Composition Dependent Physicochemical Property Data for the Binary

Article pubs.acs.org/jced

Composition Dependent Physicochemical Property Data for the Binary System Water and the Ionic Liquid 1‑Butyl-3methylimidazolium Methanesulfonate ([C4mim][MeSO3]) Annegret Stark,† Anthony W. Zidell,‡ Joseph W. Russo,‡ and Markus M. Hoffmann*,‡ †

Institute for Industrial Chemistry, University of Leipzig, Linnéstr. 3-4, 04107 Leipzig, Germany Department of Chemistry, The College at Brockport, State University of New York, Brockport, New York 14420, United States

‡

S Supporting Information *

ABSTRACT: A number of composition dependent physicochemical properties were measured for the water− [C4mim][MeSO3] binary system at 85 °C. Specifically, we provide data for density, heat capacity, heats of dissolution, conductivity, viscosity, as well as the self-diffusion coefficients of the cation, anion, and water. Ionic liquids based on the MeSO3 anion constitute a suitable medium for various chemistries. This study addresses the need to provide physicochemical data on [C4mim][MeSO3] and the water− [C4mim][MeSO3] binary systems that are currently unavailable. Additional analysis of the combined data is also presented to provide further insight into the molecular level behavior of the water−[C4mim][MeSO3] binary system.

■

INTRODUCTION Ionic liquids (ILs) are salts that are liquid below 100 °C. The interest in ILs has continued to rise nearly exponentially in time. A SciFinder search for just review articles related to IL research and published in 2011 produced hundreds of hits concerning areas such as synthesis,1−4 catalysis,5,6 materials science,7−9 analytical chemistry,10−12 and many other applications for the chemical industry.13 The particular IL 1-butyl-3methylimidazolium methanesulfonate, [C4mim][MeSO3], has become of interest as it has among other ILs proven to be a suitable medium for sugar chemistry.14−16 In our own recent research we have studied oxidations17 and condensation reactions such as the esterification18 and the conversion of fructose to 5-hydroxymethylfurfural (HMF) in [C4mim][MeSO3], finding that the water content has a large effect on the yield and selectivity of these reactions.19 This motivated us to gather composition dependent physicochemical property data for the water−[C4mim][MeSO3] binary system to better understand the role of water on the reaction progress. There is a lack of available data for physicochemical properties of [C4mim][MeSO3] that is most likely due to its relatively high melting point which is near 80 °C. The high melting point of [C4mim][MeSO3] is also one reason why we focused first in a prior study on the binary system of water and [C2mim][MeSO3] that has a melting point near 40 °C and is a metastable liquid at room temperature.20 At the time of writing this report there were no entries for [C4mim][MeSO3] in the IUPAC Ionic Liquids Database project “ILThermo” accessible on the Internet through the National Institute of Standard & Technology Web site.21 Searching the literature for studies involving [C4mim][MeSO3], we found only a few related works that should be mentioned here. The Coutinho group and © XXXX American Chemical Society

collaborators have explored the aqueous phase behavior of [C4mim][MeSO3] (and a number of other ILs) with aqueous salt solutions for the application of extracting vanillin,22 liquid extractive fermentation,23 and lipase purification.24 Surface tension and chemical shift data for aqueous solutions of [C4mim][MeSO3] have also been reported.25 For the neat [C4mim][MeSO3] there are detailed sum frequency generation and X-ray studies available by the Baldali group providing insights as to the structural differences between the bulk and the surface.26,27 The polarity of [C4mim][MeSO3] has also been assessed including the provision of Kamlet−Taft empirical parameters.28−31 Given the lack of knowledge about the physicochemical properties of [C4mim][MeSO3], we conducted a composition dependent study (measured at 85 °C) with two batches of differing purity level as further specified in the Experimental Methods. For simplicity we will in the remainder of this report refer to the [C4mim][MeSO3] ILs with different purity levels as “pure IL” and “impure IL”.

■

EXPERIMENTAL METHODS The “pure IL” [C4mim][MeSO3] (CAS no. 342789-81-5) of 100 w = 99 purity (specified as water 100 w = 0.29, halide 95, with water (100 w < 2.0), halide (100 w 1.5), and 1-methylimidazole (100 w < 0.17), hence fulfilling manufacturer specifications. Considering that residual halide was found, a synthetic strategy involving ion exchange between 1-butyl-3-methylimidazolium chloride and methanesulfonic acid during the preparation of the IL appears likely. This is supported by the fact that the pH value of a 0.2 M aqueous solution of the impure IL was determined to be 1.6, hence indicating the presence of acid. The deionized water used was purified by UV purification systems (either Barnstead or Millipore). When preparing the aqueous IL samples it was noticed that samples remained liquid at room temperature if the mole fraction of water, xw, was >0.3 for the impure IL and >0.5 for the pure IL. The Karl Fischer (KF) titrations were performed either on a Denver Instrument, model 275 KF or with an AQUA 30 titrator from Electrochemie Halle (ECH) GmbH. Reproducibility was 0.5 % or better in either case. It should be noted that water determinations of the ILs were carried out on a regular basis prior to measurements and that the respective value was taken into account in the calculations of the molar or mass fractions of water in the mixture. The heats of dissolution and the conductivity were measured concurrently. Using a Mettler Toledo RC 1 reaction calorimeter the heats of dissolution were measured in adiabatic mode. The conductivity measurements were performed with a LF 340 conductivity meter with a TetraCon measuring cell from WTW for which a glass attachment for the reaction calorimeter was devised. The conductivity cell was calibrated against aqueous solutions of KCl.32 The precision of the conductivity measurements are discussed in the Results and Discussion section. Prior to initially loading the reaction calorimeter with 100 cm3 of IL, the reaction calorimeter was flushed for 20 min with dry argon. The exact weight of the IL loaded into the calorimetric was determined by mass difference to ±0.01 g. The water content of the IL loaded into the calorimeter was measured per Karl Fischer titration using a portion of the remaining IL. While the large weight of the loaded IL prohibited the use of an analytical balance the amount of subsequently added water was determined by mass difference with an analytical balance to ± 1·10−5 g. The exact procedure for how known amounts of water were incrementally added to the IL is described in detail in a prior report.20 The temperature measurements were recorded to a precision of 0.001 °C. The temperature change caused by the water addition was determined by the equal area method. Despite the relatively high temperature of 85 °C, we observed no visible water condensation in the reaction calorimeter headspace for xw < 0.4. At higher water content, some water evaporated from the IL solution and condensed on the colder reaction calorimeter headspace. The loss of water was accounted for by the concurrently measured drop in conductivity. Overall, the measurement uncertainty for the molar heats of dissolution is estimated to be ± 100 J·mol−1 in the range of 0.1 < xw < 0.4. For xw > 0.4 uncertainty increases to ± 300 J·mol−1 due to the uncertainty for accounting the loss of evaporated water. For xw < 0.1, the uncertainty is also estimated to be larger (± 500 J·mol−1) because as xw approaches zero the uncertainty of delivering minute amounts of water increases. Nevertheless, we were able to reproduce the heats of dissolution of KCl33 and [C4mim][BF4] in water34 within 5 % of the literature values. Heat capacities were measured independently of the reaction calorimeter experiments by differential scanning calorimetry35

on a DSC 7 from Perkin-Elmer. IL−water samples were prepared first in 5 cm3 vials and heated to 85 °C for thorough mixing. For samples that remained liquid at room temperature, 10−15 mg of sample was placed into an aluminum pan using disposable Wiretol micropipets and then immediately sealed with an aluminum lid using a hand press. (The press and the aluminum pans and lids were all obtained from Perkin-Elmer). The transfer of the solid samples proved to be more challenging and we were able to only transfer up to 5 mg of sample without the risk of getting sample in contact with the seal area. An analytical balance with a precision of 1·10−5 g was used to weigh the empty and loaded, sealed pans. The weight variations of the empty aluminum pans were kept within 5·10−5 g for a measurement series. Independent exposure tests of an aluminum pan and lid to wet IL for several days and at 90 °C for several hours did not result in any visual alteration of the aluminum pan and lid, and their weight remained the same, indicating that the IL solutions did not corrode the aluminum material. For the heat capacity measurements, the heating profile was as follows: the sample was held at 80 °C for 5 min then heated to 90 °C at a rate of 1 °C·min−1, and finally held at 90 °C for another 5 min. Each measurement series including the calibration runs with empty pans and pans filled with a sapphire crystal standard were completed within the same day. From prior work with the related water−[C2mim][MeSO3] binary system, we found the reproducibility of the DSC scans to be within 3 %.20 However, the uncertainty for the heat capacity measurements from samples that were solid at room temperature have a much higher uncertainty as shown in the Results and Discussion section because of the small sample amount we were able to transfer. Prior to preparing the NMR samples, dissolved gases in the water and the IL were first separately removed with seven freeze−pump−thaw cycles. The NMR sample solutions were then prepared in a nitrogen glovebox and transferred into a melting capillary tube of 0.8 mm inner diameter immediately after heating to prevent premature solidification within the needle. The prepared melting capillary tube was then hooked up to a vacuum line and subsequently flame-sealed. The sealed melting tube was next placed into a standard 5 mm NMR tube that was then filled with 0.75 cm3 of deuterated dimethyl sulfoxide, DMSO-d6 (Norell Inc.) as the lock solvent because its residual signal does not interfere with the IL resonances and it has a high boiling point. The DMSO-d6 was added to the NMR tube which was then also flame-sealed. The NMR data were obtained on an Avance 300 NMR instrument from Bruker-BioSpin using a broadband probe with temperature held at 85.0 ± 0.1 °C. Sample rotation was turned off during acquisition. Self-diffusion coefficients, D, were measured by linearly varying the pulsed magnetic field gradient strength, g, between (5.6 and 56) G·cm−1, in 16 increments (16 or 32 scans each) using a pulse program originally published by Jershow and Mueller.36,37 This particular pulse program reduces systematic errors in D from eddy currents and thermal gradients. The obtained stimulated echo intensity, I, was then fitted to the function shown in eq 1 2 2 2

I(g ) = I0e−Dγ g

δ ((4Δ− δ)/ π 2)

(1)

where I0 is the stimulated echo intensity in absence of any gradient, γ is the gyromagnetic ratio for protons, δ is the length of the gradient pulse that was of sine shape, and Δ is the diffusion time that was held constant at 0.1 ms. Delays for B

dx.doi.org/10.1021/je300045c | J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

Article

gradient recovery and eddy currents were set to (0.1 and 5) ms, respectively. The values for δ ranged from (1.5 to 5.5) ms. Diffusion measurements were sometimes repeated with δ values optimized for the water signal to obtain more reliable selfdiffusion coefficients for water as it self-diffused about three times faster than the cation or anion. However, for some samples the water resonance partially overlapped with an IL resonance. In these cases the water intensities were found through deconvolution of the overlapping signals. Since the IL cation shows eight lines in the proton spectrum the standard deviation from the average of the IL cation self-diffusion coefficient can be used as a measure of reproducibility. Standard deviations of the IL cation self-diffusion coefficients were all within 6 % of the average value (typically even within 3 %). Viscosities of the impure IL were determined with a Brookfield LVDV-III UCP rheometer using the disk spindle CP-40 that has a measurement range from (0.15 to 3065) mPa·s. The sample holder and spindle were made from stainless steel. The sample holder was temperature-controlled by a water circulating bath to within (0.01 up to 80) °C. The calibration of the rheometer was periodically checked with Brookfield standard oils. The Newtonian fluid behavior of the [C4mim][MeSO3] was confirmed as observed by a linear dependence between sheer stress and sheer rate. The accuracy of the measurements were tested against literature values for [C2mim][BF4]38,39 and were found to be within 3 %. Further details on the measurement procedure are provided elsewhere.20 Viscosities and densities of the pure sample were measured in parallel with an Anton Paar rolling ball viscometer and a DMA 4100 vibrating tube density meter. Solutions with xw > 0.9 were measured with a 1.6 mm capillary; otherwise a 1.8 mm capillary was used. The rolling ball viscometer was calibrated against neat water. All measurements were averaged from a total of six repetitions, two measurements each at three different angles. The standard deviation of the 6 repetitions was always within 6 % (mostly within 2 %). The viscosity of the pure IL sample xw = 0.9 was measured with both capillaries and agreement was within 3 %. All density measurements were obtained with internal sample viscosity compensation. Both instruments use a Peltier system to control temperature to within 0.02 °C. As an accuracy check the density of pure water was measured in the range of (20 to 85) °C, and the results agreed with the literature values40 within 0.001 g·cm−3.

Figure 1. Specific heats for the water−[C4mim][MeSO3] binary system as a function of water mass fraction, cw, at 85 °C: ○, pure IL; ∇, impure IL. The solid line is a linear least-squares fit for the pure IL data set that includes the value for pure water taken from ref 40.

definitive conclusion can be drawn if the impurities in the impure IL caused the specific heats for the binary system to be slightly higher than for the binary system with pure IL. Therefore, the least-squares fit parameters for the fit line in Figure 1 were used for all subsequent evaluations of the heats of dissolution discussed in the next subsection. Heats of Dissolution. The addition of water into [C4mim][MeSO3] was observed to be exothermic as was previously noted for the similar IL [C2mim][MeSO3].20 Figure 2 shows the incremental heat released per mol of added water

Figure 2. Incremental heat released per molar addition of water into the water−[C4mim][MeSO3] binary system at 85 °C: ∇, ○, and □, impure IL; Δ and ◊, pure IL. The lines are linear least-squares fits to the data.

■

to the water−[C4mim][MeSO3] binary system at 85 °C, with three runs using the impure IL and two runs using the pure IL. The data shown in Figure 2 is also provided in Table S2 in the Supporting Information. The solid lines represent linear leastsquares fits through the combined data from the two different purity grade ILs. The parameters of the least-squares fits are summarized in Table 1 along with R2 values and standard deviations for each of the two combined data sets. In our previous study for the related water−[C2mim][MeSO3] system the heat released per added mol of water

RESULTS AND DISCUSSION Specific Heat. The specific heats for binary systems composed of water and pure or impure IL are shown in Figure 1 as a function of mass fraction, cw, of water and are also supplied in tabular form in the Supporting Information, Table S1. The literature value for pure water is included.40 For reasons described in the Experimental Methods, uncertainty is largest for the specific heats at the smallest water mass fractions. Nevertheless, within the scatter of the data, ideal behavior, that is a linear relationship of specific heat and water mass fraction, is indicated. Thus, a least-squares fit was applied for the pure IL data set and is included in Figure 1. Specifically, the values for slope and intercept are (2.24 ± 0.12 and 1.92 ± 0.04) J·g−1·K−1 with an overall standard deviation of ± 0.11 J·g−1·K−1 and R2 value of 0.979. The specific heats for the data from impure IL are all larger than the fit line except for one data point at the lowest water mass fraction. However, the deviations are all within the standard deviation of ± 0.11 J·g−1·K−1 and no

Table 1. Linear Fit Coefficients for the Heat Released per Added Moles of Water slope [C4mim][MeSO3] pure IL impure IL C

−1

kJ·mol

−5.04 ± 0.15 −7.36 ± 0.35

σ

intercept kJ·mol

−1

4.63 ± 0.05 6.50 ± 0.10

R

2

0.988 0.971

kJ·mol−1 0.15 0.35



Article

showed a maximum of 4.77 kJ·mol−1 near xw = 0.1 at 85 °C,20 which in value is similar to the heat released per molar addition of water at xw values approaching zero in Figure 2. However, for [C4mim][MeSO3]−water mixtures, such a maximum at a low xw is not distinctive, although the values for the heat released per mol of added water tend to be lower than what the shown linear fit line predicts for xw approaching 0. Also, at xw = 1, the heat released should approach zero, but extrapolation of the linear fit trend to xw = 1 results in values for the heat released of −0.41 kJ·mol−1 for the pure IL and −0.86 kJ·mol−1 for the impure IL. Therefore, the true heat released per molar addition of water is likely to deviate from linearity as xw approaches 1. Nevertheless, for the measured data range and within the scatter of the data, the linear fit functions are reasonable with standard deviations of ± 0.15 kJ·mol−1 for the pure IL and ± 0.34 kJ·mol−1 for the impure IL. As clearly evident in Figure 2 the heat released per molar addition of water into the impure IL is systematically larger than that of the pure IL. Given that one major impurity found in the impure IL is acid this observation can be easily explained by the significant contribution to the heats of dissolution from the present acid impurity. To obtain the excess molar enthalpy of mixing from the data in Figure 2, the incremental heats need to be summed up and the obtained sum be divided by the total number of moles of both water and IL. For one of the experimental runs for each of the binary systems with pure IL (data set ∇ in Figure 2) and impure IL (data set Δ in Figure 2), there was already significant amounts of water present (∼0.1 xw) before the first addition of water. Thus, these two data sets could not be used to directly obtain the excess molar enthalpy of mixing because a significant heat contribution from adding the first approximately 0.1 xw of water was not measured. However, one can also evaluate incrementally the excess molar enthalpy of mixing from the linear fit lines to the data sets in Figure 2. For a given arbitrary initial number of moles of neat IL, nIL, the number of moles of water that needs to be added to obtain a final desired water mole fraction is

nw =

x wnIL 1 − xw

showing the excess molar enthalpy of mixing, HE, for the binary system of water with pure and with impure IL. Although the data sets for pure and impure IL vary in magnitude, both HE minima are near xw = 0.6, with HE = −1.6 for the pure IL. For the related system water−[C2mim][MeSO3], the minimum of the excess molar enthalpy of mixing at 85 °C was also found at xw = 0.6 with HE = −2.0.20 The smaller amount of heat liberated when water is dissolved in [C4mim][MeSO3] compared to [C2mim][MeSO3] could be explained by the [C4mim][MeSO3] with the longer aliphatic side chain requiring more energy to realign in order to accommodate the added water. Density. Density data at 85 °C were only obtained for samples where the aqueous IL solutions are liquid at room temperature and can be injected into a density meter. To evaluate excess molar volumes from the density data, the density of the neat IL at 85 °C must be known besides the density of pure water at 85 °C, which was taken from ref 38. Unfortunately, to the best of our knowledge, experimental density data for the neat [C4mim][MeSO3] have not been reported yet. However, there is a study by Blesic et al.41 reporting physicochemical properties of a number of 1-alkyl-3methylimidazolium alkanesulfonate ILs. The authors could show that molar volumes are additive with respect to the number of carbon atoms present in the alkyl chains, that is, there is a linear relationship between the molar volume of the neat IL at a given temperature and the number of carbon atoms in the alkyl side chain. Thus, we interpolated from their reported temperature dependent density data of other neat 1alkyl-3-methylimidazolium alkanesulfonate ILs the density of neat [C4mim][MeSO3] to be 1.134 g·cm−3 at 85 °C. Table 2 summarizes the measured densities and excess molar volumes for the solutions with pure and impure IL. The densities of the pure and impure IL are very similar resulting only in relatively Table 2. Density, ρ, and Excess Molar Volume, VE, for Water−[C4mim][MeSO3] at 85.00 ± 0.02 °C and Ambient Pressure as a Function of Water Mole Fraction, xw

(2)

ρb xwa

With the obtained nw and the linear fit equations for the data sets of Figure 2, the amount of released heat can be evaluated, and the excess molar enthalpy of mixing is obtained from the cumulative heats divided by the total number of moles of both water and IL. In this fashion we obtained the lines in Figure 3

VE

g·cm−3

cm3·mol−1 c

0.975 0.950 0.925 0.900 0.852 0.801 0.749 0.702 0.654 0.559 0.975 0.958 0.937 0.914 0.895 0.860 0.975

Figure 3. Excess molar enthalpy of mixing, HE, for the water− [C4mim][MeSO3] binary system 85 °C: ○ and □, impure IL; ◊, pure IL. The lines were obtained from the linear least-squares fits in Figure 2 and eq 2 as described in the text.

Pure IL 1.0039 1.0297 1.0482 1.0623 1.0812 1.0949 1.1041 1.1103 1.1144 1.1207 Impure ILc 1.0010 1.0200 1.0386 1.0547 1.0641 1.0815 1.0010

0.0334 −0.0024 −0.0413 −0.0823 −0.1507 −0.2404 −0.2931 −0.3339 −0.3106 −0.3025 0.0997 0.0797 0.0308 −0.0517 −0.0628 −0.2571 0.0997

Mole fraction precision is ± 0.001. bInstrument precision is ± 0.0010 g·cm−3. cAs defined in the Experimental Methods.

a

D



Article

small differences for the excess molar volumes shown in Figure 4. Therefore, the combined data sets were used to fit the excess

Figure 5. Molar conductivity as a function of water mass fraction at 85 °C for the water−[C4mim][MeSO3] binary system: ∇, ○, and □, impure IL; Δ and ◊, pure IL. The third order polynomial fit line was obtained from the combined five data sets shown.

Figure 4. Excess molar volume for the water−[C4mim][MeSO3] binary system at 85 °C for the pure IL, ○; and the impure IL, □. The solid line is the Redlich−Kister fit function according to eq 3 and using the parameters in Table 3.

more scatter in the conductivity and thus the molar conductivity data at mass fractions greater than about 0.2, which is mainly due to the increased uncertainty of the solution composition because of water evaporation and recondensation at the cooler reactor top of the reaction calorimeter as explained in the Experimental Methods. Regardless, it is clearly evident in Figure 5 that the presence of the impurity had no measurable effect on the solution conductivity. The third order polynomial fit lines to the data in Figures S1 and 5 were therefore obtained from the combined data sets from both pure IL and impure IL and resulted in an R2 value of 0.986 and a standard deviation of ± 0.27 S·m−1 for the conductivity data and an R2 value of 0.988 and a standard deviation of ± 0.85·10−4 S·m2·mol−1 for the molar conductivity data, as summarized in Table 4. Table 4 also includes the result of a linear fit to the raw conductivity data at water mass fractions less than 0.13 which reduces the standard deviation for this data range to ± 0.07 S·m−1. Compared to the water−[C2mim][MeSO3] system,20 the conductivity as a function of the mass fraction is similar to the water−[C4mim][MeSO3] system of this study, e.g., at cw = 0.35, σ = 12.00 ± 0.55 S·m−1 (from polynomial)20 for the former and σ = 9.86 ± 0.27 S·m−1 (from polynomial) for the latter. The difference in the conductivity is mostly due to differences in viscosity, as described below. Viscosity. The viscosity data for binary mixtures of water and the pure IL at 85 °C and impure IL at 80 °C are tabulated in the Supporting Information in Table S4. As stated in the Experimental Methods, the viscosities of the binary water−IL systems with the pure and impure IL were measured with different instruments, where for the measurements with the impure IL we were limited to 80 °C. In Figure 6, we show the viscosity data in terms of fluidity, 1/η, against mass fraction of water cw. It is well-known that the viscosity is very sensitive to some impurities,42 where low-molecular weight impurities such as acid impurities can be expected to result in increased fluidity. However, a decrease of fluidity will be caused by a lower temperature. These two competing effects of present impurities and lower temperature on the fluidity seem to approximately cancel each other and as a result the two data sets of water-pure IL and water-impure IL fall on top of each other in Figure 6. We fitted the data with pure IL to a second order polynomial that resulted in an R2 value of 0.9994 and a standard deviation of ± 0.019 mPa−1·s−1(see Table 4). The graph in Figure 6 is very similar in behavior as previously found for [C2mim][MeSO3] where at small mass fractions up to about 0.25 the

molar volumes to the Redlich−Kister polynomials shown in eq 3 with the obtained BP fitting coefficients listed in Table 3. V E = x wx IL ∑ BP (x w − x IL)P

(3)

P

Table 3. Redlich−Kister Polynomial Fitting Coefficients to the VE Data in Figure 4 P

BP

0 1 2 3 4

−0.6759 −2.9450 −1.4175 4.0301 3.1784

The standard deviation of the fit line to the data in Figure 4 was determined to be 0.047 cm3·mol−1. A minimum of VE near xw = 0.7 is evident in Figure 4. Even though we do not have experimental density data for xw < 0.5 we used eq 3 to obtain VE and subsequently molar volume data for evaluating the molar conductivity from raw conductivity data presented in the next subsection. In this regard, we note that the water− [C4mim][MeSO3] is behaving nearly ideally. Specifically, the minimum for VE in Figure 4 is only near −0.3 cm3·mol−1, which is in magnitude slightly less than the value of −0.4 cm3·mol−1 we observed in our prior study for the water-[C2mim][MeSO3] binary system at 85 °C, where the minimum occurred at xw = 0.6.20 Conductivity. In the Supporting Information, we list in Table S3 the raw conductivity data for binary mixtures of water and pure as well as impure [C4mim][MeSO3] and plot these data as a function of mass fraction of water in Figure S1. To evaluate molar conductivities, the molar IL concentration must be known which means the density must be known, which we obtained from eq 3 for VE. Table S3 includes the evaluations of VE, V, and density ρ. In Figure 5 we show the molar conductivity as a function of mass fraction of water as this graph indicates a linear trend for low water mass fractions, smaller than about 0.15, which also holds true for the raw conductivity data shown in Figure S1. We observed such linear trend at low water concentration also previously for the water−[C2mim][MeSO3] system.20 There is E



Article

Table 4. Mass Fraction of Water (cw) Dependent Properties of the Water−[C4mim][MeSO3] Binary System at 85 °C and Ambient Pressure and Constants Fitting the Data to the Polynomial a3cw3 + a2cw2 + a1cw + a0: Conductivity, σ, Molar Conductivity, Λ, Fluidity, η−1, and Self-Diffusion Coefficients of Cation, D+, and Anion, D− σ

σ −1

S·m a0 a1 a2 a3 σ R2 a

0.835 19.73

0.072a 0.988a

η−1

Λ −1

S·m

0.8108 21.407 −21.193 96.991 0.27b 0.986b

10

−4

S·m ·mol 2

−1

−1 −1

1.5668 55.256 −105.79 594.25 ± 0.85 0.988

η−1 −1 −1

D+ −10

D− 2 −1

m2·s−1

mPa ·s

mPa ·s

0.1002 0.4057 2.4888

0.0316 1.0100 1.5601

0.5480 12.323 17.435

0.4893 15.038 22.258

± 0.019c 0.9994c

± 0.002d 0.9993d

± 0.057 0.9995

± 0.070 0.9995

10

m ·s

−10

10

Fit range: 0 < cw < 0.13. bFit range: 0 < cw < 0.41. cFit range: 0.08 < cw < 1.00. dFit range: 0.08 < cw < 0.25.

Figure 7. Proton chemical shifts at 85 °C as a function of composition for the water−[C4mim][MeSO3] binary system: Δ, aromatic protons; ◊, protons from the butyl side chain of the cation; □, NCH3 protons; ○, anion protons; +, water protons for the impure IL; ×, water protons of the pure IL; *, residual signals from water in the external DMSO-d6 lock solvent (near 2.8 ppm) and the lock solvent itself (near 2.25 ppm). Solid symbols are for the impure IL and open symbols are for the pure IL. The chemical shifts are relative to the signal from the terminal methyl proton of the butyl side chain of the cation. The graph is constructed to illustrate how the proton NMR spectral lines move with changes in composition.

Figure 6. Fluidity (inverse viscosity) as a function of water mass fraction for binary water systems with the pure IL at 85 °C, ○, and the impure IL at 80 °C, □. The solid line is a second order polynomial fit for the pure IL data.

fluidity is nearly linearly dependent. Specifically, a linear leastsquares fit to the fluidity data up to cw = 0.25 results in a R2 value of 0.997. Compared to the water−[C2mim][MeSO3] system,20 the decrease in fluidity for the water−[C4mim][MeSO3] system is overall rather modest, e.g., at cw = 0.50, 1/η = 0.9667 ± 0.015 mPa−1·s−1 (at 80 °C, from polynomial)20 for the former and 1/η = 0.9253 ± 0.019 mPa−1·s−1 (at 85 °C, from polynomial) for the latter. Blesic et al. found for a number of neat 1-alkyl-3methylimidazolium alkanesulfonate ILs linear relations for the viscosity with the number of carbon atoms in the alkyl chain in the cation.41 Given the very similar behavior of the water− [C4mim][MeSO3] system compared to the water−[C2mim][MeSO3] system, such linear relations may also hold true for these ILs in binary mixtures at a fixed composition. Chemical Shift. In Figure 7 we show the effect of water on the proton NMR spectral lines of [C4mim][MeSO3] that are also tabulated in the Supporting Information, Table S5. The chemical shifts are reported relative to the chemical shift of the terminal methyl group of the butyl side chain. We note that the residual signal from the external deuterated lock solvent DMSO-d6 was consistently found at 2.235 ± 0.015 ppm for all spectral data in Figure 7. In Figure 7 it is evident that all spectral lines of the cation move toward lower chemical shifts, i.e., the cation protons become more shielded as water is added. In contrast, the singlet for the anion near 1.7 ppm moves toward larger chemical shifts, i.e., the methyl protons of the anion become less shielded as water is added. We have previously observed these trends also for adding water into [C2mim][MeSO3] and explained these observations by a competition of the water and the cation interacting with the anion where water is interacting more preferentially with the

anion (formation of solvation shells) than the cation does.20 The presence of the impurities in the impure IL has no noticeable effect on the chemical shifts of the cation and anion signals in the proton spectrum as is clearly evident in Figure 7 by the overlap of the open (pure IL) and filled (impure IL) symbols. However, the water resonance is strongly affected by the presence of the impurities in the impure IL. Especially at low water content the water resonance moves to significantly larger chemical shift values, which is indicative of the presence of acid impurities. Specifically, the difference of the water chemical shift between the pure and impure binary water−IL systems approaches approximately 2 ppm as xw approaches zero. Yasaka et al. have actually used this strong effect of acid impurities on the water chemical shift to evaluate quantitatively the amount of free acid impurity present in the IL.43 According to their calibration curve of adding CH3SO3H into [C4mim][MeSO3] at 80 °C with a molality of 0.093 mol·kg−1 of water present, we can estimate that the amount of acid impurity in the impure IL is approximately 0.045 mol·kg−1. Figure 7 shows that the chemical shift of the water resonance for the pure IL first decreases until about xw = 0.5 upon which the chemical shift of water increases. This trend is presently not understood. In the study on the parent compound [C2mim][MeSO3], the curvature pointed continuously upfield with increasing IL content, due to the breaking of the hydrogenbond network of water and incorporation of anions. At high water concentration, this is clearly also the case for [C2mim]F



Article

Also clearly evident from Figure 8 is that the impurities present in the impure IL do not affect the self-diffusion coefficients of the anion and cation. There is much more scatter in the self-diffusion coefficient data for the water because, as can be seen from the chemical shifts plotted in Figure 7, the water signal was often coinciding with one of the cation signals, and the spin echo signal from the much faster moving water signal becomes very small compared to the IL spin echo signals in particular for the samples with the lowest water content. Nevertheless, also for the water self-diffusion coefficient there is no apparent influence of the impurities. For that reason, the solid fit line for the water self-diffusion coefficient and the second order polynomial fit lines for the cation and anion selfdiffusion coefficients shown in Figure 8 were obtained from the combined data sets from the samples with pure and impure IL. The corresponding fit coefficients are included in Table 4. Overall, the self-diffusion coefficients for the cation and anion show some deviation from linearity in Figure 8, which is different from our previous findings for [C2mim][MeSO3] where linearity was observed. Otherwise, the data is very similar insofar that the water moves about 3 times as fast as the IL over the entire concentration range. The self-diffusion coefficient of water is about 10 % lower in this study for [C4mim][MeSO3] compared to our previous study on [C2mim][MeSO3], while the self-diffusion coefficients of the cation and anion are about 15 % lower. Analysis of the Combined Results. The fact that the molar conductivity, density, and fluidity of the water− [C4mim][MeSO3] binary system is continuously increasing up to the highest water mass fraction of about 0.4 (corresponding to a water mole fraction of about 0.9) indicates that the IL structure must continuously be broken and the ions solvated upon addition of water. When combining some of our results to generate Figures 9−11, which we will explain shortly,

[MeSO3]-water mixtures. A possible explanation of the increase of the water chemical shift with decreasing water content for xw < 0.5 could be the longer alkyl side chain of the cation compared to [C2mim][MeSO3], which may increase the strength of the interaction between the anion and the charge bearing imidazolium ring as witnessed by the overall downfield shift of the imidazolium-protons in the [C4mim]-based compared to the [C2mim]-based mixtures.20 As a result, small amounts of water in the IL may not be able to solvate the anion but only weaken the cation−anion interactions somewhat. Interestingly, for the water-[C2mim][MeSO3] binary system the chemical shift of the anion increased by 0.13 ppm from xw = 0.1 to the xw = 0.9,20 which is of the same magnitude for similar dilution of [C4mim][MeSO3] in water (0.10 ppm; Table S5). Self-Diffusion. In Figure 8 is shown the self-diffusion coefficients of the IL cation, IL anion and water at 85 °C as a

Figure 8. Self-diffusion coefficients at 85 °C as a function of water mass fraction, cw, for the water−[C4mim][MeSO3] binary system: □, water; ∇, anion; ○, cation. The gray symbols are for the impure and the open symbols are for the pure IL. The lines were obtained from fitting the respective combined data sets of impure IL and pure IL.

function of composition in terms of water mass fraction, cw. The data points are also presented in the Supporting Information in Table S6. The general trends in Figure 8 are similar to what we previously observed for the water− [C2mim][MeSO3] system.20 The self-diffusion coefficients increase linearly or nearly linearly with water mass fraction. The diffusion coefficient of water is always about a factor of 3 larger than that of the cation and anion (which have very similar diffusion coefficients) over the entire investigated composition range. At the largest water mass fraction, the anion is diffusing clearly faster than the cation, which is the expected behavior because the anion is smaller in size than the cation. However, with decreasing water content this difference diminishes and at water mass fractions less than 0.1 the diffusion coefficient of the anion is observed to be always slightly lower than that of the cation diffusion coefficients. This statement has to be taken with caution as all of the observed differences for cw < 0.1 fall within the standard deviation of the averaged diffusion coefficient of the cation except for the smallest cw value of 0.003 for the impure IL (see Table S6). However, this nonintuitive behavior that the larger cation is diffusing faster has been observed for a large number of neat ILs,38 and can be explained by the existence of fast fluctuating nonpolar domains made up of aggregating aliphatic portions of the cation and anion.20 If such domains are indeed present, one would expect these to be even more pronounced in the water− [C4mim][MeSO3] binary system.

Figure 9. Ratio of molar conductivity and molar conductivity evaluated from the self-diffusion coefficients of cation and anion according to eq 4 for the binary water−[C4mim][MeSO3] system as a function of water mole fraction at 85 °C: □, impure IL; ∇, pure IL.

as we have done in our previous study on the water− [C2mim][MeSO3] binary system we find strong indication also in this study that the principle architecture of the [C4mim][MeSO3] structure stays intact even when the number of water molecules exceeds the number of IL units. We refer for the detailed argumentation to our previous study and briefly reiterate the main points here after providing first several comments regarding Figures 9−11. The data points shown in Figures 9−11 and their evaluations are summarized in Table S7 in the Supporting Information. The conductivities at the water compositions for which we G



Article

interpretation that there exists a strong structural organization of the IL−water mixture. In Figure 11 we inspect the Walden rule, which states that the product of molar conductivity and viscosity is constant for

measured self-diffusion coefficients were obtained by interpolation from the fit coefficients in Table 4. To also obtain the corresponding viscosities we needed to extrapolate to smaller cw values than covered by water−pure IL data. Using the second order polynomial fit parameters in Table 4 obtained from the entire fluidity data set would result in an obvious overestimate for the fluidity of the neat IL. However, using the linear fit equation obtained from fitting just the data up to cw = 0.25 would result in an unphysical value below zero for the fluidity of the neat IL. Thus, we used the second order fit coefficients also shown in Table 4 that were obtained from fitting just the data range up to cw = 0.25. In Figure 9 we show the ratio of the measured molar conductivity and the molar conductivity, ΛNMR, evaluated from the diffusion coefficients of cation and anion, D+ and D−, according to eq 4 ΛNMR = F 2(D+ + D−)/RT

(4)

Figure 11. Walden plot for the water−[C4mim][MeSO3] binary system at 85 °C obtained from both molar conductivity and viscosity data sets, pure and impure IL. The solid line represents the ideal behavior according to the Walden rule. Since fluidity and molar conductivity increase with increasing water content, the water content associated with the individual data points increases from the lower left corner to the upper right corner of the graph. The data points representing the highest water content increasingly approach the solid line.

where F and R are the Faraday and gas constants, respectively. A value less than 1 indicates that cation and anion are not completely independently moving as freely dissociated ions but move at least in part concertedly together as ion pairs or aggregates. In Figure 9 it is evident that Λ/ΛNMR remains constant near 0.5 for all investigated composition ranges except for xw = 0.91 where Λ/ΛNMR jumps up to a value of about 0.8. The value of Λ/ΛNMR = 0.5 is actually quite similar to what has been observed for many neat ILs,38 and even a value of Λ/ ΛNMR = 0.8 is still significantly smaller than the value of 1 expected for infinite dilution, thus strongly indicating that the IL-water mixture remains structurally organized upon dilution of water up to at least xw = 0.8. In Figure 10, we plot the ratio of D+η/T from our data, which according to the Stokes−Einstein equation

aqueous solutions of strong electrolytes by plotting the log of the molar conductivity against the log of the fluidity. The data points in Figure 11 are all clearly below the solid line that represents the ideal behavior predicted by the Walden rule and are generally very close in location to what was found for the water−[C2mim][MeSO3] binary system as well as a number of neat ILs at a wide range of temperatures.44 Only the two data points for the most fluid samples appear to approach the solid line possibly indicating an onset for the IL to lose its central structural organization. Therefore, combining the observations from Figures 9−11 strongly suggests that the principle IL structure remains intact upon addition of water at least to a mole fraction of about 0.8 based on Figure 9 where an onset for an IL structural breakdown is at best only indicated with the most water rich studied sample at xw = 0.9. It has become fairly established from a number of experimental and theoretical studies that neat ILs structurally organize themselves in polar and nonpolar domains.45−50 Overall, our findings support the work by others indicating that the addition of water to the IL leads to a gradual break-up of the IL network.51−53 Our observation for both ILs [C4mim][MeSO3] and [C2mim][MeSO3] that their principle structural architecture structure remains intact upon addition of water to water mole fractions exceeding well beyond xw = 0.5 implies that the IL remains highly aggregated in the presence of such significant amounts of water. On the other hand, it has been shown that the surface tension of water reduces upon addition of ILs and the formation of aggregates of ILs in aqueous solution is indicated, which is in line with our observations.25,54 Finally, extrapolation of the various composition dependent properties to zero water content allowed us to obtain property values for the neat IL [C4mim][MeSO3] at 85 °C. These and the estimate of the neat IL density at 85 °C are for convenience summarized in Table 5 with standard deviations that were estimated from the fit uncertainties.

Figure 10. D+η/T as a function of water mole fraction for the water− [C4mim][MeSO3] binary system at 85 °C: □, impure IL; ∇, pure IL. A gradual decrease of D+η/T with increasing water mole fraction is indicated.

D=

kT cπηr

(5)

should be proportional to 1/r of the diffusing particle, where in eq 5 r is the hydrodynamic radius of the diffusing particle, and k is the Boltzmann constant. As one would expect that upon dilution in water the IL should gradually break up into smaller fragments and/or ion pairs and, eventually at sufficiently dilute concentrations, into dissociated ions, the ratio plotted in Figure 10 should increase with increasing water content. Clearly, the opposite trend is observed in Figure 10, further supporting the H



Article

Notes

Table 5. Summary of Physicochemical Properties for Neat [C4mim][MeSO3] at 85 °C and Ambient Pressure specific heat density fluidity molar conductivity cation self-diffusion coefficient anion self-diffusion coefficient

■

Csp/J·g−1·K−1 ρ/g·cm−3 η−1/mPa−1·s−1 Λ/10−4 S·m2·mol−1 D+/10−11 m2·s−1 D−/10−11 m2·s−1

1.92 1.1343 0.032 1.57 5.48 4.89

± ± ± ± ± ±

The authors declare no competing financial interest.

■

0.04 0.0023 0.002 0.85 0.57 0.70

ACKNOWLEDGMENTS We thank Dr. Flammersheim for his help with differential scanning calorimetry, Mrs. Petra Weiss for acquiring the DSC data, and Ms. Grit Sauer for acquiring the viscosity data of the impure IL samples. The authors thank Prof. B. Ondruschka for kindly providing the infrastructure during M.M.H.'s sabbatical leave.

■

CONCLUSIONS New experimental data on a number of physicochemical properties for the water−[C4mim][MeSO3] binary system, as well as extrapolated data for the neat IL were presented. The obtained data are qualitatively very similar to what we found in an earlier study on the water−[C2mim][MeSO3] binary system, and combined they are strongly suggestive that the principle structural architecture of water−[C4mim][MeSO3] stays intact upon addition of water even for compositions where the water molecules outnumber the IL molecules. Comparison to the investigation of the water−[C2mim][MeSO3] system20 shows that for neither property unexpected differences occur, and that the behavior upon dilution follows a similar pattern, taking into account the somewhat higher viscosity of [C4mim][MeSO3]. The effect of the acidic impurities in terms of comparing measurement outcomes from using IL with different purity grades are essentially (and surprisingly) absent for the specific heats, conductivity, density, and self-diffusion coefficients and 1 H NMR chemical shifts of cation, anion and water. However, the (acid) impurities present in the impure IL caused larger heats of dissolution and moved the water proton chemical shift to higher values. In addition, it is well-known that the viscosity is very sensitive to some impurities.42 However, this fact was not found in the data presented here due to the counterbalancing effects of lower measurement temperature (viscosity increasing) vs lower purity (viscosity decreasing) of the impure sample. As for why the selectivity of reactions in ILs often depends on a specific water concentration (e.g., from at xw > 0.5 a lower selectivity is observed),17−19 no final conclusion can be drawn from these results: one could have expected decisive changes of the properties at the relevant molar ratios (xw ≈ 0.5), increasing the activity of water, but this is clearly not the case.

■

ASSOCIATED CONTENT

■

AUTHOR INFORMATION

REFERENCES

(1) Sawant, A. D.; Raut, D. G.; Darvatkar, N. B.; Salunkhe, M. M. Recent Developments of Task-Specific Ionic Liquids in Organic Synthesis. Green Chem. Lett. Rev. 2011, 4, 41−54. (2) Isambert, N.; Duque, M. d. M. S.; Plaquevent, J.-C.; Genisson, Y.; Rodriguez, J.; Constantieux, T. Multicomponent Reactions and Ionic Liquids: A Perfect Synergy for Eco-Compatible Heterocyclic Synthesis. Chem. Soc. Rev. 2011, 40, 1347−1357. (3) Ahmed, E.; Ruck, M. Chemistry of Polynuclear Transition-Metal Complexes in Ionic Liquids. Dalton Trans. 2011, 40, 9347−9357. (4) Strehmel, V. Ionic Liquids in Polymer Synthesis. Chem. Eng. Technol. 2011, 83, 1443−1453. (5) Toma, S.; Sebesta, R.; Meciarova, M. Organocatalytic Reactions under Unusual Conditions. Curr. Org. Chem. 2011, 15, 2257−2281. (6) Steinrueck, H. P.; Libuda, J.; Wasserscheid, P.; Cremer, T.; Kolbeck, C.; Laurin, M.; Maier, F.; Sobota, M.; Schulz, P. S.; Stark, M. Surface Science and Model Catalysis with Ionic Liquid-Modified Materials. Adv. Mater. 2011, 23, 2571−2587. (7) Vollmer, C.; Janiak, C. Naked Metal Nanoparticles from Metal Carbonyls in Ionic Liquids: Easy Synthesis and Stabilization. Coord. Chem. Rev. 2011, 255, 2039−2057. (8) Kempe, K.; Becer, C. R.; Schubert, U. S. Microwave-Assisted Polymerizations: Recent Status and Future Perspectives. Macromolecules 2011, 44, 5825−5842. (9) Malik, M. A.; Hashim, M. A.; Nabi, F. Ionic Liquids in Supported Liquid Membrane Technology. Chem. Eng. J. 2011, 171, 242−254. (10) Opallo, M.; Lesniewski, A. A Review on Electrodes Modified with Ionic Liquids. J. Electroanal. Chem. 2011, 656, 2−16. (11) Poole, C. F.; Poole, S. K. Ionic Liquid Stationary Phases for Gas Chromatography. J. Sep. Sci. 2011, 34, 888−900. (12) Shiddiky, M. J. A.; Torriero, A. A. J. Application of Ionic Liquids in Electrochemical Sensing Systems. Biosens. Bioelectron. 2011, 26, 1775−1787. (13) Siedlecka, E. M.; Czerwicka, M.; Stolte, S.; Stepnowski, P. Stability of Ionic Liquids in Application Conditions. Curr. Org. Chem. 2011, 15, 1974−1991. (14) Wilpiszewska, K.; Spychaj, T. Ionic Liquids: Media for Starch Dissolution, Plasticization and Modification. Carbohydr. Polym. 2011, 86, 424−428. (15) Mora-Pale, M.; Meli, L.; Doherty, T. V.; Linhardt, R. J.; Dordick, J. S. Room Temperature Ionic Liquids as Emerging Solvents for the Pretreatment of Lignocellulosic Biomass. Biotechnol. Bioeng. 2011, 108, 1229−1245. (16) Osichow, A.; Mecking, S. Alkoxycarbonylation of Ethylene with Cellulose in Ionic Liquids. Chem. Commun. 2010, 46, 4980−4981. (17) Seddon, K. R.; Stark, A. Selective Catalytic Oxidation of Benzyl Alcohol and Alkylbenzenes in Ionic Liquids. Green Chem. 2002, 4 (2), 119−123. (18) Stark, A. Ionic Liquid Structure-Induced Effects on Organic Reactions. Top. Curr. Chem. 2009, 290, 41−81. (19) Stark, A.; Ondruschka, B. Ionic Liquids in the Manufacture of 5Hydroxymethylfurfural from Saccharides. An Example of the Conversion of Renewable Resources to Platform Chemicals; Wiley-VCH: Weinheim, Germany, 2010; pp 93−121. (20) Stark, A.; Zidell, A. W.; Hoffmann, M. M. Is the Ionic Liquid 1Ethyl-3-Methylimidazolium Methanesulfonate [Emim][MeSO3] Capable of Rigidly Binding Water? J. Mol. Liq. 2011, 160, 166−179.

S Supporting Information *

Data shown in the figures is available in tabulated form. This material is available free of charge via the Internet at http:// pubs.acs.org. Corresponding Author

*E-mail: mhoff[email protected]. Phone: 585-395-5598. Fax: 585-395-5805. Funding

M.M.H. acknowledges support from the American Chemical Society (PRF 46578-UFS) and The College at Brockport for a sabbatical leave at the Friedrich Schiller University Jena, Germany. The instruments on which the density and viscosity measurements were obtained were purchased from a 2009 NSF RUI grant (Award No. 0842960). A.S. is indebted to the DFG for funding (STA 1027/2-2) within the priority programme SPP 1191 on Ionic Liquids. I



Article

(21) Ionic Liquids Database- (IL Thermo) NIST Standard Reference Database #147. http://ilthermo.boulder.nist.gov/ILThermo/ mainmenu.uix (10/31/2011). (22) Claudio, A. F. M.; Freire, M. G.; Freire, C. S. R.; Silvestre, A. J. D.; Coutinho, J. A. P. Extraction of Vanillin Using Ionic-Liquid-Based Aqueous Two-Phase Systems. Sep. Purif. Technol. 2010, 75, 39−47. (23) Ventura, S. P. M.; Neves, C. M. S. S.; Freire, M. G.; Marrucho, I. M.; Oliveira, J.; Coutinho, J. A. P. Evaluation of Anion Influence on the Formation and Extraction Capacity of Ionic-Liquid-Based Aqueous Biphasic Systems. J. Phys. Chem. B 2009, 113 (27), 9304−9310. (24) Ventura, S. P. M.; Sousa, S. G.; Freire, M. G.; Serafim, L. S.; Lima, A. S.; Coutinho, J. A. P. Design of Ionic Liquids for Lipase Purification. J. Chromatogr., B: Anal. Technol. Biomed. Life Sci. 2011, 879, 2679−2687. (25) Russo, J. W.; Hoffmann, M. M. Measurements of Surface Tension and Chemical Shift on Several Binary Mixtures of Water and Ionic Liquids and Their Comparison for Assessing Aggregation. J. Chem. Eng. Data 2011, 56, 3703−3710. (26) Rivera-Rubero, S.; Baldelli, S. Surface Characterization of 1Butyl-3-Methylimidazolium Br−, I−, PF6−, BF4−, (CF3SO2)2N−, SCN−, CH3SO3−, CH3SO4−, and (CN)2N− Ionic Liquids by Sum Frequency Generation. J. Phys. Chem. B 2006, 110 (10), 4756−4765. (27) Santos, C. S.; Rivera-Rubero, S.; Dibrov, S.; Baldelli, S. Ions at the Surface of a Room-Temperature Ionic Liquid. J. Phys. Chem. C 2007, 111, 7682−7691. (28) Lungwitz, R.; Spange, S. A Hydrogen Bond Accepting (HBA) Scale for Anions, Including Room Temperature Ionic Liquids. New J. Chem. 2008, 32, 392−394. (29) Lungwitz, R.; Strehmel, V.; Spange, S. The Dipolarity/ Polarizability of 1-Alkyl-3-Methylimidazolium Ionic Liquids as Function of Anion Structure and the Alkyl Chain Length. New J. Chem. 2010, 34, 1135−1140. (30) Sellin, M.; Ondruschka, B.; Stark, A. Hydrogen Bond Acceptor Properties of Ionic Liquids and Their Effect on Cellulose Solubility. ACS Symp. Ser. 2010, 1033, (Cellulose Solvents) 121−135. (31) Ab Rani, M. A.; Brant, A.; Crowhurst, L.; Dolan, A.; Lui, M.; Hassan, N. H.; Hallett, J. P.; Hunt, P. A.; Niedermeyer, H.; PerezArlandis, J. M.; Schrems, M.; Welton, T.; Wilding, R. Understanding the Polarity of Ionic Liquids. Phys. Chem. Chem. Phys. 2011, 13 (37), 16831−16840. (32) Wu, Y. C.; Koch, W. F.; Pratt, K. W. Proposed New Electrolytic Conductivity Primary Standards for KCl Solutions. J. Res. Nat. Inst. Stand. Technol. 1991, 96 (2), 191−201. (33) Colin, E.; Clarke, W.; Glew, D. N. Evaluation of the Thermodynamic Functions of Aqueous Sodium Chloride from Equilibrium and Calorimetric Measurements Below 154 °C. J. Phys. Chem. Ref. Data 1985, 14 (2), 489−610. (34) Guan, W.; Wang, H.; Li, L.; Zhang, Q.-G.; Yang, J.-Z. Enthalpy of Solution of the Ionic Liquid BMIBF4 in Water. Thermochim. Acta 2005, 437 (1−2), 196−197. (35) Höhne, G. W. H.; Hemminger, W.; Flammersheim, H.-J. Differential Scanning Calorimetry, 2nd ed.; Springer: Berlin, Germany, 2003. (36) Jerschow, A.; Mueller, N. 3D Diffussion-Ordered TOCSY for Slowly Diffusing Molecules. J. Magn. Reson. A 1996, 123, 222−225. (37) Jerschow, A.; Mueller, N. Suppression of Convection Artifacts in Stimulated-Echo Diffusion Experiements. Double-Stimulated-Echo Experiments. J. Magn. Reson. 1997, 125, 372−375. (38) Tokuda, H.; Tsuzuki, S.; Susan, M. A. B. H.; Hayamizu, K.; Watanabe, M. How Ionic Are Room-Temperature Ionic Liquids? An Indicator of the Physicochemical Properties. J. Phys. Chem. B 2006, 110 (39), 19593−19600. (39) Seddon, K. R.; Stark, A.; Torres, M. J. Viscosity and Density of 1-Alkyl-3-Methylimidazolium Ionic Liquids. In Clean SolventsAlternative Media for Chemical Reactions and Processing; Abraham, M. A., Moens, L., Eds.; ACS Symposium Series 819; American Chemical Society: Washington DC, 2002; Vol. 819, pp 34−49. (40) Lide, D. R., CRC Handbook of Chemistry and Physics, 83rd ed.; CRC Press: Boca Raton, FL, 2003.

(41) Blesic, M.; Swadzba-Kwasny, M.; Belhocine, T.; Gunaratne, H. Q. N.; Lopes, J. N. C.; Gomes, M. F. C.; Padua, A. A. H.; Seddon, K. R.; Rebelo, L. P. N. 1-Alkyl-3-Methylimidazolium Alkanesulfonate Ionic Liquids, [CnH2n+1mim][CkH2k+1SO3]: Synthesis and Physicochemical Properties. Phys. Chem. Chem. Phys. 2009, 11 (39), 8939− 8948. (42) Seddon, K. R.; Stark, A.; Torres, M.-J. Influence of Chloride, Water, and Organic Solvents on the Physical Properties of Ionic Liquids. Pure Appl. Chem. 2000, 72 (12), 2275−2287. (43) Yasaka, Y.; Wakai, C.; Matubayasi, N.; Nakahara, M. Water as an in Situ NMR Indicator for Impurity Acids in Ionic Liquids. Anal. Chem. 2009, 81 (1), 400−407. (44) Xu, W.; Cooper, E. I.; Angell, C. A. Ionic Liquids: Ion Mobilities, Glass Temperatures, and Fragilities. J. Phys. Chem. B 2003, 107, 6170−6178. (45) Turton, D. A.; Hunger, J.; Stoppa, A.; Hefter, G.; Thomann, A.; Walther, M.; Buchner, R.; Wynne, K. Dynamics of Imidazolium Ionic Liquids Form a Combined Dielectric Relaxation and Optical Kerr Effect Study: Evidence Got Mesoscopic Aggregation. J. Am. Chem. Soc. 2009, 131, 11140−11146. (46) Rodriguez, V.; Grondin, J.; Adamietz, F.; Danten, Y. Local Structure in Ionic Liquids Investigated by Hyper-Rayleigh Scattering. J. Phys. Chem. B 2010, 114 (46), 15057−15065. (47) Greaves, T. L.; Kennedy, D. F.; Kirby, N.; Drummond, C. J. Nanostructure Changes in Protic Ionic Liquids (PILs) through Adding Solutes and Mixing Pils. Phys. Chem. Chem. Phys. 2011, 13 (30), 13501−13509. (48) Wang, Y.; Voth, G. A. Tail Aggregation and Domain Diffusion in Ionic Liquids. J. Phys. Chem. B 2006, 110 (37), 18601−18608. (49) Wang, Y.; Voth, G. A. Unique Spatial Heterogeneity in Ionic Liquids. J. Am. Chem. Soc. 2005, 127 (35), 12192−12193. (50) Padua, A. A. H.; Costa Gomes, M. F.; Canongia Lopes, J. N. A. Molecular Solutes in Ionic Liquids: A Structural Perspective. Acc. Chem. Res. 2007, 40 (11), 1087−1096. (51) Zhang, L.; Xu, Z.; Wang, Y.; Li, H. Prediction of the Solvation and Structural Properties of Ionic Liquids in Water by TwoDimensional Correlation Spectroscopy. J. Phys. Chem. B 2008, 112 (20), 6411−6419. (52) Hanke, C. G.; Lynden-Bell, R. M. A Simulation Study of WaterDialkylimidazolium Ionic Liquid Mixtures. J. Phys. Chem. B 2003, 107 (39), 10873−10878. (53) Rollet, A.-L.; Porion, P.; Vaultier, M.; Billard, I.; Deschamps, M.; Bessada, C.; Jouvensal, L. Anomalous Diffusion of Water in [Bmim][TFSI] Room-Temperature Ionic Liquid. J. Phys. Chem. B 2007, 111, 11888−11891. (54) Singh, T.; Kumar, A. Aggregation Behavior of Ionic Liquids in Aqueous Solutions: Effect of Alkyl Chain Length, Cations, and Anions. J. Phys. Chem. B 2007, 111 (27), 7843−7851.

J


Composition Dependent Physicochemical Property Data for the Binary

Recommend Documents