Article pubs.acs.org/JPCB
Dissociation of Equimolar Mixtures of Aqueous Carboxylic Acids in Ionic Liquids: Role of Specific Interactions Shashi Kant Shukla and Anil Kumar* Physical and Materials Chemistry Division, CSIR−National Chemical Laboratory, Pune 411008, India S Supporting Information *
ABSTRACT: Hammett acidity function observes the effect of protonation/deprotonation on the optical density/absorbance of spectrophotometric indicator. In this work, the Hammett acidity, H0, of equimolar mixtures of aqueous HCOOH, CH3COOH, and CH3CH2COOH was measured in 1methylimidazolium-, 1-methylpyrrolidinium-, and 1-methylpiperidinium-based protic ionic liquids (PILs) and 1-butyl-3-methylimidazolium-based aprotic ionic liquid (AIL) with formate (HCOO−) anion. Higher H0 values were observed for the equimolar mixtures of aqueous carboxylic acids in protic ionic liquids compared with those of the aprotic ionic liquid because of the involvement of the stronger specific interactions between the conjugate acid of ionic liquid and conjugate base of carboxylic acids as suggested by the hard−soft acid base (HSAB) theory. The different H0 values for the equimolar mixtures of aqueous carboxylic acids in protic and aprotic ionic liquids were noted to depend on the activation energy of proton transfer (Ea,H+). The higher activation energy of proton transfer was obtained in AIL, indicating lower ability to form specific interactions with solute than that of PILs. Thermodynamic parameters determined by the “indicator overlapping method” further confirmed the involvement of the secondary interactions in the dissociation of carboxylic acids. On the basis of the thermodynamic parameter values, the potential of different ionic liquids in the dissociation of carboxylic acids was observed to depend on the hydrogen bond donor acidity (α) and hydrogen bond acceptor basicity (β), characteristics of specific interactions.
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INTRODUCTION The concept of acid−base catalysis has played a key role in the understanding and development of many organic and biological processes.1 Precise determination of the acidity level/strength of an acid assists in understanding the mechanistic aspects of these processes. The most common scales employed in the acidity level determination of an acid are pKa, pH, and acidity function.2 The development of the Hammett structure−activity equation and Brønsted catalysis equation, which laid the foundation of physical−organic chemistry, is based on the pKa scale of acidity measurement.3,4 Although the pH and pKa scales measure the strength of an acid in a medium, acidity function has a slight advantage over these scales because of its possible bearing on the acid catalysis. The Hammett acidity function, H0, developed by Hammett and Deyrup, is used to extend the pH limit, over the range of 13 < pH < 1, on either end in aqueous and nonaqueous media.5,6 The strength of an acid in a medium depends on the availability of a proton provided by the acid rather than the concentration of acid as suggested by the solvation enthalpy (ΔH) of the proton.7 Thus, the magnitude of solute−solvent interactions regulates the extent of dissociation of an acid in a medium. Knowledge about the nature of solute−solvent interactions between acid and solvent can therefore be of significant value in the design of efficient media for attaining maximum strength even for weak acids. In this way, involvement of the corrosive acids can be avoided in various processes by employing the efficient media along with the moderate acid. © 2015 American Chemical Society
Thermodynamic parameters have been used to predict the nature of solute−solvent interactions in the dissociation of weak acid.8 Thermal change directly affects the magnitude of solute−solvent interactions. Thus, the thermal effect on a physical change can be used to comprehend the nature of the solute−solvent interactions involved in the process. The thermodynamic parameters are obtained from the measurements of dissociation constant (Ka) weak acids at different temperatures. The Ka values of weak acids are measured either by potentiometry or by the EMF method. The EMF method fails in the case of strong acids, whereas potentiometry does not yield reliable values of the Ka values for weak acids and bases.9 In the current work, we propose a method based on the “indicator overlapping method” for the determination of Ka of weak acids. This method depends on the protonation of indicator at different temperatures. The general criterion for the selection of indicator is that its pKa should be at least 2 units less than that of the acid to avoid complete protonation of the indicator. In view of alarming situations caused by the use of volatile organic compounds in chemical synthesis, intense efforts are being made to use ionic liquids because of their negligible vapor pressure, high thermal stability, low melting points, good solvating capabilities, wide electrochemical window, inflammability, intrinsic ionic conductivities, recyclability, etc.10−13 Received: January 3, 2015 Revised: March 9, 2015 Published: April 3, 2015 5537
DOI: 10.1021/acs.jpcb.5b00056 J. Phys. Chem. B 2015, 119, 5537−5545
Article
The Journal of Physical Chemistry B Because of their solvent−catalyst duality, shorter reaction time, good yields, lower temperature, and high selectivity even for sterically crowded substrates, protic ionic liquids (PILs) have become a suitable choice for various acid-catalyzed reactions over conventional catalyst.14,15 The efficacy of PILs in various applications largely depends on the ionic content/ ionicity.16 Belieres and Angell have observed that the ionicity in PILs relies on the pKa difference (ΔpKa) between acid and base.17 However, MacFarlane et al. have shown that the ΔpKa is a poor estimate of the ionicity in PILs.18−20 The equilibrium study suggests that ΔpKa = 4 is sufficient for complete proton transfer, whereas ΔpKa > 3 was adequate to produce 100% ionicity in ionic solids.21,22 A strong network of hydrogen bonds in PILs prohibits the back-transfer of proton and inhibits the formation of neutral acid and base.23 Recently, a general account of various properties and applications of PILs was presented by Greaves and Drummond.24 The optimum yield and selectivity in PILs can be attributed to the involvement of both acidic and basic moieties, contrary to the traditional catalyst or the combination of aprotic ionic liquid (AIL) and acid.15 The feeble action of AIL and acid together for the catalytic process can attributed to the lower strength of acid in AIL. About a decade ago, Thomazeau et al. measured the strength of bistriflylimide acid (HN(Tf)2) and triflic acid (HOTf) in terms of the H0 of bistriflylimide acid (HN(Tf)2) and triflic acid (HOTf) in imidazolium-based ILs containing tetrafluoroborate ([BF4]−) and bis(trifluoromethyl-1-sulphonyl) imide ([N(Tf)2]−) anions.25 This study observed the insignificant contribution of C2−H (the most acidic proton on the imidazolium cation) on the H0 value. Most recently, a standard pKa scale was developed for substituted benzoic acids in room temperature ILs.26 In a different study, lower pKa values have been reported for the aliphatic and aromatic carboxylic acids in AILs than that in water.27,28 The lower potential of AILs in the dissociation of acid is also observed during the comparison of the relative stability of the acidic/ basic form of several indicators in water at different pH values and in AILs.21 The noteworthy contribution of this study was the greater stability of the basic form of indicators in AILs than in water because of the higher basicity of the former. In recent years, we have been involved in the investigation of several physicochemical and physical organic issues in ILs.29−32 Of late, we have documented the higher acidic strength for weak carboxylic acids in PILs compared to that in water.33 The greater dissociation of carboxylic acids was noted in PILs containing a more acidic cation along with the less basic anion or vice versa, that is, for unstable PILs. An unstable PIL leads to the greater dissociation and consequently higher strength of carboxylic acids. The secondary interactions between the PIL cation and the conjugate anion of acid can be understood in terms of the “hard−soft acid−base” (HSAB) principle.33 In the present investigation, we address the issue of the acidity determination of carboxylic acids in PIL and AIL using the Hammett acidity relation. We have attempted to relate the Hammett acidity change in different ILs in terms of the solute− solvent interactions operating between acid and IL. Efforts have also been made to account for the principal reason for the large difference in the H0 values for carboxylic acids in PIL and AIL. We have attempted to solve these issues by measuring the H0 for carboxylic acids in PIL and AIL, determining thermodynamic parameters, activation energy of proton transfer, and polarity. The structures of ILs employed in the study are shown in Figure 1.
Figure 1. Structures of the ILs employed in the study.
The structures of different probes used for the polarity and Hammett acidity (4-nitroaniline, pKa = 0.99) measurements are shown in Figure 2. The acronyms for ILs used in the current study are summarized in Table 1.
Figure 2. Structures of probes used for the polarity and Hammett acidity measurements.
Table 1. Acronyms of ILs Employed in the Study acronym
name of ionic liquid
[HmIm][HCOO] [HPyrr][HCOO] [HPyrd][HCOO] [BMIm][HCOO]
1-methylimidazolium formate 1-methylpyrrolidinium formate 1-methylpiperidinium formate 1-butyl-3-methylimidazolium formate
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EXPERIMENTAL SECTION Materials. The spectroscopic indicator dye 2,6-dichloro-4(2,4,6-triphenylpyridinium-1-yl)phenolate (1) was purchased from Fluka Analytical (98.5% purity). Reichardt’s dye (2) and 4-nitroaniline (3) were obtained from M/s Sigma-Aldrich. N,NDiethyl-4-nitroaniline (4) was purchased from Frinton Laboratories (99% purity) (Figure 2). 1-Methylimidazole, 1methylpyrrolidine, 1-methylpiperidine, and n-bromobutane were procured from Sigma and distilled prior to use. Formic, acetic, and propionic acids were obtained from Merck and used as obtained. Deionized water possessing a specific conductance of 3. For pyrrolidinium- and piperidinium-based PILs ΔpKa is >6. On the basis of these reports, it can be assumed that the PILs used in this work have high ionic content. Measurement of Acidity Using the Hammett Equation. The Hammett acidity of HCOOH, CH3COOH, and CH3CH2COOH in different ILs was measured according to the procedure used in the literature.33 In principle, the indicator ratio was measured using the indicator overlapping method suggested by Gilbert and group by comparing the reference absorbance with those obtained at different concentrations of carboxylic acids.25 The strength of HCOOH, CH3COOH, and CH3CH2COOH in PILs and AIL is calculated by Hammett equation H0 = pK (I)aq + log[I]s /[HI+]s
a polarity scale of the less basic probe 2,6-dichloro-4-(2,4,6triphenylpyridinium-1-yl)phenolate (pKa = 4.78).38 pH Measurement. The pH of an equimolar mixture of carboxylic acids and water in ILs was measured by using a calibrated pH meter (DPH 504). Determination of Thermodynamic Parameters by Indicator Overlapping Method. A small volume of 4nitroaniline (10−4 M) prepared in methanol was added to a round-bottom flask, and methanol was removed using high vacuum prior to the addition of 1 mL of IL. The sample was filled inside a cuvette, and absorbance was recorded at 298 K. This absorbance was taken as reference value. The reference absorbance denotes the presence of the 100% unprotonated form of indicator and thus negates any aid provided by the medium on the H0. All acids were mixed in equimolar proportion with deionized water. The H0 of carboxylic acids in ILs was independent of the presence of a small amount of water as dissociation is guided by the basicity (β) of medium, and the β of studied ILs (Table 5) is higher than that of water (β = 0.49).21 The calculated amount of acid was added to the cuvette, and absorbance was noted at different temperatures ranging from 298 to 338 K. Similarly, the absorbance at different concentrations in the temperature range from 298 to 338 K was also recorded. The indicator ratio ([HI+]/[I]) was calculated by the procedure given in Table 2. Table 2. Calculation of the Equilibrium Constant for HCOOH in [HPyrr][HCOO] at 0.48 M in the Temperature Range from 298 to 338 K by Indicator Overlapping Method
(1)
where pK(I)aq is the protonation constant of 4-nitroaniline (pK(I)aq = 0.99) in aqueous solution and [I]s and [HI+]s are the concentrations of nonprotonated and protonated forms of the indicator in solvated state. From eq 1, it is clear that a strong acid possesses a lower H0 value. Equation 1 can also be written as H0 = −log a(H+aq) − log γ(I)/γ(HI+) − log Γ(I) /Γ(HI+)
a
(2)
where γ(I) and γ(HI+) are the activity coefficients of the unprotonated and protonated forms of indicator and Γ(I), Γ(HI+) are the transfer activity coefficients of indicator from water to the PIL. In a dilute solution of indicator (10−6 M), the ratio γ(I)/γ(HI+) remains constant. However, the ratio of the transfer activity coefficients of the two forms of indicator depends on their solvation. Because it is difficult to maintain a similar environment for the solvation of protonated and unprotonated forms of indicator, the ratio Γ(I)/Γ(HI+) cannot become unity. For structurally similar indicators, it was assumed that the ratio Γ(I)/Γ(HI+) remains constant. The H0 values of carboxylic acids in ILs are within reproducibility of ±0.05. Polarity Measurement. The polarity of [HPyrr][HCOO], [HPyrd][HCOO], and [BMIM][HCOO] was determined by using the UV−visible spectroscopic method described in the literature.38 The longest wavelength intramolecular charge transfer spectra of different probes dissolved in 1 mL of an ionic liquid was recorded. The polarity of the ionic liquids and Kamlet−Taft parameters were calculated using an earlier work.31 PILs have lower pKa values than AILs owing to the presence of acidic N−H protons. This causes the disappearance of the long-wavelength intramolecular charge transfer (ICT) absorption band for Reichardt’s dye 30 (pKa = 8.6) in PILs. The ET(30) values for all PILs were thus derived from ET(33),
T (K)
Amax
[I] (%)
[HI+ (%)
[HI+]/[I]
Keqa
298 298 308 318 328 338
1.420 1.334 1.301 1.269 1.230 1.190
100.0 93.9 91.6 89.4 86.6 83.8
0 6.1 8.4 10.6 13.4 16.2
0 0.064 0.091 0.119 0.154 0.193
0 0.13 0.19 0.25 0.32 0.40
Keq values are precise to ±0.03.
The protonation equilibrium of indicator is I + HX ↔ HI+ + X−
(3)
The equilibrium constant (Keq) for the dissociation of weak acids is given as Keq =
[HI+][X−] [I][HX]
(4)
+
The ratio [HI ]/[I] was obtained by comparing the optical density/absorbance of the protonated form of indicator against the reference value. The ratio [X−]/[HX] was calculated from the pH data of carboxylic acids in PILs. However, a very small extent of carboxylic acid remains in the dissociated form in PILs as suggested by their degree of dissociation (α) values calculated from the pH measurements. The degree of dissociation for HCOOH, CH3COOH, and CH3CH2COOH in [HmIm][HCOO] and [HPyrr][HCOO] was in the order of ∼10−7, whereas in [HPyrd][HCOO] and [BMIm][HCOO] it was noted to be on the order of ∼10−6 (Tables T1−T4 in the Supporting Information) within the experimental temperature range of 298−338 K. The equilibrium constant was calculated as shown in Table 2. 5539
DOI: 10.1021/acs.jpcb.5b00056 J. Phys. Chem. B 2015, 119, 5537−5545
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The Journal of Physical Chemistry B
regioselectivity of the product in PILs.43 The high polarity of PIL helps in the aggregation of surfactant compared to aprotic ILs.44 Thus, the Hammett acidity, which measures the ability of a solvent to donate a proton to a base, would be different for a protic acid in PILs and AILs. Lower H 0 values for HCOOH, CH 3 COOH, and CH3CH2COOH were obtained in PILs than in AIL (Figure 4). The CH+ values denote the concentrations of added carboxylic acid. This suggests the greater dissociation of carboxylic acids in PIL in comparison with AIL. For comparison, the H0 values of HCOOH, CH3COOH, and CH3CH2COOH in PILs are taken from our earlier work.33 From these plots, it is evident that the H0 values for all carboxylic acids are higher in PILs as compared to those in AIL. The higher dissociation of carboxylic acids in PILs is independent of their hyperpolarity as lower H0 values were obtained in the polar PILs. For example, the maximum dissociation for all carboxylic acids is obtained in [HmIm][HCOO] (ET(30) = 56 kcal mol−1) as compared with [HPyrd][HCOO] (ET(30) = 61.4 kcal mol−1) despite the lower polarity of the former PIL. This indicates involvement of nonelectrostatic interactions between PIL and carboxylic acid. Recently, it was reported that the dissociation of carboxylic acid in PILs is facilitated by the hard−soft acid−base (HSAB) combination between the PIL cation and the conjugate anion of acid.26 The increasing strength of carboxylic acids in PILs follows the order
Similarly, the temperature-dependent Keq values were obtained at different concentrations of carboxylic acid (0.79, 1.07, and 1.4 M) and were further used for the determination of the thermodynamic parameters. The linear dependence of Keq on temperature is given by the van’t Hoff equation: ΔG° = − RT ln Keq
(5)
or ΔG° (6) RT At a given temperature, ΔG° depends on the ΔH° and ΔS° ln Keq = −
as ΔG° = ΔH ° − T ΔS°
(7)
Substituting eq 7 into eq 5, the linear form of the van’t Hoff equation leads to ΔH ° ΔS° + (8) RT R A plot of ln Keq against 1/T gives a straight line with slope = −ΔH°/R and intercept = ΔS°/R. A negative slope between ln Keq and 1/T signifies an endothermic nature of carboxylic acid dissociation, whereas a positive slope implies an exothermic nature of the dissociation process as suggested by eq 8. In PILs and AIL, dissociation of carboxylic acids was endothermic in nature. A typical plot of ln Keq against 1/T for carboxylic acid is shown in Figure 3. The plots of ln Keq versus 1/T for other ln Keq = −
CH3CH 2COOH > CH3COOH > HCOOH
Thus, the conjugate acidity of the PIL cation acts as a decisive factor in controlling the H0 of carboxylic acids. The order of conjugate acidity of 1-methylimidazole (pKa = 6.95), 1methylpyrrolidine (pKa = 10.5), and 1-methylpiperidine (pKa = 10.08) bases, [HmIm]+, [HPyrr]+, and [HPyrd]+, respectively, is [HmIm]+ > [HPyrd]+ > [HPyrr]+
The strength of HCOOH in PILs varies according to the conjugate acidity of the PIL cations. The maximum strength for HCOOH was achieved in [HmIm][HCOO] and the minimum strength in [HPyrr][HCOO]. Similarly, the strength of CH3COOH in PILs was noted to follow the order of the conjugate acidity of cations. However, a very small difference in the H0 values was observed for CH3COOH in [HPyrd][HCOO] and [HPyrr][HCOO]. The small difference in H0 is probably due to the onset of steric repulsion between the larger [HPyrd]+ (six membered) and CH3COO−, which opposes the HSAB interaction and therefore lowers the H0. The strength of CH3CH2COOH in PILs was noted to follow the order
Figure 3. Plot of ln Keq versus 1/T for CH3COOH in (a) [HPyrr][HCOO] and (b) [BMIM][HCOO] at 0.48 M (■), 0.79 M (○), 1.07 M (▲), and 1.4 M (▽).
systems are shown in the Supporting Information (Figures S1− S4). The thermodynamic parameters ΔH° and ΔS° for the dissociation of carboxylic acids were calculated from the slope and intercept values, respectively.
[HmIm]+ > [HPyrr]+ > [HPyrd]+
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The above ordering suggests the dominance of steric repulsion between CH3CH2COO− and [HPyrd]+ over the favorable HSAB interactions. Contrary to the PIL, which causes higher dissociation of carboxylic acids (lower H0 value), lower dissociation was observed in [BMIM][HCOO] for all carboxylic acids (higher H0 value) as evidenced by Figure 4. This suggests the minor contribution of any conducive mechanism involved in the dissociation of carboxylic acid in AIL. The lower strength of acid albeit in the presence of basic anion (HCOO−) indicates the dominant role of IL cation in the dissociation of acid.
RESULTS AND DISCUSSION The solvation behaviors of protic and aprotic ILs is different from one another as indicated by their polarity values.39−41 For the aprotic class of ILs, polarity (ET(30)) ranges from 47 to 59 kcal mol−1, which is similar to those of alcohols; however, the ET(30) for the protic class of ILs varies from 50 to 65 kcal mol−1, which is comparable to that of water.24,42 In the polar Diels−Alder reactions, the electrophilic character of the dienophile was observed to be significantly higher in PILs than in AILs as evidenced by the higher reactivity and 5540
DOI: 10.1021/acs.jpcb.5b00056 J. Phys. Chem. B 2015, 119, 5537−5545
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Figure 4. Plots of H0 versus ln CH+ for (a) HCOOH, (b) CH3COOH, and (c) CH3CH2COOH in [HmIm][HCOO] (black squares), [HPyrr][HCOO] (red circles), [HPyrd][HCOO] (green triangles), and [BMIM][HCOO] (blue triangles).
[BMIM]+ and CH3CH2COO− and hence results in lower H0 for CH3CH2COOH. The above analysis implies that although IL acts as a strong Coulomb medium, specific interaction plays a significant role in governing the solute−solvent interactions. The strength of specific interactions is noted to depend on the activation energy of proton transfer, Ea,H+, which relies on the viscosity of the medium. Activation Energy of Proton Transfer, Ea,H+. The activation energy of proton transfer (Ea,H+) is a measure of the hindrance provided by the medium in the protonation of indicator. Numerically, it is equivalent to the amount of energy required to release the proton from the solvated state. The high viscosity of IL retards the diffusion of reacting species and thereby lowers the protonation of indicator. Thus, Ea,H + increases with the viscosity of the medium. The Ea,H+ was obtained from the temperature-dependent H0 measurements. When the temperature-dependent H0 was fitted in the Arrhenius-type equation, a linear relationship was observed between lnH0 and 1/T with a negative slope (Figure 6). Ea,H+ was calculated from the slope of linear fit
Therefore, it can be presumed that the H0 in ILs is a function of the conjugate acidity of the cation. In other words, higher dissociation of carboxylic acids in ILs is because of the greater stabilization of the conjugate base of carboxylic acids rather than the stabilization of the proton. The order of the H0 for HCOOH, CH3COOH, and CH3CH2COOH in [BMIM][HCOOH] is shown in Figure 5. The H0 values for HCOOH, CH3COOH, and CH3CH2COOH in [BMIM][HCOO] are noted to vary in the order CH3COOH > CH3CH 2COOH > HCOOH
ln H0 = ln(H0)0 − Ea,H+ /RT
(9)
where (H0)0 is the Hammett acidity at room temperature, Ea,H+ is the activation energy of proton transfer, and R is the universal gas constant. Figure 5. Plots of H0 versus ln CH+ for HCOOH (dotted squares), CH3COOH (dotted circles), and CH3CH2COOH (dotted triangles) in [BMIM][HCOO].
The above arrangement of carboxylic acid dissociation in [BMIM][HCOO] reveals neither the complete dominance of the electrostatic interactions nor the specific interactions. The higher H0 for CH3COOH than of HCOOH in [BMIM][HCOO] indicates involvement of the specific interactions in the dissociation. Because [BMIM][HCOO] is synthesized via quaternization of base rather than the transfer of proton, its cation ([BMIM]+) is expected to possess lower ability to bind with the conjugate base of carboxylic acids than that of PIL cation. A higher H 0 for CH 3 COOH ahead of the CH3CH2COOH can be explained due to the steric repulsion between the [BMIM]+ and CH3CH2COO−.33 The steric repulsion opposes the specific interaction between the
Figure 6. Arrhenius-type plots for HCOOH in (a) [HmIm][HCOO] and [BMIM][HCOO] at 0.48 M (black squares), 0.79 M (red circles), 1.07 M (green triangles), and 1.4 M (blue triangles). 5541
DOI: 10.1021/acs.jpcb.5b00056 J. Phys. Chem. B 2015, 119, 5537−5545
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The Journal of Physical Chemistry B The Ea,H+ values for HCOOH in different ILs are given in Table 3. As shown in Table 3, Ea,H+ for HCOOH in all PILs
solvent interactions between carboxylic acids and ILs can also be shown by the thermodynamic parameters (ΔH°, ΔS°). Thermodynamic Parameters and Dissociation of Carboxylic Acids in ILs. Thermodynamic parameters (ΔH°, ΔS°) arise from solute−solvent interactions and help in comprehending the molecular level interactions. The value of the thermodynamic parameters also helps in predicting the feasibility of protonation of the indicator. The driving force for the protonation of indicator in a medium depends on the affinity of solute by medium. The magnitudes of ΔH° and ΔS° were observed to vary with the nature of ILs. The ΔH° and ΔS° values at different concentrations of HCOOH, CH3COOH, and CH3CH2COOH in [HmIm][HCOO], [HPyrr][HCOO], [HPyrd][HCOO], and [BMIM][HCOO] are given in Table 4. A linear relationship was observed between ΔH° and ΔS° against lnCH+ (Figure 7). This indicates that the nature of carboxylic acid−IL interactions remains independent of the concentration of acid. The efficacy of the thermodynamic parameters in the dissociation of carboxylic acids is discussed in different classes of ILs. The standard entropy change, ΔS°, refers to the entropy change during the protonation of the indicator. In PILs, all carboxylic acids have positive ΔS° values. Large ΔS° values for carboxylic acids in PILs are in contradiction with the generalization that “all ionization of uncharged acids in aqueous solution leads standard entropy change, ΔS°, nearly −22 cal K−1 mol−1.”46 A large ΔS° value indicates the higher dissociation of carboxylic acids in PILs than in aqueous
Table 3. Activation Energy (Ea,H+) for HCOOH in PILs and AIL Ea,H+ (kJ mol−1) IL
0.48 M
0.79 M
1.07 M
1.4 M
[HmIm][HCOO] [HPyrr][HCOO] [HPyrd][HCOO] [BMIM][HCOO]
5.37 5.14 5.47 10.49
4.46 4.06 4.09 9.21
3.87 3.44 3.49 5.53
3.80 3.12 3.12 4.75
does not vary significantly, but for AIL it is nearly 2-fold. The high Ea,H+ value for HCOOH in [BMIM][HCOO] denotes the greater hindrance in the diffusion of reacting species that lowers the protonation of indicator and thus lowers the strength for HCOOH (higher H0). Similar observations were noted for CH3COOH and CH3CH2COOH in PILs and AIL. It is noted by several workers that the diffusion-controlled process depends on the viscosity of the medium.32,45 Thus, the high Ea,H+ value in a medium indicates the highly viscous nature of the medium. Generally, AILs possess higher viscosity than PILs. For AILs, viscosity ranges between 60 and 10000 cP (1 cP = 1 mPa·s), whereas for PILs it varies between 6 and 30 cP. Hence, the different H 0 values for HCOOH, CH 3 COOH, and CH3CH2COOH in PILs and AIL are also confirmed by the Ea,H+ in these media. The presence of intermolecular solute−
Table 4. Thermodynamic Parameters for the Dissociation of HCOOH, CH3COOH, and CH3CH2COOH in Ionic Liquids at Different Concentrations IL [HmIm][HCOO]
carboxylic acid
thermodynamic parametera
0.48 M
0.79 M
1.07 M
1.4 M
ΔH° ΔS° ΔH° ΔS° ΔH° ΔS°
47.7 0.13 38.7 0.10 49.0 0.13
33.2 0.09 30.4 0.08 35.9 0.10
24.7 0.07 27.6 0.07 30.3 0.08
20.6 0.06 23.0 0.06 26.9 0.07
ΔH° ΔS° ΔH° ΔS° ΔH° ΔS°
43.3 0.14 43.1 0.14 52.3 0.13
29.7 0.10 31.2 0.11 26.8 0.10
22.5 0.08 23.9 0.09 21.4 0.08
18.2 0.07 20.3 0.08 15.6 0.06
ΔH° ΔS° ΔH° ΔS° ΔH° ΔS°
26.0 0.09 44.3 0.15 38.6 0.13
17.6 0.07 36.8 0.12 26.5 0.09
14.1 0.05 33.8 0.11 20.9 0.08
12.1 0.04 31.3 0.10 20.8 0.07
ΔH° ΔS° ΔH° ΔS° ΔH° ΔS°
46.0 0.12 16.9 0.04 26.4 0.07
37.2 0.09 13.2 0.03 14.0 0.03
19.0 0.04 10.7 0.02 12.3 0.02
14.8 0.03 9.4 0.01 10.8 0.01
HCOOH CH3COOH CH3CH2COOH
[HPyrr][HCOO]
HCOOH CH3COOH CH3CH2COOH
[HPyrd][HCOO]
HCOOH CH3COOH CH3CH2COOH
[BMIM][HCOO]
HCOOH CH3COOH CH3CH2COOH
a
Units for ΔH° in kJ K−1 mol−1 and for ΔH° in kJ mol−1 5542
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represent the solvation characteristics and can therefore be used for delineating the intermolecular interactions with solute. Thermodynamic Parameters and Polarity of ILs. The different arrangements of PILs in the dissociation of HCOOH, CH3COOH, and CH3CH2COOH based on the ΔH° values indicate specific interactions between carboxylic acid and PILs. The solvation ability of PILs is represented by the polarity parameters. The polarity parameters that are used to define the solvation behavior of ILs are the normalized electronic transition energy (ETN), hydrogen bond donor acidity (α), hydrogen bond acceptor basicity (β), and dipolarity/polarizability or polarity index (π*). The ETN value denotes overall polarity exhibited by medium, whereas α, β, and π* represent the hydrogen bond donating ability of the cation, the hydrogen bond accepting ability of the anion, and the electrostatic strength of the medium, respectively. The polarity parameters of different ILs are shown in Table 5. The standard enthalpy
Figure 7. Plots of (a) ΔS° versus lnCH+ and (b) ΔH° versus lnCH+° for HCOOH in [HmIm][HCOO] (■), [HPyrr][HCOO] (○), and [HPyrd][HCOO] (▼).
Table 5. Polarity Parameters of Protic Ionic Liquids at 298.15 K
solutions. However, the contribution of ΔS° in promoting the dissociation of carboxylic acids is very small in comparison to the ΔH°; that is, dissociation of acids is guided by the ΔH° values. The dissociation of carboxylic acids in PILs is mainly governed by the enthalpy change rather than the entropy change. Furthermore, the solute−solvent interactions can be perceived more accurately by the enthalpy change rather than the entropy change and hence are employed in the detection of intermolecular interactions. The enthalpy required in the dissociation of carboxylic acids varies with the nature of PILs as shown in Table 4. The PILs, which assist in the dissociation of carboxylic acids, require the lower ΔH°. Thus, different PILs can be arranged in the order based on their ability to dissociate carboxylic acids. The dissociation of carboxylic acids in PIL is favored either by the greater stabilization of the conjugate anion of acid or by the dissociation of proton. The orders of different PILs in promoting the dissociation of HCOOH, CH3COOH, and CH3CH2COOH based on the ΔH° values are
IL [HmIm][HCOO] [HPyrr][HCOO] [HPyrd][HCOO] [BMIm][HCOO]
b
ETNa
π*
A
B
56.0 49.3 61.4 51.4
1.10 0.99 0.98 1.17
0.81 0.45 1.25 0.46
0.81 0.78 0.88 0.68
The uncertainty in the ETN values as reported was ±0.04. bReference 38.
a
change, ΔH°, which measures solute−solvent interactions, can be correlated with the polarity parameters of ILs. The arrangement of PILs in dissociating the HCOOH and CH3CH2COOH is in the inverse order of the polarity index (π*), whereas for CH3COOH it is in the increasing order of π*. The efficiency of PILs in controlling the ΔH° for HCOOH and CH3CH2COOH is in the opposite order of π*, which suggests the lower relevance of the electrostatic field effect in controlling the dissociation of acid. This implies the involvement of specific interactions in the dissociation of carboxylic acids. The polarity parameters that represent specific interactions are α and β, although direct dependence on either α or β in the dissociation of HCOOH and CH3CH2COOH is not revealed by the arrangements of PILs. However, the dependence of ΔH° on α or β can be explained in an alternate way. It is proposed by several workers that in ILs the hydrogen bond donor ability of the cation (α) and the hydrogen bond acceptor ability of the anion (β) remain in equilibrium with each other.47,48 The resultant of these two determines the behavior of IL. Therefore, a solute will always experience lesser affinity either toward the cation or anion than that predicted by the Kamlet−Taft parameters (α, β, π*). In other words, α is partially reduced by β and vice versa. Thus, in all ILs, a competition exists between the α and β values in the presence of solute. In view of these counterbalanced interactions present in ILs, the arrangement of different PILs in affecting the ΔH° values should be correlated with the relative α and β values. The higher dissociation of carboxylic acids in PILs is due to either the stabilization of proton (H+) or the stabilization of conjugate base of acid. In [HPyrd][HCOO], α is higher than β; thus, the dissociation of HCOOH and CH3CH2COOH is controlled
[HPyrd]+ > [HPyrr]+ > [HmIm]+ [HmIm]+ > [HPyrr]+ > [HPyrd]+ [HPyrd]+ > [HPyrr]+ > [HmIm]+
As apparent from these arrangements of PILs based on the ΔH° values, there is no definite order of PILs in which they interact with carboxylic acids. The potential of PILs based on the ΔH° values is different from those obtained from the H0 values for HCOOH, CH3COOH, and CH3CH2COOH at 298 K. This discrepancy in the dissociating carboxylic acids in PILs might be due to the thermal effects that were not accounted for in the determination of the H0. However, unexpectedly very low ΔH° values for all carboxylic acids were obtained in [BMIm][HCOO]. This points out that the nature of intermolecular interactions between carboxylic acids and [BMIM][HCOO] is similar to PILs and carboxylic acids based on the ΔH° values. The unusual behavior of [BMIm][HCOO] is in contradiction with the very large H0 values for all carboxylic acids. Thus, the intermolecular force between the [BMIM][HCOO] and carboxylic acids cannot be predicted from the ΔH° values. The ambiguous behavior of [BMIm][HCOO] can be explained by employing the polarity parameters, which control the solute−solvent interactions. The polarity parameters of ILs 5543
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by α. In [HPyrr][HCOO], α is lower than β; hence, dissociation of HCOOH and CH3CH2COOH is favored by β. In [HmIm][HCOO] both α and β are equal in magnitude; therefore, the least dissociation of carboxylic acids and consequently the higher ΔH° for HCOOH and CH3CH2COOH are observed in [HmIm][HCOO]. Hence, the arrangement of different PILs based on the ΔH° for HCOOH and CH3CH2COOH is according to their ability to bind with carboxylic acids by specific interactions. Surprisingly, a different order of PILs in the increasing order of their π* values was noted to be effective in the dissociation of CH3COOH. This indicates the involvement of Coulombic forces in the dissociation of CH3COOH, which seems highly improbable. The arrangement of PILs for CH3COOH in the increasing order of π* value might be due to the higher dipole moment of CH3COOH (1.74 D) than those of HCOOH (1.41 D) and CH3CH2COOH (0.63 D). The strong bond dipoles present in the CH3COOH combine well with the Coulombic region created inside PILs. However, similar observations could not made for the AILs as, except for 1-butyl-3-methylimidazolium formate ([BMIm][HCOO]), other ILs, for example, 1-butyl-3-methylpyrrolidinium formate ([BMPyrr][HCOO]) and 1-butyl-3-methylpiperidine formate ([BMPyrd][HCOO]), were solid at room temperature. However, as can be seen from the polarity parameters for [BMIm][HCOO], α and β are small as compared to those of PILs. A small difference in the α and β indicates the lower tendency of [BMIm][HCOO] toward the specific interactions with solute than of PILs. Therefore, the polarity parameters of [BMIm][HCOO] also support the lower strength of carboxylic acids. Further investigations are required to explore the role of different protic and aprotic ILs having similar cations/anions in the dissociation of weak acids.
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Article
ASSOCIATED CONTENT
S Supporting Information *
Additional tables and drawings associated with the determination of thermodynamic parameters; degree of dissociation (α) and ln Keq versus 1/T for HCOOH, CH3COOH, and CH3CH2COOH in ILs. This material is available free of charge via the Internet at http://pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Author
*(A.K.) E-mail: [email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS S.K.S. is grateful to UGC, New Delhi, for awarding him a research fellowship. A.K. thanks DST, New Delhi, for a J. C. Bose National Fellowship (SR/S2/JCB-26/2009).
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REFERENCES
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CONCLUSIONS
We have compared the potential of protic and aprotic ionic liquids in the dissociation of HCOOH, CH3COOH, and CH3CH2COOH measured by the Hammett acidity function. Protic ionic liquids were observed as promising candidates in the dissociation of carboxylic acids because of higher acidic strength than their aprotic counterparts. The greater dissociation of carboxylic acids in protic ionic liquids was noted to depend on the intensity of specific interactions between them. The activation energy of proton transfer (Ea,H+) further fortifies this view. Higher Ea,H+ values for carboxylic acids obtained in aprotic ionic liquid than in protic ionic liquid lower the possibility of specific interactions. Thermodynamic parameters (ΔH°, ΔS°), calculated by using a novel “indicator overlapping method”, also signify the participation of specific solute− solvent interactions in the dissociation of carboxylic acids. The ability of different protic ionic liquids as observed by ΔH° value in dissociating carboxylic acids was observed to depend on the relative hydrogen bond donor (α) and hydrogen bond acceptor (β) tendencies. The observations drawn above will be useful in designing efficient media for weak acids. This will help in avoiding the use of corrosive acids for various applications. To the best of our knowledge, this is the first report comparing the potential of protic ionic liquid and aprotic ionic liquid in the dissociation of carboxylic acids. 5544
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Dissociation of Equimolar Mixtures of Aqueous Carboxylic Acids in Ionic Liquids: Role of Specific Interactions Shashi Kant Shukla and Anil Kumar* Physical and Materials Chemistry Division, CSIR−National Chemical Laboratory, Pune 411008, India S Supporting Information *
ABSTRACT: Hammett acidity function observes the effect of protonation/deprotonation on the optical density/absorbance of spectrophotometric indicator. In this work, the Hammett acidity, H0, of equimolar mixtures of aqueous HCOOH, CH3COOH, and CH3CH2COOH was measured in 1methylimidazolium-, 1-methylpyrrolidinium-, and 1-methylpiperidinium-based protic ionic liquids (PILs) and 1-butyl-3-methylimidazolium-based aprotic ionic liquid (AIL) with formate (HCOO−) anion. Higher H0 values were observed for the equimolar mixtures of aqueous carboxylic acids in protic ionic liquids compared with those of the aprotic ionic liquid because of the involvement of the stronger specific interactions between the conjugate acid of ionic liquid and conjugate base of carboxylic acids as suggested by the hard−soft acid base (HSAB) theory. The different H0 values for the equimolar mixtures of aqueous carboxylic acids in protic and aprotic ionic liquids were noted to depend on the activation energy of proton transfer (Ea,H+). The higher activation energy of proton transfer was obtained in AIL, indicating lower ability to form specific interactions with solute than that of PILs. Thermodynamic parameters determined by the “indicator overlapping method” further confirmed the involvement of the secondary interactions in the dissociation of carboxylic acids. On the basis of the thermodynamic parameter values, the potential of different ionic liquids in the dissociation of carboxylic acids was observed to depend on the hydrogen bond donor acidity (α) and hydrogen bond acceptor basicity (β), characteristics of specific interactions.
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INTRODUCTION The concept of acid−base catalysis has played a key role in the understanding and development of many organic and biological processes.1 Precise determination of the acidity level/strength of an acid assists in understanding the mechanistic aspects of these processes. The most common scales employed in the acidity level determination of an acid are pKa, pH, and acidity function.2 The development of the Hammett structure−activity equation and Brønsted catalysis equation, which laid the foundation of physical−organic chemistry, is based on the pKa scale of acidity measurement.3,4 Although the pH and pKa scales measure the strength of an acid in a medium, acidity function has a slight advantage over these scales because of its possible bearing on the acid catalysis. The Hammett acidity function, H0, developed by Hammett and Deyrup, is used to extend the pH limit, over the range of 13 < pH < 1, on either end in aqueous and nonaqueous media.5,6 The strength of an acid in a medium depends on the availability of a proton provided by the acid rather than the concentration of acid as suggested by the solvation enthalpy (ΔH) of the proton.7 Thus, the magnitude of solute−solvent interactions regulates the extent of dissociation of an acid in a medium. Knowledge about the nature of solute−solvent interactions between acid and solvent can therefore be of significant value in the design of efficient media for attaining maximum strength even for weak acids. In this way, involvement of the corrosive acids can be avoided in various processes by employing the efficient media along with the moderate acid. © 2015 American Chemical Society
Thermodynamic parameters have been used to predict the nature of solute−solvent interactions in the dissociation of weak acid.8 Thermal change directly affects the magnitude of solute−solvent interactions. Thus, the thermal effect on a physical change can be used to comprehend the nature of the solute−solvent interactions involved in the process. The thermodynamic parameters are obtained from the measurements of dissociation constant (Ka) weak acids at different temperatures. The Ka values of weak acids are measured either by potentiometry or by the EMF method. The EMF method fails in the case of strong acids, whereas potentiometry does not yield reliable values of the Ka values for weak acids and bases.9 In the current work, we propose a method based on the “indicator overlapping method” for the determination of Ka of weak acids. This method depends on the protonation of indicator at different temperatures. The general criterion for the selection of indicator is that its pKa should be at least 2 units less than that of the acid to avoid complete protonation of the indicator. In view of alarming situations caused by the use of volatile organic compounds in chemical synthesis, intense efforts are being made to use ionic liquids because of their negligible vapor pressure, high thermal stability, low melting points, good solvating capabilities, wide electrochemical window, inflammability, intrinsic ionic conductivities, recyclability, etc.10−13 Received: January 3, 2015 Revised: March 9, 2015 Published: April 3, 2015 5537
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Article
The Journal of Physical Chemistry B Because of their solvent−catalyst duality, shorter reaction time, good yields, lower temperature, and high selectivity even for sterically crowded substrates, protic ionic liquids (PILs) have become a suitable choice for various acid-catalyzed reactions over conventional catalyst.14,15 The efficacy of PILs in various applications largely depends on the ionic content/ ionicity.16 Belieres and Angell have observed that the ionicity in PILs relies on the pKa difference (ΔpKa) between acid and base.17 However, MacFarlane et al. have shown that the ΔpKa is a poor estimate of the ionicity in PILs.18−20 The equilibrium study suggests that ΔpKa = 4 is sufficient for complete proton transfer, whereas ΔpKa > 3 was adequate to produce 100% ionicity in ionic solids.21,22 A strong network of hydrogen bonds in PILs prohibits the back-transfer of proton and inhibits the formation of neutral acid and base.23 Recently, a general account of various properties and applications of PILs was presented by Greaves and Drummond.24 The optimum yield and selectivity in PILs can be attributed to the involvement of both acidic and basic moieties, contrary to the traditional catalyst or the combination of aprotic ionic liquid (AIL) and acid.15 The feeble action of AIL and acid together for the catalytic process can attributed to the lower strength of acid in AIL. About a decade ago, Thomazeau et al. measured the strength of bistriflylimide acid (HN(Tf)2) and triflic acid (HOTf) in terms of the H0 of bistriflylimide acid (HN(Tf)2) and triflic acid (HOTf) in imidazolium-based ILs containing tetrafluoroborate ([BF4]−) and bis(trifluoromethyl-1-sulphonyl) imide ([N(Tf)2]−) anions.25 This study observed the insignificant contribution of C2−H (the most acidic proton on the imidazolium cation) on the H0 value. Most recently, a standard pKa scale was developed for substituted benzoic acids in room temperature ILs.26 In a different study, lower pKa values have been reported for the aliphatic and aromatic carboxylic acids in AILs than that in water.27,28 The lower potential of AILs in the dissociation of acid is also observed during the comparison of the relative stability of the acidic/ basic form of several indicators in water at different pH values and in AILs.21 The noteworthy contribution of this study was the greater stability of the basic form of indicators in AILs than in water because of the higher basicity of the former. In recent years, we have been involved in the investigation of several physicochemical and physical organic issues in ILs.29−32 Of late, we have documented the higher acidic strength for weak carboxylic acids in PILs compared to that in water.33 The greater dissociation of carboxylic acids was noted in PILs containing a more acidic cation along with the less basic anion or vice versa, that is, for unstable PILs. An unstable PIL leads to the greater dissociation and consequently higher strength of carboxylic acids. The secondary interactions between the PIL cation and the conjugate anion of acid can be understood in terms of the “hard−soft acid−base” (HSAB) principle.33 In the present investigation, we address the issue of the acidity determination of carboxylic acids in PIL and AIL using the Hammett acidity relation. We have attempted to relate the Hammett acidity change in different ILs in terms of the solute− solvent interactions operating between acid and IL. Efforts have also been made to account for the principal reason for the large difference in the H0 values for carboxylic acids in PIL and AIL. We have attempted to solve these issues by measuring the H0 for carboxylic acids in PIL and AIL, determining thermodynamic parameters, activation energy of proton transfer, and polarity. The structures of ILs employed in the study are shown in Figure 1.
Figure 1. Structures of the ILs employed in the study.
The structures of different probes used for the polarity and Hammett acidity (4-nitroaniline, pKa = 0.99) measurements are shown in Figure 2. The acronyms for ILs used in the current study are summarized in Table 1.
Figure 2. Structures of probes used for the polarity and Hammett acidity measurements.
Table 1. Acronyms of ILs Employed in the Study acronym
name of ionic liquid
[HmIm][HCOO] [HPyrr][HCOO] [HPyrd][HCOO] [BMIm][HCOO]
1-methylimidazolium formate 1-methylpyrrolidinium formate 1-methylpiperidinium formate 1-butyl-3-methylimidazolium formate
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EXPERIMENTAL SECTION Materials. The spectroscopic indicator dye 2,6-dichloro-4(2,4,6-triphenylpyridinium-1-yl)phenolate (1) was purchased from Fluka Analytical (98.5% purity). Reichardt’s dye (2) and 4-nitroaniline (3) were obtained from M/s Sigma-Aldrich. N,NDiethyl-4-nitroaniline (4) was purchased from Frinton Laboratories (99% purity) (Figure 2). 1-Methylimidazole, 1methylpyrrolidine, 1-methylpiperidine, and n-bromobutane were procured from Sigma and distilled prior to use. Formic, acetic, and propionic acids were obtained from Merck and used as obtained. Deionized water possessing a specific conductance of 3. For pyrrolidinium- and piperidinium-based PILs ΔpKa is >6. On the basis of these reports, it can be assumed that the PILs used in this work have high ionic content. Measurement of Acidity Using the Hammett Equation. The Hammett acidity of HCOOH, CH3COOH, and CH3CH2COOH in different ILs was measured according to the procedure used in the literature.33 In principle, the indicator ratio was measured using the indicator overlapping method suggested by Gilbert and group by comparing the reference absorbance with those obtained at different concentrations of carboxylic acids.25 The strength of HCOOH, CH3COOH, and CH3CH2COOH in PILs and AIL is calculated by Hammett equation H0 = pK (I)aq + log[I]s /[HI+]s
a polarity scale of the less basic probe 2,6-dichloro-4-(2,4,6triphenylpyridinium-1-yl)phenolate (pKa = 4.78).38 pH Measurement. The pH of an equimolar mixture of carboxylic acids and water in ILs was measured by using a calibrated pH meter (DPH 504). Determination of Thermodynamic Parameters by Indicator Overlapping Method. A small volume of 4nitroaniline (10−4 M) prepared in methanol was added to a round-bottom flask, and methanol was removed using high vacuum prior to the addition of 1 mL of IL. The sample was filled inside a cuvette, and absorbance was recorded at 298 K. This absorbance was taken as reference value. The reference absorbance denotes the presence of the 100% unprotonated form of indicator and thus negates any aid provided by the medium on the H0. All acids were mixed in equimolar proportion with deionized water. The H0 of carboxylic acids in ILs was independent of the presence of a small amount of water as dissociation is guided by the basicity (β) of medium, and the β of studied ILs (Table 5) is higher than that of water (β = 0.49).21 The calculated amount of acid was added to the cuvette, and absorbance was noted at different temperatures ranging from 298 to 338 K. Similarly, the absorbance at different concentrations in the temperature range from 298 to 338 K was also recorded. The indicator ratio ([HI+]/[I]) was calculated by the procedure given in Table 2. Table 2. Calculation of the Equilibrium Constant for HCOOH in [HPyrr][HCOO] at 0.48 M in the Temperature Range from 298 to 338 K by Indicator Overlapping Method
(1)
where pK(I)aq is the protonation constant of 4-nitroaniline (pK(I)aq = 0.99) in aqueous solution and [I]s and [HI+]s are the concentrations of nonprotonated and protonated forms of the indicator in solvated state. From eq 1, it is clear that a strong acid possesses a lower H0 value. Equation 1 can also be written as H0 = −log a(H+aq) − log γ(I)/γ(HI+) − log Γ(I) /Γ(HI+)
a
(2)
where γ(I) and γ(HI+) are the activity coefficients of the unprotonated and protonated forms of indicator and Γ(I), Γ(HI+) are the transfer activity coefficients of indicator from water to the PIL. In a dilute solution of indicator (10−6 M), the ratio γ(I)/γ(HI+) remains constant. However, the ratio of the transfer activity coefficients of the two forms of indicator depends on their solvation. Because it is difficult to maintain a similar environment for the solvation of protonated and unprotonated forms of indicator, the ratio Γ(I)/Γ(HI+) cannot become unity. For structurally similar indicators, it was assumed that the ratio Γ(I)/Γ(HI+) remains constant. The H0 values of carboxylic acids in ILs are within reproducibility of ±0.05. Polarity Measurement. The polarity of [HPyrr][HCOO], [HPyrd][HCOO], and [BMIM][HCOO] was determined by using the UV−visible spectroscopic method described in the literature.38 The longest wavelength intramolecular charge transfer spectra of different probes dissolved in 1 mL of an ionic liquid was recorded. The polarity of the ionic liquids and Kamlet−Taft parameters were calculated using an earlier work.31 PILs have lower pKa values than AILs owing to the presence of acidic N−H protons. This causes the disappearance of the long-wavelength intramolecular charge transfer (ICT) absorption band for Reichardt’s dye 30 (pKa = 8.6) in PILs. The ET(30) values for all PILs were thus derived from ET(33),
T (K)
Amax
[I] (%)
[HI+ (%)
[HI+]/[I]
Keqa
298 298 308 318 328 338
1.420 1.334 1.301 1.269 1.230 1.190
100.0 93.9 91.6 89.4 86.6 83.8
0 6.1 8.4 10.6 13.4 16.2
0 0.064 0.091 0.119 0.154 0.193
0 0.13 0.19 0.25 0.32 0.40
Keq values are precise to ±0.03.
The protonation equilibrium of indicator is I + HX ↔ HI+ + X−
(3)
The equilibrium constant (Keq) for the dissociation of weak acids is given as Keq =
[HI+][X−] [I][HX]
(4)
+
The ratio [HI ]/[I] was obtained by comparing the optical density/absorbance of the protonated form of indicator against the reference value. The ratio [X−]/[HX] was calculated from the pH data of carboxylic acids in PILs. However, a very small extent of carboxylic acid remains in the dissociated form in PILs as suggested by their degree of dissociation (α) values calculated from the pH measurements. The degree of dissociation for HCOOH, CH3COOH, and CH3CH2COOH in [HmIm][HCOO] and [HPyrr][HCOO] was in the order of ∼10−7, whereas in [HPyrd][HCOO] and [BMIm][HCOO] it was noted to be on the order of ∼10−6 (Tables T1−T4 in the Supporting Information) within the experimental temperature range of 298−338 K. The equilibrium constant was calculated as shown in Table 2. 5539
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regioselectivity of the product in PILs.43 The high polarity of PIL helps in the aggregation of surfactant compared to aprotic ILs.44 Thus, the Hammett acidity, which measures the ability of a solvent to donate a proton to a base, would be different for a protic acid in PILs and AILs. Lower H 0 values for HCOOH, CH 3 COOH, and CH3CH2COOH were obtained in PILs than in AIL (Figure 4). The CH+ values denote the concentrations of added carboxylic acid. This suggests the greater dissociation of carboxylic acids in PIL in comparison with AIL. For comparison, the H0 values of HCOOH, CH3COOH, and CH3CH2COOH in PILs are taken from our earlier work.33 From these plots, it is evident that the H0 values for all carboxylic acids are higher in PILs as compared to those in AIL. The higher dissociation of carboxylic acids in PILs is independent of their hyperpolarity as lower H0 values were obtained in the polar PILs. For example, the maximum dissociation for all carboxylic acids is obtained in [HmIm][HCOO] (ET(30) = 56 kcal mol−1) as compared with [HPyrd][HCOO] (ET(30) = 61.4 kcal mol−1) despite the lower polarity of the former PIL. This indicates involvement of nonelectrostatic interactions between PIL and carboxylic acid. Recently, it was reported that the dissociation of carboxylic acid in PILs is facilitated by the hard−soft acid−base (HSAB) combination between the PIL cation and the conjugate anion of acid.26 The increasing strength of carboxylic acids in PILs follows the order
Similarly, the temperature-dependent Keq values were obtained at different concentrations of carboxylic acid (0.79, 1.07, and 1.4 M) and were further used for the determination of the thermodynamic parameters. The linear dependence of Keq on temperature is given by the van’t Hoff equation: ΔG° = − RT ln Keq
(5)
or ΔG° (6) RT At a given temperature, ΔG° depends on the ΔH° and ΔS° ln Keq = −
as ΔG° = ΔH ° − T ΔS°
(7)
Substituting eq 7 into eq 5, the linear form of the van’t Hoff equation leads to ΔH ° ΔS° + (8) RT R A plot of ln Keq against 1/T gives a straight line with slope = −ΔH°/R and intercept = ΔS°/R. A negative slope between ln Keq and 1/T signifies an endothermic nature of carboxylic acid dissociation, whereas a positive slope implies an exothermic nature of the dissociation process as suggested by eq 8. In PILs and AIL, dissociation of carboxylic acids was endothermic in nature. A typical plot of ln Keq against 1/T for carboxylic acid is shown in Figure 3. The plots of ln Keq versus 1/T for other ln Keq = −
CH3CH 2COOH > CH3COOH > HCOOH
Thus, the conjugate acidity of the PIL cation acts as a decisive factor in controlling the H0 of carboxylic acids. The order of conjugate acidity of 1-methylimidazole (pKa = 6.95), 1methylpyrrolidine (pKa = 10.5), and 1-methylpiperidine (pKa = 10.08) bases, [HmIm]+, [HPyrr]+, and [HPyrd]+, respectively, is [HmIm]+ > [HPyrd]+ > [HPyrr]+
The strength of HCOOH in PILs varies according to the conjugate acidity of the PIL cations. The maximum strength for HCOOH was achieved in [HmIm][HCOO] and the minimum strength in [HPyrr][HCOO]. Similarly, the strength of CH3COOH in PILs was noted to follow the order of the conjugate acidity of cations. However, a very small difference in the H0 values was observed for CH3COOH in [HPyrd][HCOO] and [HPyrr][HCOO]. The small difference in H0 is probably due to the onset of steric repulsion between the larger [HPyrd]+ (six membered) and CH3COO−, which opposes the HSAB interaction and therefore lowers the H0. The strength of CH3CH2COOH in PILs was noted to follow the order
Figure 3. Plot of ln Keq versus 1/T for CH3COOH in (a) [HPyrr][HCOO] and (b) [BMIM][HCOO] at 0.48 M (■), 0.79 M (○), 1.07 M (▲), and 1.4 M (▽).
systems are shown in the Supporting Information (Figures S1− S4). The thermodynamic parameters ΔH° and ΔS° for the dissociation of carboxylic acids were calculated from the slope and intercept values, respectively.
[HmIm]+ > [HPyrr]+ > [HPyrd]+
■
The above ordering suggests the dominance of steric repulsion between CH3CH2COO− and [HPyrd]+ over the favorable HSAB interactions. Contrary to the PIL, which causes higher dissociation of carboxylic acids (lower H0 value), lower dissociation was observed in [BMIM][HCOO] for all carboxylic acids (higher H0 value) as evidenced by Figure 4. This suggests the minor contribution of any conducive mechanism involved in the dissociation of carboxylic acid in AIL. The lower strength of acid albeit in the presence of basic anion (HCOO−) indicates the dominant role of IL cation in the dissociation of acid.
RESULTS AND DISCUSSION The solvation behaviors of protic and aprotic ILs is different from one another as indicated by their polarity values.39−41 For the aprotic class of ILs, polarity (ET(30)) ranges from 47 to 59 kcal mol−1, which is similar to those of alcohols; however, the ET(30) for the protic class of ILs varies from 50 to 65 kcal mol−1, which is comparable to that of water.24,42 In the polar Diels−Alder reactions, the electrophilic character of the dienophile was observed to be significantly higher in PILs than in AILs as evidenced by the higher reactivity and 5540
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Figure 4. Plots of H0 versus ln CH+ for (a) HCOOH, (b) CH3COOH, and (c) CH3CH2COOH in [HmIm][HCOO] (black squares), [HPyrr][HCOO] (red circles), [HPyrd][HCOO] (green triangles), and [BMIM][HCOO] (blue triangles).
[BMIM]+ and CH3CH2COO− and hence results in lower H0 for CH3CH2COOH. The above analysis implies that although IL acts as a strong Coulomb medium, specific interaction plays a significant role in governing the solute−solvent interactions. The strength of specific interactions is noted to depend on the activation energy of proton transfer, Ea,H+, which relies on the viscosity of the medium. Activation Energy of Proton Transfer, Ea,H+. The activation energy of proton transfer (Ea,H+) is a measure of the hindrance provided by the medium in the protonation of indicator. Numerically, it is equivalent to the amount of energy required to release the proton from the solvated state. The high viscosity of IL retards the diffusion of reacting species and thereby lowers the protonation of indicator. Thus, Ea,H + increases with the viscosity of the medium. The Ea,H+ was obtained from the temperature-dependent H0 measurements. When the temperature-dependent H0 was fitted in the Arrhenius-type equation, a linear relationship was observed between lnH0 and 1/T with a negative slope (Figure 6). Ea,H+ was calculated from the slope of linear fit
Therefore, it can be presumed that the H0 in ILs is a function of the conjugate acidity of the cation. In other words, higher dissociation of carboxylic acids in ILs is because of the greater stabilization of the conjugate base of carboxylic acids rather than the stabilization of the proton. The order of the H0 for HCOOH, CH3COOH, and CH3CH2COOH in [BMIM][HCOOH] is shown in Figure 5. The H0 values for HCOOH, CH3COOH, and CH3CH2COOH in [BMIM][HCOO] are noted to vary in the order CH3COOH > CH3CH 2COOH > HCOOH
ln H0 = ln(H0)0 − Ea,H+ /RT
(9)
where (H0)0 is the Hammett acidity at room temperature, Ea,H+ is the activation energy of proton transfer, and R is the universal gas constant. Figure 5. Plots of H0 versus ln CH+ for HCOOH (dotted squares), CH3COOH (dotted circles), and CH3CH2COOH (dotted triangles) in [BMIM][HCOO].
The above arrangement of carboxylic acid dissociation in [BMIM][HCOO] reveals neither the complete dominance of the electrostatic interactions nor the specific interactions. The higher H0 for CH3COOH than of HCOOH in [BMIM][HCOO] indicates involvement of the specific interactions in the dissociation. Because [BMIM][HCOO] is synthesized via quaternization of base rather than the transfer of proton, its cation ([BMIM]+) is expected to possess lower ability to bind with the conjugate base of carboxylic acids than that of PIL cation. A higher H 0 for CH 3 COOH ahead of the CH3CH2COOH can be explained due to the steric repulsion between the [BMIM]+ and CH3CH2COO−.33 The steric repulsion opposes the specific interaction between the
Figure 6. Arrhenius-type plots for HCOOH in (a) [HmIm][HCOO] and [BMIM][HCOO] at 0.48 M (black squares), 0.79 M (red circles), 1.07 M (green triangles), and 1.4 M (blue triangles). 5541
DOI: 10.1021/acs.jpcb.5b00056 J. Phys. Chem. B 2015, 119, 5537−5545
Article
The Journal of Physical Chemistry B The Ea,H+ values for HCOOH in different ILs are given in Table 3. As shown in Table 3, Ea,H+ for HCOOH in all PILs
solvent interactions between carboxylic acids and ILs can also be shown by the thermodynamic parameters (ΔH°, ΔS°). Thermodynamic Parameters and Dissociation of Carboxylic Acids in ILs. Thermodynamic parameters (ΔH°, ΔS°) arise from solute−solvent interactions and help in comprehending the molecular level interactions. The value of the thermodynamic parameters also helps in predicting the feasibility of protonation of the indicator. The driving force for the protonation of indicator in a medium depends on the affinity of solute by medium. The magnitudes of ΔH° and ΔS° were observed to vary with the nature of ILs. The ΔH° and ΔS° values at different concentrations of HCOOH, CH3COOH, and CH3CH2COOH in [HmIm][HCOO], [HPyrr][HCOO], [HPyrd][HCOO], and [BMIM][HCOO] are given in Table 4. A linear relationship was observed between ΔH° and ΔS° against lnCH+ (Figure 7). This indicates that the nature of carboxylic acid−IL interactions remains independent of the concentration of acid. The efficacy of the thermodynamic parameters in the dissociation of carboxylic acids is discussed in different classes of ILs. The standard entropy change, ΔS°, refers to the entropy change during the protonation of the indicator. In PILs, all carboxylic acids have positive ΔS° values. Large ΔS° values for carboxylic acids in PILs are in contradiction with the generalization that “all ionization of uncharged acids in aqueous solution leads standard entropy change, ΔS°, nearly −22 cal K−1 mol−1.”46 A large ΔS° value indicates the higher dissociation of carboxylic acids in PILs than in aqueous
Table 3. Activation Energy (Ea,H+) for HCOOH in PILs and AIL Ea,H+ (kJ mol−1) IL
0.48 M
0.79 M
1.07 M
1.4 M
[HmIm][HCOO] [HPyrr][HCOO] [HPyrd][HCOO] [BMIM][HCOO]
5.37 5.14 5.47 10.49
4.46 4.06 4.09 9.21
3.87 3.44 3.49 5.53
3.80 3.12 3.12 4.75
does not vary significantly, but for AIL it is nearly 2-fold. The high Ea,H+ value for HCOOH in [BMIM][HCOO] denotes the greater hindrance in the diffusion of reacting species that lowers the protonation of indicator and thus lowers the strength for HCOOH (higher H0). Similar observations were noted for CH3COOH and CH3CH2COOH in PILs and AIL. It is noted by several workers that the diffusion-controlled process depends on the viscosity of the medium.32,45 Thus, the high Ea,H+ value in a medium indicates the highly viscous nature of the medium. Generally, AILs possess higher viscosity than PILs. For AILs, viscosity ranges between 60 and 10000 cP (1 cP = 1 mPa·s), whereas for PILs it varies between 6 and 30 cP. Hence, the different H 0 values for HCOOH, CH 3 COOH, and CH3CH2COOH in PILs and AIL are also confirmed by the Ea,H+ in these media. The presence of intermolecular solute−
Table 4. Thermodynamic Parameters for the Dissociation of HCOOH, CH3COOH, and CH3CH2COOH in Ionic Liquids at Different Concentrations IL [HmIm][HCOO]
carboxylic acid
thermodynamic parametera
0.48 M
0.79 M
1.07 M
1.4 M
ΔH° ΔS° ΔH° ΔS° ΔH° ΔS°
47.7 0.13 38.7 0.10 49.0 0.13
33.2 0.09 30.4 0.08 35.9 0.10
24.7 0.07 27.6 0.07 30.3 0.08
20.6 0.06 23.0 0.06 26.9 0.07
ΔH° ΔS° ΔH° ΔS° ΔH° ΔS°
43.3 0.14 43.1 0.14 52.3 0.13
29.7 0.10 31.2 0.11 26.8 0.10
22.5 0.08 23.9 0.09 21.4 0.08
18.2 0.07 20.3 0.08 15.6 0.06
ΔH° ΔS° ΔH° ΔS° ΔH° ΔS°
26.0 0.09 44.3 0.15 38.6 0.13
17.6 0.07 36.8 0.12 26.5 0.09
14.1 0.05 33.8 0.11 20.9 0.08
12.1 0.04 31.3 0.10 20.8 0.07
ΔH° ΔS° ΔH° ΔS° ΔH° ΔS°
46.0 0.12 16.9 0.04 26.4 0.07
37.2 0.09 13.2 0.03 14.0 0.03
19.0 0.04 10.7 0.02 12.3 0.02
14.8 0.03 9.4 0.01 10.8 0.01
HCOOH CH3COOH CH3CH2COOH
[HPyrr][HCOO]
HCOOH CH3COOH CH3CH2COOH
[HPyrd][HCOO]
HCOOH CH3COOH CH3CH2COOH
[BMIM][HCOO]
HCOOH CH3COOH CH3CH2COOH
a
Units for ΔH° in kJ K−1 mol−1 and for ΔH° in kJ mol−1 5542
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represent the solvation characteristics and can therefore be used for delineating the intermolecular interactions with solute. Thermodynamic Parameters and Polarity of ILs. The different arrangements of PILs in the dissociation of HCOOH, CH3COOH, and CH3CH2COOH based on the ΔH° values indicate specific interactions between carboxylic acid and PILs. The solvation ability of PILs is represented by the polarity parameters. The polarity parameters that are used to define the solvation behavior of ILs are the normalized electronic transition energy (ETN), hydrogen bond donor acidity (α), hydrogen bond acceptor basicity (β), and dipolarity/polarizability or polarity index (π*). The ETN value denotes overall polarity exhibited by medium, whereas α, β, and π* represent the hydrogen bond donating ability of the cation, the hydrogen bond accepting ability of the anion, and the electrostatic strength of the medium, respectively. The polarity parameters of different ILs are shown in Table 5. The standard enthalpy
Figure 7. Plots of (a) ΔS° versus lnCH+ and (b) ΔH° versus lnCH+° for HCOOH in [HmIm][HCOO] (■), [HPyrr][HCOO] (○), and [HPyrd][HCOO] (▼).
Table 5. Polarity Parameters of Protic Ionic Liquids at 298.15 K
solutions. However, the contribution of ΔS° in promoting the dissociation of carboxylic acids is very small in comparison to the ΔH°; that is, dissociation of acids is guided by the ΔH° values. The dissociation of carboxylic acids in PILs is mainly governed by the enthalpy change rather than the entropy change. Furthermore, the solute−solvent interactions can be perceived more accurately by the enthalpy change rather than the entropy change and hence are employed in the detection of intermolecular interactions. The enthalpy required in the dissociation of carboxylic acids varies with the nature of PILs as shown in Table 4. The PILs, which assist in the dissociation of carboxylic acids, require the lower ΔH°. Thus, different PILs can be arranged in the order based on their ability to dissociate carboxylic acids. The dissociation of carboxylic acids in PIL is favored either by the greater stabilization of the conjugate anion of acid or by the dissociation of proton. The orders of different PILs in promoting the dissociation of HCOOH, CH3COOH, and CH3CH2COOH based on the ΔH° values are
IL [HmIm][HCOO] [HPyrr][HCOO] [HPyrd][HCOO] [BMIm][HCOO]
b
ETNa
π*
A
B
56.0 49.3 61.4 51.4
1.10 0.99 0.98 1.17
0.81 0.45 1.25 0.46
0.81 0.78 0.88 0.68
The uncertainty in the ETN values as reported was ±0.04. bReference 38.
a
change, ΔH°, which measures solute−solvent interactions, can be correlated with the polarity parameters of ILs. The arrangement of PILs in dissociating the HCOOH and CH3CH2COOH is in the inverse order of the polarity index (π*), whereas for CH3COOH it is in the increasing order of π*. The efficiency of PILs in controlling the ΔH° for HCOOH and CH3CH2COOH is in the opposite order of π*, which suggests the lower relevance of the electrostatic field effect in controlling the dissociation of acid. This implies the involvement of specific interactions in the dissociation of carboxylic acids. The polarity parameters that represent specific interactions are α and β, although direct dependence on either α or β in the dissociation of HCOOH and CH3CH2COOH is not revealed by the arrangements of PILs. However, the dependence of ΔH° on α or β can be explained in an alternate way. It is proposed by several workers that in ILs the hydrogen bond donor ability of the cation (α) and the hydrogen bond acceptor ability of the anion (β) remain in equilibrium with each other.47,48 The resultant of these two determines the behavior of IL. Therefore, a solute will always experience lesser affinity either toward the cation or anion than that predicted by the Kamlet−Taft parameters (α, β, π*). In other words, α is partially reduced by β and vice versa. Thus, in all ILs, a competition exists between the α and β values in the presence of solute. In view of these counterbalanced interactions present in ILs, the arrangement of different PILs in affecting the ΔH° values should be correlated with the relative α and β values. The higher dissociation of carboxylic acids in PILs is due to either the stabilization of proton (H+) or the stabilization of conjugate base of acid. In [HPyrd][HCOO], α is higher than β; thus, the dissociation of HCOOH and CH3CH2COOH is controlled
[HPyrd]+ > [HPyrr]+ > [HmIm]+ [HmIm]+ > [HPyrr]+ > [HPyrd]+ [HPyrd]+ > [HPyrr]+ > [HmIm]+
As apparent from these arrangements of PILs based on the ΔH° values, there is no definite order of PILs in which they interact with carboxylic acids. The potential of PILs based on the ΔH° values is different from those obtained from the H0 values for HCOOH, CH3COOH, and CH3CH2COOH at 298 K. This discrepancy in the dissociating carboxylic acids in PILs might be due to the thermal effects that were not accounted for in the determination of the H0. However, unexpectedly very low ΔH° values for all carboxylic acids were obtained in [BMIm][HCOO]. This points out that the nature of intermolecular interactions between carboxylic acids and [BMIM][HCOO] is similar to PILs and carboxylic acids based on the ΔH° values. The unusual behavior of [BMIm][HCOO] is in contradiction with the very large H0 values for all carboxylic acids. Thus, the intermolecular force between the [BMIM][HCOO] and carboxylic acids cannot be predicted from the ΔH° values. The ambiguous behavior of [BMIm][HCOO] can be explained by employing the polarity parameters, which control the solute−solvent interactions. The polarity parameters of ILs 5543
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The Journal of Physical Chemistry B
■
by α. In [HPyrr][HCOO], α is lower than β; hence, dissociation of HCOOH and CH3CH2COOH is favored by β. In [HmIm][HCOO] both α and β are equal in magnitude; therefore, the least dissociation of carboxylic acids and consequently the higher ΔH° for HCOOH and CH3CH2COOH are observed in [HmIm][HCOO]. Hence, the arrangement of different PILs based on the ΔH° for HCOOH and CH3CH2COOH is according to their ability to bind with carboxylic acids by specific interactions. Surprisingly, a different order of PILs in the increasing order of their π* values was noted to be effective in the dissociation of CH3COOH. This indicates the involvement of Coulombic forces in the dissociation of CH3COOH, which seems highly improbable. The arrangement of PILs for CH3COOH in the increasing order of π* value might be due to the higher dipole moment of CH3COOH (1.74 D) than those of HCOOH (1.41 D) and CH3CH2COOH (0.63 D). The strong bond dipoles present in the CH3COOH combine well with the Coulombic region created inside PILs. However, similar observations could not made for the AILs as, except for 1-butyl-3-methylimidazolium formate ([BMIm][HCOO]), other ILs, for example, 1-butyl-3-methylpyrrolidinium formate ([BMPyrr][HCOO]) and 1-butyl-3-methylpiperidine formate ([BMPyrd][HCOO]), were solid at room temperature. However, as can be seen from the polarity parameters for [BMIm][HCOO], α and β are small as compared to those of PILs. A small difference in the α and β indicates the lower tendency of [BMIm][HCOO] toward the specific interactions with solute than of PILs. Therefore, the polarity parameters of [BMIm][HCOO] also support the lower strength of carboxylic acids. Further investigations are required to explore the role of different protic and aprotic ILs having similar cations/anions in the dissociation of weak acids.
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Article
ASSOCIATED CONTENT
S Supporting Information *
Additional tables and drawings associated with the determination of thermodynamic parameters; degree of dissociation (α) and ln Keq versus 1/T for HCOOH, CH3COOH, and CH3CH2COOH in ILs. This material is available free of charge via the Internet at http://pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Author
*(A.K.) E-mail: [email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS S.K.S. is grateful to UGC, New Delhi, for awarding him a research fellowship. A.K. thanks DST, New Delhi, for a J. C. Bose National Fellowship (SR/S2/JCB-26/2009).
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REFERENCES
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CONCLUSIONS
We have compared the potential of protic and aprotic ionic liquids in the dissociation of HCOOH, CH3COOH, and CH3CH2COOH measured by the Hammett acidity function. Protic ionic liquids were observed as promising candidates in the dissociation of carboxylic acids because of higher acidic strength than their aprotic counterparts. The greater dissociation of carboxylic acids in protic ionic liquids was noted to depend on the intensity of specific interactions between them. The activation energy of proton transfer (Ea,H+) further fortifies this view. Higher Ea,H+ values for carboxylic acids obtained in aprotic ionic liquid than in protic ionic liquid lower the possibility of specific interactions. Thermodynamic parameters (ΔH°, ΔS°), calculated by using a novel “indicator overlapping method”, also signify the participation of specific solute− solvent interactions in the dissociation of carboxylic acids. The ability of different protic ionic liquids as observed by ΔH° value in dissociating carboxylic acids was observed to depend on the relative hydrogen bond donor (α) and hydrogen bond acceptor (β) tendencies. The observations drawn above will be useful in designing efficient media for weak acids. This will help in avoiding the use of corrosive acids for various applications. To the best of our knowledge, this is the first report comparing the potential of protic ionic liquid and aprotic ionic liquid in the dissociation of carboxylic acids. 5544
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