see Table 10).

A plot of the pKa values of anilines in DMSO versus the pKa values of anilinium ions in water is linear with a slope of 1.8. This allows the extrapolation of 41 š 1 for the pKa of ammonia in DMSO from that of NH4C in water (9.27)127. Alternatively, the pKa of ammonia in DMSO has been extrapolated105 from the intersystem correlation between the DMSO acidities of NH2X and PhNHX as a value of 35.8. Extrapolation of the pKa value of ammonia from the Taft-like dual substituent parameter (DSP) of NH2X DMSO acidities gave a similar value of 36.6.

C. NH Bond Dissociation Energy

Bordwell129 developed a method of estimating relative bond dissociation energies (BDE) for families of acids, HA, by combining equilibrium acidity constants, pKHA, with the oxidation potential of their conjugated bases, A , both measured in DMSO:

a Measured in DMSO.

bMeasured by cyclic voltammetry in DMSO. c Calculated.

dLiterature data. eReference 134. fReference 135. gReference 136.

with C D 55.4 55.9 kcal mol 1 (a value of 56 has also been used, and more recently one of 73131,132). The BDEs estimated in this way agree remarkably well with the values obtained in the gas phase (š2 kcal) (Table 11), although the relationship must be considered empirical133. The cleavage modes of radical anions have also been considered131:

In the case of anilines, path A will be favoured on the basis of the values of oxidation potentials, in analogy with the gas-phase results.

An extensive study132 has been presented on the polarizability effects of alkyl groups in RX moieties (R D Me, Et, i-Pr and t-Bu; X D CH2, S, SO2, O and N) in families of weak acids and on the stabilities of adjacent anions and radicals in DMSO solution. Some of the results related to the 9-(dialkylamino)fluorenes are given in Table 12. The increases in acidity are believed to be caused by the progressive increases in anion stabilizing

TABLE 12. Acidities and homolytic bond dissociation energies of the acidic C H bonds in 9-(dialkylamino)fluorenes (9-R1R2NFlH)

a Equilibrium acidities measured in DMSO: Reference 139.

b Oxidation potentials of the conjugated anions measured in DMSO solution and referenced to the ferrocene/ferrocenium couple.

polarizability effects of the R group that parallel increases in alkyl size. Polarizability effects stabilize anions but not the analogous radicals.

Taking advantage137 of the outlined method and considering that RSE (radical stabilization energy) is equivalent to the variation in BDE, Bordwell estimated the RSEs of radicals of the type A Cž H D, where A is an electron acceptor and D an electron donor. Comparison of the oxidation potential of ˛-H2N and ˛-R2N carbanions derived from fluorene, acetophenones and malononitrile indicates that N-alkylation has little or no effect on the stability of the corresponding ˛-amino radicals (no captodative effect). Nevertheless, when saturation effects are taken into account there is evidence of a synergistic (captodative) effect. This indicates that the presence of strong donor and acceptor groups attached to a radical centre stabilizes the radicals with synergistic effects, in accord with the previous literature data138.

D. Carboxamides and Related Compounds

A comparison of acidities of six series of analogous oxygen, nitrogen and carbon acids in DMSO solution and in the gas phase has shown that the nitrogen acids are more acidic than their carbon acid counterparts by 17 š 5 kcal mol 1 (kcal mol 1 D 1.37 pKa unit), and that the oxygen acids are more acidic than the nitrogen acids by a similar amount. Differently, carboxamides have been found only slightly more acidic than their carbon acid analogues, the ketones (1 2 kcal mol 1 in DMSO and 7 8 kcal mol 1 in the gas phase140). This behaviour has been rationalized in terms of destabilization of carboxamide conjugate bases by lone pair lone pair repulsions, together with some resonance stabilization of the carboxamide.

The replacement of the CDO bond in carboxamides by the CDS bond, as in thio-

weakening of the N H bond by 14 17 kcal mol 1, evaluated by using relation 26. This effect has been considered a consequence of large increases in ground state energies and the superior ability of sulphur than oxygen in stabilizing a negative charge or odd electron.

In the gas phase, the acidity of H C bond increases as the s character of carbon increases and the homolytic bond dissociation energy decreases as shown by considering the following compounds: H3CCH3, H2CDCH2 and HC CH with Gacid kcal mol 1 D 413, 401 and 370 and BDE kcal mol 1 D 98, 110 and 132, respectively. A similar trend with the changes in hybridization of nitrogen was not demonstrated, due to the paucity of available data. The problem could be more complicated in the case of substrates containing more acidifying centres as acetamidine, for which the following scheme can be considered:

Although an analogous scheme can be written for acetamide, recent ab initio calculations142 have questioned the importance of resonance in amides, i.e. 17 $ 170 , indicating that the nitrogen in 17 has a slight negative charge rather than the positive

charge imposed by the resonance contributor 170.

The equilibrium acidities in DMSO have been measured for acetamidine, benzamidine, N,N0 -diphenylbenzamidine, N,N-diethylbenzamidine, diphenylmethanimine, guanidine, N,N0-diphenylguanidine, N,N0-diphenylurea, and N,N0-diphenylthiourea143. Combination of the resulting pKa values with the oxidation potentials of their conjugated bases gave estimates of their BDEs. These values have been compared with those of the corresponding carboxamides and thiocarboxamides. The changes in hybridization of nitrogen between NH3 and Ph2CDNH produce a changes in acidities and BDEs similar to the increases observed for the increase in the s character of carbon. Nevertheless, the BDE

11.8 kcal mol 1 (Table 13). Replacement of the methyl group by an amino group causes a decrease in acidity. The effects of substituting one phenyl group at the acidic sites produce an increase in acidity, not different from the increase obtained by introducing a phenyl group at each of the nitrogen atoms. This is due to the strong steric interaction between the two phenyl groups that allow only one phenyl group to overlap with the negative charge or odd electron.

The equilibrium acidities of aldoximes, amidoximes and ketoximes144 have also been measured in DMSO solution. It has been found that syn- and anti-benzaldoximes have nearly the same acidities, contrary to the finding in aqueous solution where the anti isomers are less acidic. The lower acidity of the anti isomers in aqueous solution is attributable to steric inhibition of solvation of the strongly hydrogen-bonded oxide ion by the phenyl group, a factor that is absent in DMSO. The 8.3 pKa unit difference between the acidity of acetaldoxime (28.5) and benzaldoxime (syn- 20.2, anti- 20.3) in DMSO points to appreciable delocalization of the negative charge in the benzaldoximide ion into the benzene ring.

Ionization equilibria of sulphamide, phenylsulphamide, mono-, diand tri-substituted sulphamides have been considered and pKa values have been determined145 for

TABLE 13. Acidities of amidines and related compounds in DMSO

R D H, X-aryl, alkyl

It has been found that in the case of equilibrium 30 the corresponding aliphatic series exhibit decreased pKa values by about 5 pKa units146.

Under the same experimental conditions147 the following equilibria have also been considered and pKa values have been determined to be in the range 9 12.

X D Me, H, Br, Cl, OMe, OEt, NO2

R1 D H, Ac, Me

R2 D cyclo-C6H11, Bu, CH2C6H5

When cyclic sulphamides (18 and 19) were studied148, it was found that the sixmembered rings are more acidic than their acyclic analogues by ca 2.5 pKa units and that the five-membered cyclic sulphamides are ca 4 pKa units more acidic than the model open-chain counterparts.

VI. HYDROXAMIC ACIDS AS AMPHOTERIC SYSTEMS

Notwithstanding their name, hydroxamic acids (20) are amphoteric systems with behaviours that are both basic and acidic. The chemistry of hydroxamic acids and N- hydroxyimides was reviewed by Bauer and Exner149 in 1974. Since then, the basic properties of hydroxamic acid derivatives have received little attention. Given the similarity of the activity coefficient behaviour of hydroxamic acids and other carbonyl bases (amides), it was originally assumed that the carbonyl oxygen rather than the nitrogen atom was the site of protonation150,151, but some recent IR studies have revealed signs of N-protonation152. This question has been recently addressed by Bagno and colleagues153, who concluded from NMR measurements made in aqueous solution that acetohydroxamic acid undergoes carbonyl protonation as the dominant process. The protonation equilibrium of this compound was determined in aqueous sulphuric acid and the use of the excess acidity method led to mŁ D 0.25 and pKBH D 1.15. The gas-phase basicity of some

hydroxamic acid derivatives (Table 14) has been determined154, the values indicating that they are 7 kcal mol 1 weaker bases than amides. The protonation site cannot be determined from gas-phase results, but a pH-metric study of the dissociation constants of L-˛-alaninehydroxamic and ˇ-alaninehydroxamic acids has shown that the NH3C group is more acidic than the NHOH group for the ˛-derivative, while the reverse is true for the ˇ-isomer155.

H

TABLE 14. Gas-phase basicities of some hydroxamic acid derivatives (kcal mol 1)

There has been a longstanding controversy as to whether hydroxamic acids are NH or OH acids. The IR and UV evidence presented by Exner and coworkers149 indicated that they are NH acids in dioxane and aqueous alcohol solvents, and this conclusion has also been supported by a 17O NMR study156. However, other studies157 on meta- and para-substituted benzohydroxamic acids led to the conclusion that, in aqueous solution, RCONHO ions were at least as present as RCON(OH) ions. More recently, Crumbliss and coworkers158 have concluded that, in an aqueous 2M NaNO3 solution, acetoand benzohydroxamic acids act as OH rather than NH acids. Bordwell’s group159 measured the equilibrium acidities in DMSO of acetoand benzohydroxamic acids, as well as their N- and O-alkyl derivatives, and concluded that hydroxamic acids act as NH acids in non-HBD solvents, such as DMSO, DMF, CH3CN etc., but they may act primarily as OH acids in hydroxylic solvents.

The gas-phase acidities of hydroxamic acid and its N- and O-methyl derivatives have been measured using FT ion cyclotron resonance (Table 15)160. The acidity order is the same as that found in DMSO although, in this solvent, the O-methylation of acetohydroxamic acid decreases the acidity by 1 pK unit and the N-methylation by 3.6 pK units, respectively.

VII. HETEROCYCLES

Imidazole is a five-membered heterocycle with both basic and acidic properties. As part of histidine, it plays an important role in proton transfer processes at the active site of certain enzymes, and is thus considered a good model for studying enzymatic modes of action by means of computational approaches. Much work has been done to characterize this molecule and its derivatives. The proton affinity161 of imidazole has been determined experimentally by Taft and coworkers, who also determined the gas-phase basicities of pyrazole, all possible isomeric methylpyrazoles (13 derivatives), and a selected set of methylimidazoles (8 derivatives), using STO-3G fully-optimized geometries of neutral molecules and their corresponding cations. The protonation energies of imidazole, calculated subsequently by Meyer162 using the HF/6-311GŁŁ basis set with additional polarization functions on hydrogen atoms163, are lower than those previously calculated without polarization functions. The substitution of hydrogen with a methyl group causes a small increase in the deprotonation/protonation energies of 1 kcal mol 1. The overall picture indicates that imidazole and 4-methyl imidazole seem to be interchangeable in model studies of the protonation of histidine.

Knowing the equilibrium constants for the protonation of the amino and imidazole nitrogens, and the kinetics of deprotonation of the C(2) carbon of imidazole, it was possible to investigate nitrogen-protonation microequilibria and the C(2)-deprotonation microkinetics of histidine, histamine and other related compounds164.