Nucleophilic Reactions at Cationic Centers in Ionic Liquids and

nucleophiles (n-butylamine, di-n-butylamine, and tri-n-butylamine) in several molecular solvents (toluene, dichloromethane, tetrahydrofuran, acetonitr...
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Ind. Eng. Chem. Res. 2008, 47, 638-644

Nucleophilic Reactions at Cationic Centers in Ionic Liquids and Molecular Solvents Giuseppe Ranieri, Jason P. Hallett, and Tom Welton* Department of Chemistry, Imperial College London, South Kensington Campus, London SW7 2AZ, United Kingdom

We report a quantitative comparison of the rates of nucleophilic reactions in ionic liquids and molecular solvents taking place at a cationic center. Cationic sulfonium electrophiles were reacted with three amine nucleophiles (n-butylamine, di-n-butylamine, and tri-n-butylamine) in several molecular solvents (toluene, dichloromethane, tetrahydrofuran, acetonitrile, and methanol) and ionic liquids {1-butyl-1-methylpyrrolidinium bis(trifluoromethanesulfonyl)imide, 1-butyl-1-methylpyrrolidinium trifluoromethanesulfonate, 1-butyl-3-methylimidazolium bis(trifluoromethanesulfonyl)imide, and 1-butyl-3-methylimidazolium trifluoromethanesulfonate}. The solvent effects on these reactions are examined using a linear solvation energy relationship based on the Kamlet-Taft solvent scales (R, β, and π*). These correlations reveal that hydrogen-bonding interactions provide the dominant effects in determining the rate of reaction. In particular, hydrogen bonds donated by the sulfonium electrophile to the solvent are the most important controlling factor on the rate of nucleophilic reaction. 1. Introduction Ionic liquids are an exciting class of solvents for organic reactions1-8 because of a variety of interesting properties, such as negligible vapor pressure, ability to dissolve organic, inorganic, and polymeric materials, and high thermal stability. They are also widely considered to be “green” solvents, and they are being investigated as a potential environmentally friendly alternative to volatile organic compounds. However, it should be noted that this is a matter of some current contention.9 The wide range of cations and anions available for use, combined with the often stark changes in solvent properties that different combinations of these bring about, has led to ionic liquids being labeled “designer solvents”.10 Our purpose in this and other recent studies11-16 is to elucidate the effect of select cations and anions on reactive processes in ionic liquids in order to enable optimum “designing” of ionic liquid solvents for specific applications. This will allow a predictive basis for deciding when ionic liquids provide substantial processing advantages over traditional molecular solvents as well as for the development of a quantitative scientific approach to determining what effect a particular ionic liquid will have on a given chemical transformation. Nucleophilic substitution reactions provide a familiar avenue for examination of solvent effects on chemical reactions.17,18 In molecular solvents, these studies originated in the work of Hughes and Ingold.19-23 The Hughes-Ingold theory describes solvent effects on chemical reactions using a simple qualitative solvation model considering only pure electrostatic interactions between ions or dipolar molecules in initial and transition states. According to this model, high-polarity solvents accelerate reactions in which charge density is created, decelerate reactions in which charge density is destroyed, and have a negligible effect on reactions that involve little or no change in the charge density on going from reactants to the transition state. Most deviations from the behavior laid out in the Hughes-Ingold rules can be attributed to specific interactions, such as hydrogen bonding, which are not taken into account. Applications of the Hughes* To whom correspondence should be addressed. E-mail: t.welton@ imperial.ac.uk.

Ingold approach to ionic liquids for a variety of reaction classes have had some limited success.17,24-28 We have previously reported quantitative kinetic studies of nucleophilic substitutions by both halides14-16 and amines13 in various ionic liquids and also detailed the application of the Hughes-Ingold approach for SN2 reactions in ionic liquids wherein a neutral electrophile, methyl-p-nitrobenzenesulfonate, with a range of neutral (mono-, di-, and tributylamine) and anionic (various halides and other ions) nucleophiles, was employed.11 It was concluded that a pure Hughes-Ingold view of the system was not sufficient, as hydrogen-bonding interactions were often the dominant solvent-solute effects governing the behavior of the system. Instead, a Kamlet-Taft linear solvation energy relationship (LSER)18 approach was utilized to effectively describe the solvent effects on these nucleophilic reactions. For all reactions with the neutral methyl-p-nitrobenzenesulfonate, hydrogen-bond donation from the solvent to the nucleophile (represented by the solvent’s Kamlet-Taft R value) was found to be more important than simple dipolar solventsolute interactions with respect to the rate of reaction; increasing R significantly reduces the rate of nucleophilic reaction by deactivating the nucleophile. In the case of the anionic nucleophilic reactions, increasing the hydrogen-bond accepting ability of the solvent (β) was found to have a small positive effect on the reaction rate, as did increasing the solvent dipolarity/ polarizability (π*); however, these effects were concluded to be much less important than the hydrogen-bond donor properties of the solvent. In addition, a hard/soft analysis of the anions was undertaken, resulting in a classification of the ionic liquids studied as “hard” solvents. In the case of the amine-based nucleophiles reacting with methyl-p-nitrobenzenesulfonate,13 the dominant effect on the rate of reaction was again found to be the solvent’s hydrogenbond donor (R) ability, which acts to slow the reaction by reducing the nucleophilicity of the amine. Increasing the solvent’s hydrogen-bond acceptor (β) ability was found to accelerate the reactions. Dipolarity/polarizability (π*) was found to be far less important, giving slightly increased rates of reaction for greater values of π*. The current work expands on these previous studies by examining the role of a charged electrophile within the framework of SN2 reactions performed

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in ionic liquids and determines how accurately a Kamlet-Taft description of the solvent effects captures this system. The results are also of fundamental interest in catalysis, where the (electrophilic) catalytic center often carries a positive charge. The Kamlet-Taft model comprises three main solvent descriptors, π* (dipolarity/polarizability), R (hydrogen-bond donating acidity), and β (hydrogen-bond accepting basicity). These three parameters, determined from spectral shifts of solvatochromic indicators, have been collected for a variety of systems29 including ionic liquids.30-32 The Kamlet-Taft approach encompasses solute-solvent specific interactions, thus making it ideally suited for studying systems where the Hughes-Ingold model alone is ineffective. The most common form of the Kamlet-Taft LSER for quantifying solvent effects on chemical reactions is shown in eq 1. In this equation, XYZ represents the solvent-dependent property under study (in the present case, the logarithm of a rate constant), and a, b, and s are solute-specific coefficients characteristic of the process and measure the sensitivity of the property (XYZ) to each particular solvent descriptor (R, β, and π*, respectively). This approach has been used a number of times previously,33 including for the correlation of reaction rate data, as presented here. The R, β, and π* values for each of the ionic liquids (and also molecular solvents) used in the present study were reported earlier.30-32

XYZ ) XYZ0 + aR + bβ + sπ*

(1)

2. Experimental Section 2.1. Materials and Reagents. 1-Methylimidazole and 1-methylpyrrolidine were purchased from Acros Organics and distilled from potassium hydroxide; 1-chlorobutane was purchased from Acros Organics and distilled from phosphorus pentoxide. Lithium bis(trifluoromethylsulfonyl)imide and lithium trifluoromethanesulfonate were purchased from Solvent Innovation GmbH and used as received. 4-Nitrothioanisole was purchased from Avocado and used as received. All molecular solvents were purified by distillation from standard drying agents. All syntheses were performed under anaerobic conditions using standard Schlenk techniques. The preparations and spectral data of the ionic liquids have been described elsewhere.34 2.1.1. Synthesis of Dimethyl-4-nitrophenylsulfonium Methylsulfate. Dimethylsulfate (2.52 cm3, 26.7 mmol) was added dropwise at room temperature to p-nitrothioanisole (3.0 g, 17.7 mmol). The resulting solution was heated to 95 °C and stirred for 2 h. After cooling, a yellowish solid precipitated, which was recrystallized twice from isopropanol to give soft cream-white crystals. δH (ppm) (DMSO-d6): 8.42 (4H, 2 d, CH-2,3,5,6), 3.37 (3H, s, CH3-OSO3), 3.33 (6H, s, -S-CH3). δC (ppm) (DMSOd6): 150.18 (CS(CH3)2), 134.10 (CNO2), 131.46 (CH-2,6), 124.86 (CH-3,5), 61.15 (CH3OSO3), 30.11 (S(CH3)2). mp ) 152 °C. UV-vis: λmax(CH2Cl2) 252 nm. Elemental analysis, found (calcd): %C ) 36.51 (36.60), %H ) 4.51 (4.44), %N ) 4.68 (4.74). 2.1.2. Synthesis of Dimethyl-4-nitrophenylsulfonium bis(Trifluoromethylsulfonyl)imide. Dimethyl-4-nitrophenylsulfonium methylsulfate (3.02 g, 10 mmol) and lithium bis(trifluoromethylsulfonyl)imide (2.94 g, 10 mmol) were stirred in water for 8 h at room temperature. From the solution, a heavier liquid separated, which was isolated and washed with CH2Cl2 and then with water. Drying under high vacuum gave the product as a viscous yellow liquid. δH (ppm) (Acetone-d6): 8.54 (4H, 2 d, CH-2,3,5,6), 3.65 (6H, s, -S-CH3). δC (ppm) (Acetone-d6): 153.09 (CS(CH3)2), 134.92 (CNO2), 133.46 (CH-2,6), 127.25 (CH-3,5), 121.98 (1J13C-19F ) 319.9 Hz,

Scheme 1. Methylation Reaction of a Nucleophile by a Sulfonium Species

[N(SO2CF3)2]-), 30.11 (S(CH3)2). UV-vis: λmax(CH3CN) ) 249.0 nm. m/z (FAB+) MS: 648, [(pNO2PhS(CH3)2)2(N(SO2CF3)2)]+, 0.65%; 184, [pNO2PhS(CH3)2]+, 100%. Elemental analysis, found (calcd): %C ) 25.80 (25.86), %H ) 2.09 (2.17), %N ) 5.89 (6.03). 2.1.3. Synthesis of Dimethyl-4-nitrophenylsulfonium Trifluoromethanesulfonate. Dimethyl-4-nitrophenylsulfonium methylsulfate (3.02 g, 10 mmol) and lithium trifluoromethanesulfonate (1.56 g, 10 mmol) were stirred in water for 3 h at room temperature. After extracting the solution with CH2Cl2 (2 × 20 cm3), the aqueous phase was concentrated and washed several times with ethyl acetate. The organic phase was then evacuated to give a yellowish liquid, which was suspended in diethyl ether to provide a white solid precipitate. δH (ppm) (D2O): 8.47 and 8.16 (4H, 2 d, CH-2,3,5,6), 3.28 (6H, s, -SCH3). δC (ppm) (D2O): 150.83 (CS(CH3)2), 132.20 (CNO2), 130.80 (CH-2,6), 125.61 (CH-3,5), 119.56 (1J13C-19F ) 315.0 Hz, [OSO2CF3]-), 28.16 (S(CH3)2). mp ) 161 °C. UV-vis: λmax(H2O) ) 252.2 nm. m/z (FAB+) MS: 517, [(pNO2PhS(CH3)2)2(OCF3SO3)]+, 3.75%; 184, [pNO2PhS(CH3)2]+, 100%; m/z (FAB-) MS: 482, [(pNO2PhS(CH3)2)(OCF3SO3)2]-, 21.9%; 149, [OCF3SO3]-, 100%. Elemental analysis, found (calcd): %C ) 32.32 (32.43), %H ) 3.11 (3.02), %N ) 4.02 (4.20). 2.2. Kinetic Procedure. A stock solution containing the nucleophile (amine) dissolved in the solvent of interest was added to a 0.5 cm path length UV/vis quartz cuvette under anaerobic conditions, after which pure solvent was added to a total volume of 1.1 cm3. After thermostatting for 5 min, the reaction was initiated by injection of 1 µL of a stock solution containing the sulfonium electrophile dissolved in the solvent of interest. Spectra were recorded at regular time intervals using a spectrophotometer fitted with a thermostatted sample holder. All reported rate constants are the average of at least three identical experiments. 2.3. Analytical. 1H NMR and 13C NMR spectra were recorded on a Bruker 400 MHz spectrometer. Chemical shifts are reported in ppm (versus tetramethylsilane) and coupling constants in Hz. FAB mass spectra were recorded on a VG AutoSpec-Q mass spectrometer. UV-vis spectra were recorded on a Perkin-Elmer UV-visible Lambda 2 spectrophotometer. A thermostatic water circulator controlled the sample-holder temperature. 3. Results and Discussion The reaction chosen for study was the SN2 reaction of the trifluoromethanesulfonate and bis(trifluoromethanesulfonyl)imide salts of dimethyl-4-nitrophenylsulfonium ([p-NO2PhS(CH3)2]+[X]-; [X]- ) [CF3SO3]-, [N(CF3SO2)2]-) by various nucleophiles, as shown in Scheme 1. These substrates are particularly convenient for assaying the process by UV spectroscopy, as each has a λmax at 252 nm and the demethylated product has a λmax at 342 nm (see Figure 1), well-above the cutoff of either the imidazolium-based (240 nm) or pyrrolidinium-based (200 nm) ionic liquids that were used as solvents. In addition, similar substrates to these have been used for literature studies of SN2 reactions in molecular solvents,35 and they are similar to the neutral electrophile used previously by

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Figure 1. Sample absorbance spectra for a sulfonium salt and its corresponding demethylated sulfide in acetonitrile.

Scheme 2. Synthesis of One of the Sulfonium Salts (Dimethyl-4-nitrophenylsulfonium Trifluoromethanesulfonate)

our group,14-16 allowing for some comparison between the systems. The sulfonium salts were synthesized as described above, following a modified literature procedure (Scheme 2).36 We chose to compare the nucleophilicities of n-butylamine, di-n-butylamine, and tri-n-butylamine in several molecular solvents and ionic liquids. The ionic liquids (ILs) chosen were 1-butyl-1-methylpyrrolidinium bis(trifluoromethanesulfonyl)imide [bmpy][N(CF3SO2)2], 1-butyl-1-methylpyrrolidinium trifluoromethanesulfonate [bmpy][CF3SO3], 1-butyl-3-methylimidazolium bis(trifluoromethanesulfonyl)imide [bmim][N(CF3SO2)2], and 1-butyl-3-methylimidazolium trifluoromethanesulfonate [bmim][CF3SO3], and the molecular solvents were dichloromethane (CH2Cl2), tetrahydrofuran (THF), acetonitrile (CH3CN), methanol (CH3OH), and toluene (C7H8). Because of the solubility differences between the two salts in the different solvents used in this study, we were forced to utilize different anions of the sulfonium salt. The trifluoromethanesulfonate salt was used with acetonitrile, tetrahydrofuran, methanol, and the two trifluoromethanesulfonate-based ionic liquids. The bis(trifluoromethanesulfonyl)imide salt was used with acetonitrile, dichloromethane, toluene, and the two bis(trifluoromethanesulfonyl)imide-based ionic liquids. No rate differences were found between studies involving the two different sulfonium salts in acetonitrile; therefore, we have treated the two electrophiles as interchangeable, which is consistent with the view that the cationic center itself acts as the electrophilic site for reaction. Unfortunately, exploration in truly nonpolar solvents (e.g., alkanes) was impossible because of the extremely low solubility of either reactant electrophile in these solvents. It is important to note that both electrophiles were found to be stable in every solvent, because no change in UV-vis absorbance was observed in the absence of the nucleophile within 24 h. The wide range of solvents involved enables us to make comparisons across a wide range of solvent characteristics. From our measurement of pseudo-first-order rate constants (∼1001000× excess nucleophilic was used), we are able to calculate

Table 1. Rate Constant Values (for Amines with the Cationic Electrophile) and Kamlet-Taft Parameters in Each Solvent k2 (M-1 s-1)

CH2Cl2 CH3CN THF CH3OH C7H8 [bmpy][N(CF3SO2)2] [bmim][N(CF3SO2)2 ] [bmpy][CF3SO3] [bmim][CF3SO3]

Kamlet-Taft paramete rs

BuNH2

Bu2NH

Bu3N

R

β

π*

0.15 0.016 0.038 0.00064 0.24 0.060 0.026 0.030 0.013

0.25 0.030 0.028 0.00097 0.14 0.13 0.068 0.038 0.016

0.068 0.0092 0.0016 0.00042 0.0067 0.089 0.025 0.0089 0.012

0.13 0.19 0.00 0.98 0.00 0.43 0.62 0.40 0.62

0.10 0.40 0.55 0.66 0.11 0.25 0.24 0.46 0.46

0.73 0.66 0.55 0.60 0.49 0.95 0.98 1.02 1.01

the bimolecular rate constants, k2, which are collected in Table 1. The Hughes-Ingold method predicts that an increase in solvent dielectric constant (polarity) would decrease the reaction rate. A cursory examination of our results (Figure 2) reveals that this is not the case. The dielectric constant increases in the order C7H8 (2.38) < THF (7.61) < CH2Cl2 (9.08) < ILs (ca. 11-15) < CH3OH (32.7) < CH3CN (37.5),37-43 and while the smallest k2 was found in CH3OH in each case, the fastest was in toluene for the n-butylamine only (and in CH2Cl2 or [bmim][N(CF3SO2)2] for the others). The ordering in between was quite random, with the second smallest k2 found for [bmim][OCF3SO3] (with BuNH2 and Bu2NH) or THF (Bu3N) despite their intermediate dielectric values. The poor correlation of rate constant with dielectric constant is typified by Figure 3, which relates these values for di-n-butylamine. This clear violation of the Hughes-Ingold rules indicates that specific interactions (e.g., hydrogen bonding) are more important than simple electrostatic forces and led us to perform an in-depth examination of k2 with respect to the Kamlet-Taft solvent descriptors (R, β, and π*) in our further analysis. Interestingly, not only do the absolute rates of the reactions change but the relative nucleophilicities and the ordering of the nucleophilicities of the different amines in the various solvents also change. For example, with the amine BuNH2, a clear ordering of solvents (C7H8 > CH2Cl2 > [bmpy][N(CF3SO2)2] > THF > [bmpy][CF3SO3] > [bmim][N(CF3SO2)2] > CH3CN > [bmim][CF3SO3] > CH3OH) is observed, wherein the ionic liquids are not even grouped together despite their similar dielectric properties.38-43 With Bu2NH as the nucleophile, the ordering is slightly different, with toluene and CH2Cl2 reversed, the rate in [bmim][N(CF3SO2)2] now faster than that in either

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Figure 2. 3-D graph of second-order reaction rate for each amine with the sulfonium electrophile in various solvents. Table 2. Results of Kamlet-Taft LSER Fits for Reactions of n-Butylamine, di-n-Butylamine, and tri-n-Butylamine with Cationic Electrophile amine

parameters

BuNH2

R, β, π*

Bu2NH

Figure 3. Natural log of second-order reaction rate constant for nucleophilic reaction between the cationic electrophile and di-n-butylamine plotted against dielectric constant for each solvent.

Figure 4. Kamlet-Taft LSER fit of ln k2 for di-n-butylamine reaction with cationic electrophile.

THF or [bmpy][OCF3SO3], and the reaction in THF now slower than that in [bmpy][OCF3SO3]. With Bu3N, the solvent effects are more muted, with the rate in [bmpy][N(CF3SO2)2] ≈ CH2Cl2 > [bmim][N(CF3SO2)2] > [bmim][OCF3SO3] ≈ CH3CN ≈ [bmpy][OCF3SO3] ≈ C7H8 > THF). The position of toluene in this case is another strong refutation of Hughes-Ingold ordering, as the least polar solvent provides the second-worst medium for nucleophilic reaction. Since the simple dielectric model proved inadequate, the LSER of eq 1 was applied to rate data from reactions of each of the nucleophiles in these solvents (Table 1). The ionic liquids used had systematically varied cations and anions to allow

Bu3N

R, β, π*

β, π*

LSER equation ln k2 ) -2.38 - 3.59R - 4.16β + 2.10π*

ln k2 ) -2.66 - 2.79R 5.01β + 2.89 π*

ln k2 ) -5.62 6.46β + 4.26 π*

p-values

R2

XYZ0 ) 0.02

0.96

R ) 0.002 β ) 0.007 π* ) 0.05 XYZ0 ) 0.001

0.99

R ) 0.0005 β ) 0.0002 π* ) 0.002 XYZ0 ) 0.002

0.87

β ) 0.002 π* ) 0.01

investigation of the effects of altering the physiochemical properties of the ionic liquid.17 For each nucleophile, ln k2 was correlated using eq 1. The error associated with each parameter was appraised in terms of the p-value, and any terms found to be statistically insignificant were eliminated. Only where an acceptable value for a, b, or s is found is the result shown. It was decided to take as acceptable only those parameters whose statistical significance p does not exceed the limit level of ∼0.05. The results of these fits, along with the associated statistical data, are shown in Table 2. These results are graphically depicted in Figure 4 for the example of di-n-butylamine. In contrast to the very poor relationship observed in Figure 3, the KamletTaft approach provides an excellent correlation for our reaction rate data. In all cases, the rate constant is shown to be governed by the hydrogen-bond acceptor property of the solvent (β), with large negative coefficients in the LSER (b). This means that higher values of β for the solvent will produce a large decrease in k2. This is in complete contrast to the results that we found for the same amines when reacted with a neutral electrophile.13 This can be explained by taking into account the relatively high acidity of the methyl protons in the sulfonium cation; e.g., this species can be readily deprotonated by alkaline substances such as hydroxide or alkoxide salts to produce the corresponding

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Figure 5. Possible hydrogen bonding interactions in the activated complex of the SN2 reaction of dibutylamine and the dimethyl-4-nitrophenylsulfonium ion.

ylide. Solvents with a high β value will hydrogen bond with the methyl protons, thus attenuating the electrophilicity of the sulfonium ion (the electron density at the reactive center will increase). A consequence will be the deactivation of the substrate, resulting in a deceleration of the overall process. Although some of the bond lengths might be considered borderline, hydrogen bonding between methyl-substituted sulfonium cations and a variety of anions have been noted in a number of crystal structures44-47 and in the liquid phase.47 The product sulfide, with the absence of a positive charge, has methyl protons that are expected to be far less acidic (e.g., the methylene protons in the benzyl phenyl methyl sulfonium ion have a pKa of 16.3 in dimethyl sulfoxide (DMSO), whereas those in benzyl phenyl sulfide have a pKa of 30.8)48 and poorer hydrogen-bond donors than those of the sulfonium salt. Consequently, during the activation process, these protons are reducing in acidity and the hydrogen bonds from the solvent to these protons in the activated complex are weaker. Hence, the activated complex is less stabilized by hydrogen-bond acceptor solvents than the starting material. This leads to greater activation energy for the reaction and reduced rate in the strong hydrogen-bond acceptor solvents. Given the argument elaborated above and because the electrophile is common to all of the reactions studied here, it might be expected that the coefficient for β should be the same for all of the reactions. This is clearly not the case, with the coefficient increasing in the order Bu3N > Bu2NH > BuNH2: as the number of protons on the amine increases, the value of b increases. Our previous investigations11 have shown that hydrogen-bond donation from the amine to the solvent will increase the nucleophilicity of these amines, which is the exact reverse of the effect of the sulfonium/sulfide hydrogen bonding. Here, the relatively weak hydrogen-bond donating protons of the amines are becoming stronger hydrogen-bond donors as the positive charge on the nitrogen develops.49 This stabilizes the activated complex for the reaction with respect to the starting materials and accelerates the reaction. While, in the reactions studied here, this is insufficient to cause an increase in the overall rates, it does moderate the deceleration caused by the decreased reactivity of the electrophile. The negative β-dependence is most pronounced in the case of tri-n-butylamine, which contains no N-protons; therefore, the moderating effect from amine-based hydrogen-bond donation is not present. The different hydrogen-bond donating interactions to and from the SN2 activated complex from and to, respectively, the solvent are depicted in Figure 5 for the case of di-n-buylamine reacting in an ionic liquid solvent. From this diagram, the various hydrogen-bonding opportunities can be readily discerned. Those between the amine and the solvent (a) increase during the activation process, stabilizing the activated complex relative to

the starting materials and increasing the rate of the reaction, while those between the sulfonium ion and the solvent (b) decrease during the activation process, stabilizing the starting materials relative to the activated complex and decreasing the rate of the reaction. In the cases of BuNH2 and Bu2NH, a correlation between k2 and -R is found; the nucleophilic attack to the sulfonium ion is, therefore, slowed by solvents with high R values. The -R dependence can be explained by the deactivation of the amine through hydrogen-bond donation by the solvent to the lone pair of electrons on the amine. However, this dependence is not observed in the case of tri-n-butylamine. We postulate that this is due to the masking of the solvent hydrogen-bond donation ability by the overwhelming negative β-dependence observed in this case. Since tri-n-butylamine is incapable of activation through N-hydrogen-bond donation to the solvent (as opposed to di-n-butylamine, Figure 5 ), the b-coefficient is much larger than that calculated for the other two amines. This much larger effect may be overwhelming any (likely small) R-effect, at least on the scale of our data set. In all experiments, a positive π* dependence was observed, which is the opposite of what we would expect from our understanding of Hughes-Ingold behavior. This suggests that π* is a poor indicator of Hughes-Ingold polarity. In our previous study of amine nucleophilicity,11 the π* effect was also observed to be positive. In this case, however, the electrophile was neutral and, therefore, the Hughes-Ingold approach would predict a positive dependence on polarity. In the present study, a cationic electrophile is employed and, therefore, the Hughes-Ingold polarity dependence should be negative. These contrasting results highlight the inability of simple polarity models to capture complex solvent effects on nucleophilic reactions and also indicate that the Hughes-Ingold concept of “polarity” may have more to do with specific, hydrogen-bonding type interactions than with a Kamlet-Taft π* concept of dipolarity/polarizability. With the exceptions of THF and toluene, in all the studied solvents, the highest rates observed were those of di-nbutylamine. The nucleophilicity of n-butylamine was, in most cases, similar to that observed for di-n-butylamine, while for the tertiary amine, the decrease in nucleophilicity was more significant (reactions generally 1 order of magnitude slower). This result reflects the basicity trend (pKa) Bu2NH (11.0) > BuNH2 (10.7) > Bu3N (10.0).50 It is important to notice that the ∆pKa between mono- and di-n-butylamine is much smaller than the ∆pKa between these two amines and tri-n-butylamine. This provides an explanation for the poor nucleophilicity of trin-butylamine compared with the partially substituted amines. There is a clear cation/anion effect when only the ionic liquid solvents are considered. For all three nucleophiles, the ionic liquids with [bmpy]+ cations provide a superior medium to those containing the same anion with a [bmim]+ cation (except in one case where they were approximately equal). A cation paired with the [N(CF3SO2)2]- anion also provides superior rates to one coupled with the [CF3SO3]- anion. This is consistent with the above analysis because the [N(CF3SO2)2]- anion is less basic (lower β) than [CF3SO3]-, and [bmpy]+-based ionic liquids are less acidic (lower R) than [bmim]+-based ILs.30 4. Conclusion This work explored nucleophilic reactions involving a positively charged electrophile and neutral nucleophiles in ionic liquids and molecular solvents. The experimental results demonstrated that a simple Hughes-Ingold polarity viewpoint (using

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dielectric constant as a measure of solvent polarity) is inadequate to describe the system, while the Kamlet-Taft solvent parameters performed well, because specific hydrogen-bonding interactions are the dominant effect on reaction rates. In particular, there is a dominant β effect derived from the ability of the solvent to accept hydrogen bonds from different sites in the transition state of the reaction. Solvents with a low β value were found to accelerate the reaction between each of three different amines and a sulfonium ion. A less-intense solvent hydrogenbond donating (R) effect was observed for two of the amines (n-butylamine and di-n-butylamine). We postulate that the much stronger β effect masks this result in the case of tri-n-butylamine. Additionally, a positive dependence on the solvent’s dipolarity/ polarizability (π*) was observed, along with effects deriving from the ionic liquid’s constituent ions and nature of the amine nucleophile. The ionic liquids [bmpy][N(CF3SO2)2] and [bmim][N(CF3SO2)2] performed best among the ionic solvents employed, and besides low β values, they possess the processing advantages that only an ionic liquid can offer. Acknowledgment We wish to thank the EPSRC (G.R.) for a studentship and the Marshall-Sherfield Foundation (J.P.H.) for a research fellowship. Supporting Information Available: The Supporting Information includes details of the synthetic procedures and assays of ionic liquids and all kinetic data (PDF). This material is available free of charge via the Internet at http://pubs.acs.org. Literature Cited (1) Zhao, H.; Malhotra, S. V. Applications of Ionic Liquids in Organic Synthesis. Aldrichimica Acta 2002, 35, 75. (2) Olivier-Bourbigou, H.; Magna, L. Ionic Liquids: Perspectives for Organic and Catalytic Reactions. J. Mol. Catal., A 2002, 182, 419. (3) Zhao, D.; Wu, M.; Kou, Y.; Min, E. Ionic Liquids: Applications in Catalysis. Catal. Today 2002, 74, 157. (4) Sheldon, R. Catalytic Reactions in Ionic Liquids. Chem. Commun. 2001, 2399. (5) Gordon, C. M. New Developments in Catalysis Using Ionic Liquids. Appl. Catal., A 2001, 222, 101. (6) Wasserscheid, P.; Keim, W. Ionic LiquidssNew “Solutions” for Transition Metal Catalysis. Angew. Chem., Int. Ed. 2000, 39, 3772. (7) Welton, T. Room-Temperature Ionic Liquids. Solvents for Synthesis and Catalysis. Chem. ReV. 1999, 99, 2071. (8) Ionic Liquids in Synthesis; Welton, T., Wasserscheid, P., Eds.; VCHWiley: Weinheim, Germany, 2002. (9) Scammells, P. J.; Scott, J. L.; Singer, R. D. Ionic Liquids: The Neglected Issues. Aust. J. Chem. 2005, 58, 155. (10) Freemantle, M. Designer SolventssIonic Liquids May Boost Clean Technology Development. Chem. Eng. News 1998, 76 (13), 32. (11) Crowhurst, L.; Falcone, R.; Lancaster, N. L.; Llopsis-Mestre, V.; Welton, T. Using Kamlet-Taft Solvent Descriptors To Explain the Reactivity of Anionic Nucleophiles in Ionic Liquids. J. Org. Chem. 2006, 71, 8847. (12) Anderson, J. L.; Ding, J.; Welton, T.; Armstrong, D. W. Characterizing Ionic Liquids On the Basis of Multiple Solvation Interactions. J. Am. Chem. Soc. 2002, 124, 14247. (13) Crowhurst, L.; Lancaster, N. L.; Perez Arlandis, J. M.; Welton, T. Manipulating Solute Nucleophilicity with Room Temperature Ionic Liquids. J. Am. Chem. Soc. 2004, 126, 11549. (14) Lancaster, N. L.; Welton, T. Nucleophilicity in Ionic Liquids. 3. Anion Effects on Halide Nucleophilicity in a Series of 1-Butyl-3methylimidazolium Ionic Liquids. J. Org. Chem. 2004, 69, 5986. (15) Lancaster, N. L.; Salter, P. A.; Welton, T.; Young, G. B. Nucleophilicity in Ionic Liquids. 2. Cation Effects on Halide Nucleophilicity in a Series of Bis(Trifluoromethylsulfonyl)imide Ionic Liquids. J. Org. Chem. 2002, 67, 8855. (16) Lancaster, N. L.; Welton, T.; Young, G. B. A Study of Halide Nucleophilicity in Ionic Liquids. J. Chem. Soc., Perkin Trans. 2001, 2, 2267.

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ReceiVed for reView May 4, 2007 ReVised manuscript receiVed June 27, 2007 Accepted July 10, 2007 IE070632V