J. Phys. Chem. 1996, 100, 17539-17543
17539
Redox and Acidity Properties of Alkyl- and Arylamine Radical Cations and the Corresponding Aminyl Radicals1 Mats Jonsson,*,2 Danial D. M. Wayner, and Janusz Lusztyk Steacie Institute for Molecular Sciences, National Research Council of Canada, Ottawa, Ontario, K1A 0R6 Canada ReceiVed: May 6, 1996; In Final Form: August 13, 1996X
In this work the pKas of six alkyl- and arylamine radical cations in aqueous solution have been determined by means of laser flash photolysis. The corresponding N-nitrosamines were used as precursors for the aminyl radicals which, upon protonation, formed the amine radical cations. The following pKas were obtained: 3.6 ( 0.2, 7.6 ( 0.3, 6.8 ( 0.5, 5.3 ( 0.5, 5.5 ( 0.5, and 5.8 ( 0.5 for the radical cations of diphenylamine, N-methylaniline, dimethylamine, diethylamine, pyrrolidine, and piperidine, respectively. In addition, the peak oxidation potentials of diphenylamine, N-methylaniline, aniline, diethylamine, pyrrolidine, and piperidine have been measured in aqueous solution and in acetonitrile using cyclic voltammetry. Furthermore, the oneelectron reduction potentials of the diphenylaminyl radical and the N-methylanilinyl radical in acetonitrile were measured using photomodulation voltammetry. The results of this study and of previously published studies are discussed in terms of relative substituent and solvent effects.
Introduction
SCHEME 1
Aminyl radicals and their protonated counterparts, amine radical cations, are important intermediates in organic synthesis3 as well as in biological systems, e.g., enzymatic oxidation of amines.4-8 Reactions between enzyme model compounds and a large number of amines have been studied extensively and it has been found that the primary step in the reaction between cytochrome P-450 and amines is electron transfer to give the amine radical cation.5-7 The reducing properties of amines, arylamines in particular, make them suitable as antioxidants and they are used as such in polymer materials; thus, oxidative degradation of polymers leads to formation of aminyl radicals or amine radical cations. Aminyl radicals are also formed upon photolytic and thermal decomposition of N-nitrosamines and N-haloamines.3 Studies by Chow et al.,3 Horner et al.,9 and Wagner et al.10 have given some insight into the reactivity of aminyl radicals and amine radical cations. Since these radicals are involved in such a variety of natural and synthetic processes it is interesting to elucidate the intrinsic thermochemical properties that govern their reactivity. The thermochemical properties that are of primary interest here are the one-electron reduction potential of the radicals (reaction 1a) and the radical cations (reaction 1b) and the pKa of the neutral amines (reaction 2a) and the radical cations (reaction 2b). The four reactions are related in a thermochemical cycle (Scheme 1). The reduction potentials and the pKas of the radical cations are the solution counterparts to the gas-phase ionization potentials of the amines and the proton affinities of the aminyl radicals. The gas-phase ionization potential, Igas, of the amine and the one-electron reduction potential, E°, of the amine radical cation are approximately related via eq 3 where the constant, 4.44 ((0.02) eV,11 is the absolute potential of the hydrogen electrode in water.12
∆G°solv (R2NH) - ∆G°solv (R2NH•+) Igas ≈ 4.44 + E° + F (3) Consequently, it is possible to estimate the difference in free X
Abstract published in AdVance ACS Abstracts, October 1, 1996.
S0022-3654(96)01286-5 CCC: $12.00
energy of solvation between an amine and the corresponding amine radical cation from the gas-phase ionization potential of the amine and the one-electron reduction potential of the amine radical cation. It should be noted that the ionization potential is the enthalpy of ionization at 0 K; thus, the ionization entropy and the temperature correction are neglected in eq 3. However, these corrections are assumed to be fairly small. Another important parameter is the N-H bond dissociation enthalpy of the corresponding amine (reaction 4) which, to a great extent, governs the reactivity of amines and aminyl radicals in hydrogen atom transfer reactions. It should be noted, however, that the rates of hydrogen atom transfer reactions can be significantly affected by the solvent13 and also that aminyl radicals in general are rather weak hydrogen atom abstractors.
R2NH f R2N• + H•
(4)
All these thermochemical properties are internally related via Hess’ law and thus, taking solvation into account, the N-H bond dissociation enthalpy of the amine can be calculated from the pKa of the amine in combination with the reduction potential of the aminyl radical or from the pKa of the amine radical cation in combination with the reduction potential of the amine radical cation.14 Equation 5 can be used to calculate the gas-phase bond dissociation enthalpy, D, from the redox and acidity properties in solution.
D ) 96.48E° + 5.70pKa + C
(5)
C in eq 5 is a constant that depends on the family of compounds and the solvent. The critical point here is the assumption of equal solvent effects on substituted and unsubstituted members of the family.15 It should be noted that N-substituted amines © 1996 American Chemical Society
17540 J. Phys. Chem., Vol. 100, No. 44, 1996 are not to be considered as one uniform family of compounds as will be shown later on. In recent years, several studies of the redox and acidity properties of aromatic aminyl radicals and amine radical cations have been conducted using experimental techniques such as cyclic voltammetry16 and pulse radiolysis.17-20 These studies have resulted in increased understanding of the substituent effects on the one-electron reduction potentials and pKas of arylamine radical cations and of the N-H bond dissociation enthalpies of the corresponding amines. Armstrong et al.21 have also performed a high-level quantum chemical study of the redox and acidity properties of alkylaminyl radicals and alkylamine radical cations where they pointed out some possible problems with previously determined reduction potentials and pKas of alkylamine radical cations. Possible problems with the electrochemistry of amines are that the oxidation processes are irreversible for most amines and also that the amines, or rather the oxidized amines, tend to polymerize on the electrodes which results in nonthermodynamical oxidation/reduction potentials. However, the latter problem can be reduced by carefully polishing and cleaning the working electrode after each cycle. In this work, we have determined the pKas of six alkyl- and arylamine radical cations in aqueous solution using laser flash photolysis. We have also measured the oxidation potentials of six alkyl- and arylamines in acetonitrile and in aqueous solution using cyclic voltammetry. Reduction potentials of the diphenylaminyl radical and methylphenylaminyl radical were measured using photomodulation voltammetry. On the basis of these results and literature data, structural and solvation effects on the thermochemical properties of aminyl radicals (reduction potentials) and amine radical cations (reduction potentials and pKas) are discussed. Experimental Section All chemicals were of the purest grade available (Sigma and Aldrich) and were used as supplied with the exception of aniline which was distilled and diphenylamine which was recrystallized prior to use. The purity of piperidine, pyrrolidine, diethylamine, and N-methylaniline was 99.5+% (Aldrich). For the laser flash photolysis experiments, the solvent, water (Omnisolv), was used as supplied. For the electrochemical experiments, water was deionized and distilled and acetontrile (Omnisolv) was distilled from CaH2 under 1 atm of argon prior to use. The supporting electrolyte TBAP was recrystallized three times from 10% hexane in ethyl acetate and dried in a vacuum oven (40 °C, 10 Torr). pKa Measurements. Aminyl radicals were generated upon 266 nm photolysis of N-nitrosamines in aqueous solutions using a Lumonics HY-750 Nd:YAG laser (8 ns pulse width; ca. 45 mJ/pulse). The UV-vis computer-controlled detection system has been described previously.22 The optical spectra of amine radical cations and the corresponding aminyl radicals differ markedly, making determination of radical pKa values possible. The pKa curve is obtained simply by measuring the absorbance at a wavelength where the spectral difference is optimal for a number of different pH values.19 The pH was varied using phosphate buffers, citrate, trifluoroacetic acid, acetate, carbonate, and NaOH. The solutions were purged with nitrogen. Cyclic Voltammetry. Cyclic voltammetry was performed with a PAR 273A potentiostat/galvanostat interfaced to a 386 based PC using the EG&G Model 270 software package. The cell was a standard three-electrode setup using a 3 mm diameter glassy carbon working electrode, a platinum coil counter electrode, and a reference electrode consisting of a silver wire in a glass tube with a Luggin capillary containing the solvent
Jonsson et al. TABLE 1: pKa Values and Measuring Wavelengths of Amine Radical Cations in Aqueous Solution amine
pKa
λ (nm)
diphenylamine
3.6 ( 0.2 3.926 7.6 ( 0.3 7.526 7.0517 6.8 ( 0.5 5.3 ( 0.5 5.5 ( 0.5 5.8 ( 0.5
650
N-methylaniline aniline dimethylamine diethylamine pyrrolidine piperidine
450 300 300 320 300
and supporting electrolyte used in the experiment. For the measurements in acetonitrile and water, 0.1 M tetrabutylammonium perchlorate (TBAP) and 1 M KCl, respectively, were used as supporting electrolytes. The aqueous solutions were buffered with 0.1 M phosphate buffers to avoid protonation of the arylamines. For the experiments involving dialkylamines the pH was adjusted to ca. 12.5 by adding KOH. All solutions were purged with argon for 20 min. The scan rate in all experiments was 200 mV/s, and the working electrode was carefully polished and cleaned after each cycle. Potential calibration in acetonitrile and water was accomplished using ferrocene (E° ) 0.44 V vs SCE23) and ferricyanide (E° ) 0.37 V vs NHE24), respectively. Photomodulation Voltammetry. The instrument has been described in detail previously.25 Aminyl radicals were generated in the electrochemical cell by photolysis, through an optically transparent gold electrode, of the corresponding N-nitrosamines in acetonitrile containing TBAP (0.1 M). The output from the lamp was modulated with a light chopper so the light intensity (and therefore the radical concentration) rose and fell as a sine wave. The electrochemical cell was fully iR compensated and was controlled with a PAR Model 174 polarographic analyzer. The ac component of the faradaic current was detected with a Stanford Research Systems Model SR530 lock-in amplifier, and plots of the ac current, as a function of potential, were recorded on a HP 7045 B X-Y recorder. All experiments were made at 100 Hz modulation frequency, and the scan rate was 20 mV/ s. In these experiments, ferrocene also was used for calibration. Results and Discussion pKas of Amine Radical Cations in Water. The aqueous pKas of the amine radical cations determined by means of laser flash photolysis are given in Table 1 along with the experimental uncertainty and the measuring wavelength. For comparison, the pKas of the diphenylamine, N-methylaniline, and aniline radical cations determined by means of pulse radiolysis in water are also given. The corresponding values determined by means of indirect electrochemical methods in DMSO are 2.7, 4.2, and 6.4, respectively.16 A typical pKa curve is given in Figure 1. The experimental uncertainty is somewhat greater for the alkylamines than for the arylamines which can simply be accounted for by the relatively weak absorbances of the alkylamine radical cations. As can be seen, the pKas of the diphenylamine radical cation and the N-methylaniline radical cation determined by laser flash photolysis are essentially identical to the values obtained by pulse radiolysis.26 The two methods are thus complementary and either one of them can be used for diaryl- and alkylarylamines. A potential problem when determining the pKas of alkylamine radical cations is that deprotonation can occur on both the nitrogen and the R-carbon.27 In the latter case, an aminoalkyl radical is formed upon deprotonation of the radical cation. Unless proton equilibrium for one of the two radical
pKa of Alkyl- and Arylamine Radical Cations
Figure 1. pH titration curve for the diphenylamine radical cation (650 nm).
types is much slower than for the other type, the radical that corresponds to the lowest pKa should be formed upon deprotonation of the radical cation. Since O2 reacts much more rapidly with carbon-centered radicals than with nitrogen-centered radicals, the fact that O2 saturation of the solution does not affect the determined pKa of the diethylamine radical cation indicates that deprotonation takes place predominantly at the nitrogen rather than the R-carbon under the present conditions (≈10 mM buffer concentration, pH 3-9). An interesting observation here is that the pKas of the dialkylamine radical cations do not vary dramatically. The pKa for the dimethylamine radical cation, 6.8 ( 0.4, is within the range suggested by Fessenden and Neta from ESR studies, 6.57.5.28 Armstrong et al.21 have suggested a pKa of 8 based on thermochemical calculations. Judging from the direct experimental determination of the pKa presented in this paper, a pKa of 8 is clearly too high. Horner et al.9 have also determined the pKas of two dialkylamine radical cations using a kinetic method where the observed rate constants for reactions of the amine radical cations were studied as a function of pH. The pKas thus determined are 6.3 ( 0.1 and 7.41 ( 0.06 for the N-(2,2-diphenylethyl)-N-ethylamine radical cation and the N-(6,6diphenyl-5-hexenyl)-N-methylamine radical cation, respectively. These two radical cations are, from an acidity point of view, expected to be similar to the diethylamine radical cation and the ethylmethylamine radical cation. The pKa values are, however, ca. 0.5-1 pK units higher than what would be expected from the data presented in this paper. For the arylamine radical cations, we see that the additional N-methyl substituent (N-methylaniline) only has a marginal effect on the pKa compared to the effect of a single N-phenyl substituent (aniline). The effect of an additional N-phenyl substituent (diphenylamine) is considerable, however. The pKas of N-substituted aniline radical cations roughly depend on the field/inductive constant29 of the N-substituent judging from these data. One-Electron Reduction Potentials of Amine Radical Cations. The one-electron reduction potentials of the amine radical cations (peak oxidation potentials of amines) determined using cyclic voltammetry are given in Table 2 along with the gas-phase ionization potentials30 of the corresponding amines. For comparison, the one-electron reduction potentials of the diphenylamine, N-methylaniline, and aniline radical cations determined by means of pulse radiolysis in water19,26 and by means of cyclic voltammetry in DMSO16 are also given in the table. A literature value for dimethylamine in aqueous solution is also given.3 The potentials measured in acetonitrile and DMSO were converted into V vs aqueous NHE by using the known potential of ferrocene vs SCE23 in each of the solvents and the difference between SCE and aqueous NHE.
J. Phys. Chem., Vol. 100, No. 44, 1996 17541 As can be seen, the peak potentials of the arylamines measured in aqueous solution by means of cyclic voltammetry are about 200 mV lower than the corresponding one-electron reduction potentials determined by means of pulse radiolysis. However, the relative trends are the same. The reason for this discrepancy is probably that in the electrochemical experiments the peak potentials are disturbed by the rapid deprotonation of the arylamine radical cations facilitated by the protic solvent.31 The peak oxidation potentials of diethylamine, pyrrolidine, and piperidine in aqueous solution are more or less identical; however, it should be stressed that the potentials are very close to the practical upper limit of cyclic voltammetry experiments in water. Interestingly, these potentials are very close to the previously reported potential of dimethylamine, 1.27 V vs NHE,3 and to a very recently pulse radiolytically determined value for dimethylamine, 1.30 ( 0.05 V vs NHE.26 The fact that the potential determined electrochemically agrees so well with the potential determined by means of pulse radiolysis may simply be coincidental since follow-up reactions of the radical cations will result in kinetic shifts in the peak potential to more negative values (up to ca. 200 mV).31 This kinetic shift can often be attenuated or even eliminated by unfavorable heterogeneous kinetics. If these factors are similar (i.e., within an order of magnitude) for all of the amines, then the relative peak potentials will reflect the relative standard potentials. The experimentally determined oxidation potentials of dimethylamine are much higher than the corresponding value calculated by Armstrong et al. (0.80 V vs NHE).21 As a general approximation, it seems reasonable to set the one-electron reduction potential of a dialkylamine radical cation in aqueous solution to 1.30 ( 0.05 V vs NHE. It is interesting to note that the potentials given in Table 2 correlate with the reactivity of some of these amine radical cations toward alkenes.10 Free Energies of Solvation. From the oxidation potentials of the amines (assuming that the peak potentials in DMSO and acetonitrile are close to the standard potentials) and the corresponding gas-phase ionization potentials, we can estimate the differences in solvation free energy between the amines and the amine radical cations using eq 3. The calculated solvation energy differences are given in Table 2. Judging from the results of the measurements in acetonitrile and water, the general trend seems to be that the difference in free energy of solvation, ∆∆Gsolv, increases with increasing gas-phase ionization potential of the amine. The rationale for this trend is probably the following: for amines with higher ionization potentials the charge of the amine radical cation formed upon ionization is more localized on the nitrogen than for the amines with lower ionization potentials.32 As the charge becomes more localized on the nitrogen, the solvation of the amine radical cation tends to increase. It is not unreasonable to suggest that, as the ionization potential increases, the N-H bond polarity decreases and the pKa will increase.33 This, in turn, makes hydrogen bonding between the amine and the solvent less favorable.34 The solvation free energy differences, ∆∆Gsolv, for dimethylamine in aqueous solution can be estimated to be 244 kJ/mol from the experimentally determined oxidation potential of 1.27 V vs NHE. The value obtained from the pulse radiolytically determined potential is 241 kJ/mol. The ∆∆Gsolv values for diethylamine, pyrrolidine, and piperidine do not vary significantly from the value for dimethylamine. One-Electron Reduction Potentials of Aminyl Radicals in Acetonitrile. The only aminyl radical reduction potentials that were practically possible to measure using photomodulation voltammetry were those for the diphenylaminyl radical (-0.16 V vs NHE) and the N-methylanilinyl radical (-0.54 V vs NHE).
17542 J. Phys. Chem., Vol. 100, No. 44, 1996
Jonsson et al.
TABLE 2: Gas-Phase Ionization Potentials of Amines and One-Electron Reduction Potentials and Solvation Energies of the Corresponding Amine Radical Cations in Aqueous Solution, Acetonitrile, and DMSO H2O 29
amine
IP (eV)
diphenylamine
7.16 ( 0.04
N-methylaniline
7.35 ( 0.03
aniline
7.720 ( 0.002
dimethylamine diethylamine pyrrolidine piperidine
8.24 ( 0.08 8.01 ( 0.01 8.0 8.20 ( 0.05
a
E°a
MeCN
DMSO
∆∆Gsolv
E°a
166
1.24 ( 0.05
143
1.1416
189
1.14 ( 0.05
171
1.1316
172
218
1.29 ( 0.05
192
1.1316
207
244 213 222 233
1.38 ( 0.05 1.09 ( 0.05 1.36 ( 0.05
211 238 232
b
0.83 (1.00)26 0.73 (0.95)26 0.83 (1.02)19 1.273 1.36 ( 0.05 1.26 ( 0.05 1.34 ( 0.05
∆∆Gsolv
E°a
∆∆Gsolvb 152
b
In V vs NHE. b In kJ mol-1.
The upper limit reduction potential of the dialkylaminyl radicals in acetonitrile is -1.46 V vs NHE. Due to the instability of monosubstituted N-nitrosamines, the reduction potential of the anilinyl radical could not be measured. Bordwell et al. have measured the reduction potentials for a large number of aromatic aminyl radicals in DMSO and the values for the diphenylaminyl radical, the N-methylanilinyl radical, and the anilinyl radical are -0.18, -0.37, and -0.31 V vs NHE, respectively.16 By assuming the relative trend to be unaltered by solvent, we estimate the reduction potential for the anilinyl radical in acetonitrile to -0.41 V vs NHE. Judging from these data, the potential range seems to increase going from DMSO to acetonitrile; e.g., the difference in reduction potential between the diphenylaminyl radical and the N-methylanilinyl radical is 0.19 V in DMSO and 0.38 V in acetonitrile. As for the pKas of the corresponding radical cations, the one-electron reduction potentials of N-substituted anilinyl radicals seems depend on the field/inductive properties of the N-substituents. A linear relationship between the field/inductive constant of R and the one-electron reduction potential has recently been found for peroxyl radicals with the general formula ROO•.35 Solvent Effects on the Redox and Acidity Properties of Aromatic Amine Radical Cations. As has been shown previously, both the pKas and the one-electron reduction potentials of 4-substituted aniline radical cations are linearly dependent on the Brown substituent constant, σp+, of the + 4-substituent (eqs 6a and 6b).16,19 F+ 1 and F2 signify the + E° ) E°0 + F+ 1 σp
(6a)
+ pKa ) pKa0 + F+ 2 σp
(6b)
magnitude of the substituent effects on the one-electron reduction potential and the pKa, respectively, and E°0 and pKa0 are the one-electron reduction potential and pKa of the unsubstituted + aniline radical cation. To put F+ 1 and F2 on the same energy scale, and thus make them directly comparable, F+ 1 should be multiplied by 2.303RT/F. Experimental data on the reduction potentials of 4-substituted aniline radical cations in water, DMSO, and acetonitrile and on the pKas of 4-substituted aniline radical cations in water and DMSO indicate a solvent dependence on both the intercepts, + E°0 and pKa0, and the slopes, F+ 1 and F2 , of eqs 6a and 6b (assuming that the peak potentials in DMSO and acetonitrile are close to the standard potentials or at least reflect the general trends). The reduction potential of the unsubstituted aniline radical cation increases going from water to DMSO and to acetonitrile, 1.02, 1.13, and 1.29 V vs NHE, respectively. The 19,20 0.48,16 same trend is observed for the slopes, F+ 1 ()0.33, 36 and 0.78 in water, DMSO, and acetonitrile, respectively). The
pKa of the unsubstituted aniline radical cation is 7.0517 in water and 6.416 in DMSO and the corresponding slopes, F+ 2 , are -3.719,20 and -5.4,16 respectively. Judging from the redox data, the solvent property that seems to govern both the relative trends and the absolute potentials is the solvent dipolarity/polarizability given by the constant π* (π* ) 1.09, 1.00, and 0.75 for water, DMSO, and acetonitrile, respectively).37 The relationships between the solvent dipolarity/polarizability and both the absolute reduction potential of the unsubstituted aniline radical cation and the magnitude of the substituent effects, F+ 1 , appear to be linear. The fundamental reason for a linear relationship between the potential of the unsubstituted aniline radical cation and the solvent dipolarity/polarizability would be that ∆∆Gsolv for the aniline radical cation depends linearly on π*.37 We intend to explore this further in the near future since a quantification of the solvent dielectric effects would undoubtedly be very useful for the understanding of reactions in heterogeneous and interfacial systems, e.g., biological membranes. An interesting observation here is that the magnitude of the substituent effect, F+ 1 , on the one-electron reduction potentials of 4-substituted aniline radical cations, 4-X-C6H4NH2•+, increases with the solvent induced increase in reduction potential of the parent aniline radical cation, C6H5NH2•+. The same phenomenon has been observed when comparing the magnitude of the substituent effect on the one-electron reduction potential of substituted benzene radical cations where the parent compounds have different reduction potentials in a given solvent (water); e.g., the magnitude of the effect of 4-substitution of anisole (E° ) 1.62 V vs NHE) is larger than the effect of 4-substitution of aniline.38 Solvent-induced changes in the oneelectron reduction potentials of the parent compounds followed by changes in the magnitude of the substituent effects have been observed previously for 4-substituted aryl methyl chalcogenide radical cations.39,40 The reduction potentials of C6H5SCH3•+, C6H5SeCH3•+, and C6H5TeCH3•+ are higher in acetonitrile39 than in aqueous solution40 and the substituent effects on the reduction potentials are larger in acetonitrile than in water which is exactly the same trend as for aniline radical cations. The trends for the reduction potentials of the unsubstituted diphenylamine and N-methylaniline radical cations in water, DMSO, and acetonitrile are similar to the trends found for aniline.41 Since pKas of substituted arylamine radical cations have only been determined in aqueous solution and in DMSO, it is difficult to explore the solvent effects to the same extent as for the oneelectron reduction potentials. However, the relative pKas, i.e., the magnitude of the substituent effects on the pKa, F+ 2 , in a given solvent can be estimated from the magnitude of the substituent effects on the reduction potential in the solvent of interest and the magnitudes of the substituent effects on the reduction potential and pKa in another solvent using eq 5. The
pKa of Alkyl- and Arylamine Radical Cations rationale for this is that the magnitude of the substituent effects on the N-H bond dissociation enthalpies should be the same regardless of the solvent in which E° and pKa were determined. For anilines, the magnitude of the substituent effects (vs σp+) on the N-H bond dissociation enthalpy has been determined to be 12.1 and 12.6 from redox and acidity data in water19 and DMSO,16 respectively. The difference between these values is smaller than the experimental uncertainty; thus, the assumption on which eq 5 is based seems to be adequate for anilines. From the redox and acidity data in water and in DMSO we estimate the magnitude of the substituent effects, F+ 2 , on the pKa of 4-substituted aniline radical cations in acetonitrile to -10.8 (average of the values calculated from data in water and DMSO). Estimation of N-H Bond Dissociation Enthalpies of Arylamines from Redox and Acidity Data. As mentioned previously, the N-H bond dissociation enthalpies can be calculated using eq 5. The problem here is to determine the constant C for each type of molecule. By using the gas-phase N-H bond dissociation enthalpies of aniline, N-methylaniline, and diphenylamine, C can be calculated for each family of compounds and for each solvent. The N-H bond dissociation enthalpies of aniline, N-methylaniline, and diphenylamine have been reported to be 373,19,42 372,43 and 364 kJ/mol,42 respectively. Since the pKas of the radical cations of aniline, N-methylaniline, and diphenylamine are unknown in acetonitrile, we make the assumption of a linear correlation between the reduction potential and the pKa of the radical cations in different solvents. The thus estimated pKas of the radical cations of aniline, N-methylaniline, and diphenylamine in acetonitrile are 5.5, 4.2, and 1.8, respectively. Our extrapolated constants, C, for aniline in acetonitrile, DMSO, and water are 217.2, 227.5, and 234.4 kJ/mol, respectively, C for diphenylamines are 233.9, 238.6, and 245.3 kJ/mol, respectively, and C for N-methylaniline are 238.1, 239.0, and 237.6 kJ/mol, respectively. As can be seen, C depends both on the solvent and on the parent molecule. Using the same C for all three types of arylamines would thus not be recommended since this would result in systematic error on the order of (10 kJ/mol. Conclusions Most of the properties discussed in this paper can, at least qualitatively, be shown to depend on the substituent constants of the N-substituents (field/inductive constants, σI) or the ring substituents (Brown constants, σp+). The general and most important observation from this work is that the absolute values and the magnitude of the substituent effects on the pKas and one-electron reduction potentials of arylamine radical cations and the one-electron reduction potentials of aminyl radicals appear to be solvent dependent. This solvent dependence can, as a first approximation, be described by linear solvation energy relationships as a function of dipolarity/polarizability, π*. Similar solvent effects on both the absolute values and the relative substituent effects have also been found for other classes of aromatic radical cations, suggesting that this may be a general phenomenon for the redox properties of aromatic radical cations in solution. Warning! N-nitrosamines are suspected to be potent carcinogens which must be handled, stored, and discarded with due respect for the possible hazards involved. Acknowledgment. M.J. thanks the Swedish Natural Science Research Council for financial support. We also thank Dr.
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