Synthesis and Characterization of Thiazolium-Based Room

Aug 10, 2012 - Department of Chemistry, University of Tennessee, Knoxville, ... Rong An , Yudan Zhu , Nanhua Wu , Wenlong Xie , Jiawei Lu , Xin Feng ,...
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Synthesis and Characterization of Thiazolium-Based Room Temperature Ionic Liquids for Gas Separations Patrick C. Hillesheim,† Shannon M. Mahurin,*,† Pasquale F. Fulvio,† Joshua S. Yeary,† Yatsandra Oyola,† De-en Jiang,† and Sheng Dai*,†,‡ †

Chemical Sciences Division, Oak Ridge National Laboratory, Oak Ridge, Tennessee 37831, United States Department of Chemistry, University of Tennessee, Knoxville, Tennessee 37996, United States

‡

ABSTRACT: A series of novel thiazolium-bis(triflamide) based ionic liquids has been synthesized and characterized. Physicochemical properties of the ionic liquids such as thermal stability, phase transitions, and infrared spectra were analyzed and compared to the imidazolium-based congeners. Several unique classes of ancillary substitutions are examined with respect to impacts on overall structure, in addition to their carbon dioxide absorption properties in supported ionic-liquid membranes for gas separation. organic catalysts,18,19 and transition metal-containing catalysts.20 Given the large interest in the development of novel ILs,21 and to better understand the mechanisms of CO2 separation by these compounds, herein we report the synthesis and characterization of six new thiazolium-based ILs, which are shown in Figure 1. By maintaining the same 4-methyl-3-

1. INTRODUCTION The field of ionic liquids (ILs) has received great attention since the first reports on air and moisture stable compounds.1 Boasting many favorable properties, such as wide electrochemical windows, high thermal stabilities, negligible vapor pressures, and highly modular core structures, ILs have been successfully applied to homogeneous catalysis,2 biomass processing,3 carbon capture,4 solvent extraction,5 and energetic materials applications.6 Among the various classes of ILs, imidazolium based ILs have been one of the most largely investigated.7 Of recent interest, is the application of ILs to the field of carbon capture, specifically CO 2 separation 3 in supported membranes. 8 To date, imidazolium-based ILs,6−9 and most recently pyrrolidinium-, pyridinium-, ammonium-, and phosphonium-based ILs,9−11 have been the motifs of choice for displaying high CO2 solubility. Current experimental and computational studies on these ILs suggest that their physical properties, and consequently gas solubility, are dictated by numerous factors including the structure of the anion and cation and the interaction between the two ionic species as well as the constitution of the alkyl side chains.12 It has been proposed that the anions and cations of ILs tend to form 3-dimensional ionic networks.12 For instance, the alkyl side chains attached to the imidazolium cores, having three or more carbon atoms, segregate into hydrophobic domains. Furthermore, the properties of the latter are further influenced by both the anion and the cation core structures.12 The molecular conformation of such ionic networks, and of their hydrophobic domains, results from the charge densities, molecular asymmetry, and flexibility of the anion and the cation cores. However, while imidazolium based systems have been the main focus of such studies, other five-membered heterocycles capable of forming ILs have been neglected. First reported by Davis and Forrester,13 thiazolium-based ILs comprise a subclass with unique electronic and structural properties, which have been poorly investigated to date.14−16 Only a few examples of room temperature thiazolium-based ILs have been recently investigated for their biodegradability,17 as © 2012 American Chemical Society

Figure 1. Molecular structures of the thiazolium-based ionic liquids.

thiazolium cores, the side chain tethered to the nitrogen in these compounds was systematically modified in order to increase its hydrophobicity or add new functional alcohol or ester groups (compounds 1a−1d). Many of the functional groups can be used to enhance CO2 capture.22,23 In addition, the same rationale was applied to an even more substituted core structure, 5-(2-hydroxyethyl)-4-methylthiazolium, resulting in compounds 2a and 2b. In general, higher molecular weight substitutions in the C3-position resulted in lower thermal stability, higher glass transition temperatures, and higher viscosity. On the other hand, lower molecular weight side Received: Revised: Accepted: Published: 11530

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Temperature dependent viscosity measurements were acquired using a cone/plate viscometer (Brookfield, DVII+ Pro), which allowed the use of small sample volumes of approximately 1 mL. The sample was maintained at the desired temperature with a recirculating water bath (VWR) for at least 15 min prior to measurement, and five separate measurements at each temperature were obtained and averaged. Viscosity standards (Cannon) were used to calibrate the instrument before ionic liquid viscosities were measured. 2.4. Synthesis of Ionic Liquids. Synthesis of 3-Butyl-4methylthiazol-3-ium Bis((Trifluoromethyl)Sulfonyl)Amide (1a). To a round-bottom flask containing 1 eq of 4methylthiazole cooled in an ice bath was added 1.2 eq of 1bromobutane dropwise. After the addition was complete, the ice bath was removed, the flask was fitted with a condenser, and the reaction was heated to 60 °C overnight. The reaction was then allowed to cool to room temperature, and the unreacted reagents were decanted from the viscous oil. Ethyl acetate was added to the oil, and the mixture was stirred for ten minutes during which an off-white solid formed. The solid was filtered, washed with ethyl acetate (3 × 100 mL) and ether (3 × 100 mL), and dried under vacuum overnight yielding 3-butyl-4methylthiazol-3-ium bromide which was used as is without further purification. A 10.0 g (42.3 mmol) sample of 3-butyl-4methylthiazol-3-ium bromide was dissolved in 100 mL of water, activated carbon was added, and the mixture was heated at 70 °C overnight. The resultant clear solution was cooled to room temperature and filtered. A 14.6 g aliquot of lithium bis(trifluoromethanesulfonimide) (50.9 mmol, 1.2 eq) was added, and the mixture was stirred overnight. The aqueous layer was decanted, the remaining layer was dissolved in dichloromethane and washed with water (6 × 100 mL), the solvent was removed in vacuo, and the resultant clear liquid was dried under high vacuum for 4 days. Yield: 14.16 g (77%). 1H NMR (Methanol-d4, 400 MHz): δ = 9.97 (d, J = 2.7 Hz, 1 H), 7.90 (br. s., 1 H), 4.46 (t, J = 1.0 Hz, 2 H), 2.62 (s, 3 H), 1.93 (quin, J = 7.7 Hz, 2 H), 1.46 (dq, J = 15.0, 7.5 Hz, 2 H), 1.02 ppm (t, J = 7.5 Hz, 3 H). 13C NMR (Methanol-d4, 400 MHz): δ = 159.3, 148.2, 126.1, 122.9, 122.6, 119.7, 116.5, 54.3, 32.6, 20.6, 13.8, 13.3 ppm. ESI-MS: (+ve) 156.08 m/z [C8H15N1S1]+; ESI-MS: (-ve) 279.923 m/z [C2F6N1O4S2]−. Synthesis of 3-Benzyl-4-methylthiazol-3-ium Bis((trifluoromethyl)sulfonyl)amide (1b). Compound 1b was synthesized using benzyl bromide following the same procedure as for 1a. Yield: 13.27 g (76%). 1H NMR (Methanol-d4, 400 MHz): δ = 9.86 (s, 1 H), 7.94 (s, 1 H), 7.43 - 7.52 (m, 3 H), 7.36 (d, J = 7.3 Hz, 2 H), 5.70 (s, 2 H), 2.53 ppm (s, 3 H). 13C NMR (Methanol-d4, 101 MHz): δ = 159.7, 148.4, 133.4, 130.8, 130.8, 129.7, 126.1, 123.1, 122.9, 119.7, 116.5, 57.7, 13.6 ppm. ESI-MS: (+ve) 190.08 m/z [C8H15N1S1]+; ESI-MS: (-ve) 279.93 m/z [C2F6N1O4S2]−. Synthesis of 3-(2-Ethoxy-2-oxoethyl)-4-methylthiazol-3ium Bis((trifluoromethyl)sulfonyl)amide (1c). Compound 1c was synthesized using ethyl bromoacetate following the same procedure as for 1a. Yield: 15.89 g (91%). 1H NMR (Methanol-d4, 400 MHz): δ = 10.03 (s, 1 H), 7.95 (s, 1 H), 5.45 (m, 1 H), 4.32 (q, J = 7.2 Hz, 2 H), 2.54 (s, 3 H), 1.33 ppm (t, J = 7.2 Hz, 3 H). 13C NMR (Methanol-d4, 101 MHz): δ = 165.5, 147.2, 124.6, 121.4, 120.9, 118.2, 115.0, 62.8, 12.9, 11.6 ppm. ESI-MS: (+ve) 186.07 m/z [C8H12N1O2S1]+; ESIMS: (-ve) 279.93 m/z [C2F6N1O4S2]−. Synthesis of 3-(2-Hydroxyethyl)-4-methylthiazol-3-ium Bis((trifluoromethyl)sulfonyl)amide (1d). Compound 1d was

chains at the same position favored higher CO2 permeability and CO2/N2 selectivity, regardless of the presence of oxygen functionalities.

2. EXPERIMENTAL SECTION 2.1. Chemicals. 4-Methylthiazole was purchased from Matrix Scientific. Lithium bis(trifluoromethanesulfonimide) was purchased from 3M. All other chemicals were purchased from Sigma-Aldrich in the highest available purity and used as received. 3-Butyl-1-methyl-1H-imidazol-3-ium bis((trifluoromethyl) sulfonyl)amide) [Bmim(Tf2N)],24 3-benzyl-1-methyl-1H-imidazol-3-ium bis((trifluoromethyl)sulfonyl)amide [BzMim(Tf2N),25 and 3-(2-hydroxyethyl)-1-methyl-1Himidazol-3-ium bis((trifluoromethyl)sulfonyl)amide24 were synthesized according to established literature procedures. 2.2. Spectroscopy. FT-IR analysis was carried out on a Digilab FTS 7000 FT-IR spectrometer using an attenuated total reflectance (ATR) attachment with a diamond crystal. 1H and 13C NMR spectra were recorded on a Bruker Avance 400, at 400 MHz and 100 MHz, respectively. A JEOL (Peabody, MA) JMS-T100LC (AccuTOFTM) orthogonal time-of-flight (TOF) mass spectrometer was used to characterize the compound. 2.3. Physical Properties. Thermal stabilities were measured using a Thermal Advantage 2950 thermogravimetric analyzer (TGA) with platinum pans under flowing nitrogen atmosphere and 10 °C/min as heating rate from room temperature up to 750 °C. Melting points and glass transition temperatures were calculated from differential scanning calorimetry (DSC) measurements using a Thermal Advantage Q50 instrument using hermetically sealed aluminum pans. Each IL was subjected to three consecutive heating−cooling cycles from −90 to 200 °C, with 10 °C/min as heating rate. Gas solubility measurements were obtained using a gravimetric microbalance (Hiden Isochema, IGA). Approximately 70 mg of the ionic liquid was loaded into a quartz sample boat and sealed in the stainless steel chamber. Prior to the solubility measurement, the ionic liquid was dried and degassed at a temperature of 60 °C and a vacuum pressure of 1 mbar until the measured mass no longer decreased (typically for a minimum of 4 h). Weight measurements were then acquired at various CO2 pressures up to 10 atm at a constant temperature of 25 °C, which was maintained using a constanttemperature recirculating water bath. Gas permeability was measured by loading each ionic liquid onto a porous polyethersulfone support with a pore diameter of 100 nm and measuring gas flux through the ionic liquid membrane. The permeability measurement system and loading procedure have been previously described in detail.8 The permeability was calculated from the gas flux through the membrane using the following equation P=

τ V dP ″ · φ RTA ·ΔP0 dt

(1)

where τ is the tortuosity, ϕ is the membrane porosity, V is the permeate volume, R is the ideal gas constant, T is the absolute temperature, A is the membrane area, ΔP0 is the pressure difference, and dP′′/dt is the rate of pressure increase on the permeate side of the membrane. The tortuosity and porosity were determined using Emim(Tf2N) as a standard. CO2/N2 selectivity values were obtained by separately measuring the CO2 and N2 permeability and calculating the ratio of these values to obtain the ideal selectivity. 11531

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synthesized using 2-bromoethanol following the same procedure as for 1a. Yield: 10.1 g (53%). 1H NMR (Acetone-d6, 400 MHz): δ = 10.10 (d, J = 2.6 Hz, 1 H), 8.07 (br. s, 1 H), 4.76−4.81 (m, 2 H), 4.55 (br. s., 1 H), 4.06− 4.15 (m, 2 H), 2.73 ppm (s, 3 H). 13C NMR (Acetone-d6, 101 MHz): δ = 160.0, 148.1, 125.8, 122.6, 121.8, 119.4, 116.2, 60.4, 56.2, 13.6 ppm. ESI-MS: (+ve) 144.05 m/z [C6H10N1O1S1]+; ESI-MS: (-ve) 279.93 m/z [C2F6N1O4S2]−. Synthesis of 3-Ethyl-5-(2-hydroxyethyl)-4-methylthiazol-3ium Bis((trifluoromethyl)sulfonyl)amide (2a). A 10.0 g (39.6 mmol) aliquot of 3-ethyl-5-(2-hydroxyethyl)-4-methylthiazol-3ium bromide was dissolved in 100 mL of water to which activated carbon was added. The mixture was stirred overnight at 70 °C, allowed to cool to room temperature, and subsequently filtered. Lithium bis(trifluoromethanesulfonimide) (13.6 g, 47.5 mmol, 1.2 eq) was added, and the mixture was stirred overnight. The mixture was allowed to settle, and the aqueous layer was decanted. The resultant clear liquid was dissolved in ethyl acetate and washed with water (6 × 100 mL). The solvent was removed in vacuo, and the clear liquid was dried under high vacuum for 4 days. Yield: 10.4 g (58%) 1H NMR (Acetone-d6, 400 MHz): δ = 9.96 (s, 1 H), 4.66 (q, J = 7.2 Hz, 2 H), 4.40 (t, J = 4.9 Hz, 1 H), 3.87 (q, J = 5.2 Hz, 2 H), 3.19 (t, J = 5.6 Hz, 2 H), 2.64 (s, 3 H), 1.65 ppm (t, J = 7.3 Hz, 3 H). 13C NMR (Acetone-d6, 101 MHz): δ = 155.6, 143.0, 137.1, 125.7, 122.6, 119.4, 116.1, 61.1, 49.9, 30.6, 14.9, 11.8 ppm. ESI-MS: (+ve) 172.09 m/z [C8H14N1O1S1]+; ESI-MS: (-ve) 279.93 m/z [C2F6N1O4S2]−. Synthesis of 3-Benzyl-5-(2-hydroxyethyl)-4-methylthiazol3-ium Bis((trifluoromethyl)sulfonyl)amide (2b). Compound 2b was synthesized using 3-benzyl-5-(2-hydroxyethyl)-4-methylthiazol-3-ium bromide following the same procedure as for 2a. Yield: 12.4 g (65%). 1H NMR (Acetone-d6, 400 MHz): δ = 9.95 (s, 1 H), 7.40−7.54 (m, 5 H), 5.90 (s, 2 H), 4.43 (t, J = 4.9 Hz, 1 H), 3.88 (q, J = 5.3 Hz, 2 H), 3.20 (t, J = 5.6 Hz, 2 H), 2.57 ppm (s, 3 H). 13C NMR (Acetone-d6,101 MHz): δ = 156.9, 143.3, 137.8, 133.3, 130.4, 130.3, 129.3, 125.8, 122.6, 119.4, 116.2, 61.0, 57.7, 30.6, 12.2 ppm. ESI-MS: (+ve) 235.10 m/z [C 13 H 16 N 1 O 1 S 1 ] + ; ESI-MS: (-ve) 279.93 m/z [C2F6N1O4S2]−.

Figure 2. FTIR spectra of the thiazolium-based ionic liquids, compound 1a−d (top) and compound 2a−b (bottom). Spectra were vertically offset in increments of 150% for clarity.

distinguishing the two compounds occurs in 2b at 702 cm−1, which may arise from the benzyl moiety tethered to the nitrogen. The spectrum for compound 1b, also bearing an nbenzyl moiety, has a band at the same frequency lending further evidence to this analysis. Overall, the thiazolium-based ILs display band frequencies similar to those reported in the literature. Given the limited FTIR data available on thiazolium ILs, full analysis of the spectra requires further corroboration via computational studies.25 3.2. Physical Properties. Thermogravimetric (TG) and differential thermogravimetric (DTG) curves of all six compounds are shown in Figure 3. The majority of the compounds display a sharp decomposition step within the temperature range of 270 to 370 °C, whereas the compounds with longer side chains bound to the nitrogen in the thiazolium ring exhibit more complex decomposition profiles as well as the lowest initial decomposition temperatures. For instance, compound 1a shows a weight loss smaller than 2 wt % at 250 °C, followed by a 98 wt % weight loss at 354 °C without any residues left. The TG curve for compound 2a shows 2 overlapping decomposition steps, the first starting at 364 °C and the second at 380 °C. Compound 1d displays a single, sharp decomposition beginning at 344 °C and negligible carbon residue remaining. As for compounds 1b, 1c, and 2b, multiple

3. RESULTS AND DISCUSSION 3.1. Fourier Transform Infrared Spectroscopy (FTIR). FTIR spectra of the compounds, see Figure 2, were acquired in an effort to understand some of the structural characteristics of the cations and anions of the ILs. The spectra of the thiazolium ILs exhibit bands resembling those characteristic of imidazolium-based ILs.26−32 Additionally, the FTIR spectra of compounds 1a−1d show some distinction between these four compounds. The broad bands occurring between 3100 and 3120 cm−1 are likely due to the C−H stretch at the C2-position of the thiazole ring.30,31 The slight shift to lower wavenumbers is indicative of more acidic hydrogen at the C2-position. This concept is also evidenced by the central hydrogen appearing further downfield in the NMR spectra as compared to imidazolium based congeners.29 The characteristic functional group bands are readily observable for compounds 1c and 1d, with the expected O−H peak for 1d occurring at 3530 cm−1 and the ester carbonyl peak for 1c showing at 1747 cm−1. Analysis of compounds 2a and 2b shows virtually no change in spectra (Figure 2b). This indicates that substituents bound to the nitrogen of the thiazole ring seem to have little effect on the overall structure of the ionic liquid. The only significant band 11532

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to the thiazolium-based ILs. While the imidazolium based ILs did display higher Tdec, compounds 1a and 1d only differed by approximately 30 °C. The largest change in stability came from 1b, which decomposed at nearly 100 °C less than its imidazolium congener. Similarly compound 2b, also bearing the benzyl substituted nitrogen, showed significantly lower thermal stability as compared to the structurally similar compound 2a. Further experiments into the exact decomposition pathways will be required to properly determine the apparent effect benzyl groups are imparting with regard to their lower thermal stabilities. In addition to acquiring the decomposition temperatures for the ionic liquids, long-term thermal stability was examined for compounds 1a and 1b in an effort to examine the effects of aryl as compared to alkyl substitutions on structural stability. Despite displaying high decomposition temperatures, ionic liquids often exhibit weight loss when held at elevated temperatures for extended periods of time.2 Such long-term tests often provide additional understanding of the stability and nature of the compounds tested. Both compounds 1a and 1b were subjected to temperatures ranging from 150 to 250 °C for 6 h periods at each temperature increment under a N2 atmosphere. The imidazolium congeners, namely Bmim(Tf2N) and BzMim(Tf2N), were similarly subjected to the tests as comparisons. The results of the tests are summarized in Table 1. The initial decomposition curve for 1a was similar to that Table 1. Percentage of Weight Loss Per Hour (%wt/h) from Extended Thermal Stability Studies on Compounds 1a and 1b As Compared with Bmim(Tf2N) and BzMim(Tf2N)a

Figure 3. TG curves for the thiazolium-based ionic liquids (a) and corresponding DTG curves (b).

significant thermal decomposition steps are observed. For compound 1b, the major weight loss occurs at 271 °C corresponding to an 82 wt % change followed by a subsequent broader decomposition occurring between 500 and 515 °C. The TG curve of compound 1c shows two distinct steps, the first at 309 °C, and a secondary at approximately 368 °C, both accounting for a 94 wt % loss. Compound 2b displays three distinct decompositions. The first major decomposition begins at 296 °C, followed by two smaller decompositions at 345 and 421 °C. Compared to similarly substituted imidazole based ILs, thiazolium-based ILs display more complex decompositions. In contrast to most standard imidazolium ILs that display one-step decomposition, four of the compounds examined herein exhibit multiple decomposition steps at elevated temperatures. These observations might be indicative of thermally induced intermediary products occurring within thiazolium ILs. To better understand the thiazolium-based IL systems, the imidazolium-based congeners of compounds 1a, 1b, and 1d were synthesized and the decomposition temperatures recorded for each compound. The imidazole containing ILs, namely 3butyl-1-methyl-1H-imidazol-3-ium bis((trifluoromethyl) sulfonyl)amide [Bmim(Tf2N)] (Tdec 387 °C), 3-benzyl-1methyl-1H-imidazol-3-ium bis((trifluoromethyl)sulfonyl)amide [BzMim(Tf2N)] (Tdec 367 °C), and 3-(2-hydroxyethyl)-1methyl-1H-imidazol-3-ium bis((trifluoromethyl)sulfonyl)amide (Tdec 379 °C), all display higher thermal stability as compared

a

compd

150 °C

180 °C

1a 1b Bmim(Tf2N) BzMim(Tf2N)

0.32 0.055 2.2e−3 1.2e−3

0.037 0.21 0.022 1.6e−3

200 °C 220 °C 250 °C 0.031 1.6 0.084 0.019

0.11 7.1 0.26 0.085

0.66 4.7 1.2 0.58

total weight loss* 6.9 82.9 9.3 4.1

The * refers to weight lost over entire duration of test (30 h, %wt).

observed in the TG discussed above, with a gradual decrease in weight accounting for a 0.32%wt/h loss at 150 °C. This weight loss is in stark contrast to that observed in Bmim(Tf2N) which displayed a mere 0.0022%wt/h loss during the same period of time. Interestingly, at temperatures greater or equal to 200 °C compound 1a displayed lower weight loss than that observed in Bmim(Tf2N). For the entirety of the long-term thermal tests, 1a lost a total of approximately 6.9 wt %, while Bmim(Tf2N) exhibited a loss of 9.3 wt %. Compound 1b and BzMim(Tf2N) were also subjected to the same series of testing. For the entire duration of the test, compound 1b displayed a total weight loss of 82.9% as compared to BzMim(Tf2N), which exhibited a loss of 4.1%. While 1b exhibited nominal weight loss at lower temperatures, 200 °C. It can be concluded from the data gathered by both the standard TG curves and the long-term thermal analysis that the addition of benzyl groups has a significant effect on the thermal stability of thiazolium-based ILs. This observation is in contrast to what is observed for BzMim(Tf2N) compared to Bmim(Tf2N) where substitution of a benzyl ring for an alkyl chain increased the long-term thermal stability. 11533

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Table 2. Physicochemical Parameters for the Thiazolium-Based Ionic Liquids and Calculated Gas Absorption Dataa compd

MW (g/mol)a

D (g/cm3)b

Tg (°C)c

Tdec (°C)d

η (cP)e

PermCO2 (barrer)f

SelCO2/N2g

SolCO2 (mol/L·atm)h

KH,px (atm)i

1a 1b 1c 1d 2a 2b Emim[Tf2N]*

436.63 484.67 466.61 424.57 452.60 528.72 391.31

1.505 1.548 1.562 1.626 1.540 1.504 1.515

−68.1 −46.5 −46.9 −64.4 −40.9 −62.8 −8741

354.5 271.0 309.6 344.8 364.5 296.2 279.942

129 ± 2 443 ± 5 613 ± 9 216 ± 6 212 ± 3 1292 ± 10 3243

362 ± 17 235 ± 6 248 ± 5 284 ± 12 435 ± 20 96 ± 5 1702.437

28 ± 4 21 ± 2 18 ± 3 25 ± 2 36 ± 4 12 ± 3 23.137

0.0964 0.0869 0.0986 0.0902 0.0858 0.0720 0.10344

37 41 35 44 40 45 37.7

a − molecular weight; b − calculated density using COSMOtherm; c - glass transition temperature from DSC; d − decomposition temperature taken from the maximum of the first peak on the DTG curves; e − viscosity at 25°C; f − CO2 permeability at 25 °C; g − CO2 selectivity calculated from the permeability ratios of CO2 to N2; h − CO2 solubility 25 °C; i − Henry’s Law constant calculated according to KH,XP = p/x, where p is the pressure and x is the mol fraction of CO2 in a given IL. *Reported literature values. a

among all compounds. This IL also has the simplest structure, consisting of an alkylthiazolium without added functional groups. Conversely, ILs 1b, 1c, and 2b exhibited the highest viscosities, all having an alkyl benzyl group and/or a hydroxylterminated alkyl chain as in compound 2b. The latter further exhibits the highest viscosity among all at 1292 cP. The added aromatic group and a hydroxyl-terminated alkyl chain significantly increases the viscosity compared to the ionic liquids with short side chains. Interestingly, 1d and 2a, which differed in the position of the hydroxyl-terminated ethyl chain, in addition to an ethyl substitution, exhibited nearly identical viscosities, indicating that the positioning of the hydroxyl-alkyl chain has little effect on this physical property. In addition to the ring substitution effects, thiazolium-based ILs also exhibit higher viscosity values compared to their imidazolium analogues. For example, the recently reported imidazolium IL, BzMim(Tf2N),10 with similar ancillary groups as those in compound 1b has a viscosity of 61 cP at 25 °C, which is 86% less than that of 1b. Similarly, the viscosity of a hydroxyl-terminated alkyl chain imidazolium analogue of compound 1d has been reported to be 91 cP,24 which again is 58% lower than the viscosity of compound 1d. Figure 5b shows the viscosity values fit to the Arrhenius equation

The phase transition of the six ILs has been examined by DSC, the results of which are summarized in Table 2. The compounds displayed consistent DSC profiles over three consecutive cycles. With the exception of compound 1c, only a glass transition was observed for these compounds. Compounds 1b, 1c, and 2a exhibit the highest Tg transitions, above −47 °C. The least substituted compounds 1a and 1d and the most substituted thiazolium IL, namely compound 2b, exhibit the lowest glass transition temperatures, within −68 °C and −62 °C. Furthermore, the melting point transition for compound 1c, shown in Figure 4, happened at 40 °C, with an

ln(η) =

⎛ Ea ⎞ 1 ⎜ ⎟ + ln A ⎝ R ⎠T

(2)

where Ea is the activation energy for viscous flow, R is the universal gas constant, T is the absolute temperature, and A is a pre-exponential factor. The temperature-dependent viscosities show an excellent fit indicating Arrhenius-type behavior, which is similar to the imidazolium-based ionic liquids. 3.3. CO2 Absorption. The CO2 solubility curves are shown in Figure 6, and the calculated parameters are summarized in Table 2. The linear dependence of the CO2 absorption curves at low pressure is in agreement with Henry’s law. This is characteristic of ILs for which the primary mechanism for the CO2 solubilization is physical absorption. From Table 2, the CO2 solubility values vary from 0.0720 to 0.0986 mol·L−1·atm−1, with the highest volumetric CO2 solubility found for 1c and the lowest volumetric solubility occurring for 2b. Both ILs have comparable CO2 solubilities (and Henry’s constant values) to Emim(Tf2N) and C6mim(Tf 2 N), 0.103 mol·L −1 ·atm −1 (39 atm) and 0.0702 mol·L−1·atm−1 (44 atm), respectively.35,36 These values are comparable to 1a which is an alkyl-thiazolium ionic liquid without additional functional groups. Furthermore, the compounds with the highest viscosity values, namely compounds 1b, 1c, and 2b, have the lowest CO2 solubilities.

Figure 4. Representative DSC profile of 3-(2-ethoxy-2-oxoethyl)-4methylthiazol-3-ium bis((trifluoromethyl)sulfonyl)amide (compound 1c) showing the 3 consecutive heating and cooling cycles.

enthalpy of transition (ΔHm) of 51 kJ/mol. The melting point peak was preceded by an exothermic phase transition at approximately 5 °C. While the exact cause of this phase transition is the subject of current investigations, ILs may exhibit unusual phase transitions often attributed to molecular rearrangements or to cold crystallization in which a randomly oriented glassy state is reached prior to melting.33 To understand mass transport processes through ionic liquids, especially those involving gas separations, it is important to quantitatively determine the viscosity. Table 2 shows the viscosities of the ionic liquids at 25 °C, with full temperature-dependent viscosities shown in Figure 5a. For all 6 compounds, the viscosity values are moderately high, especially when compared to Bmim(Tf2N), which has a viscosity value of 50.4 cP at 25 °C.34 Ionic liquid 1a exhibited the lowest viscosity 11534

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Figure 5. a) Temperature dependence for the viscosity of the various synthesized thiazolium-based ionic liquids. b) Arrhenius plot of the thiazoliumbased ionic liquids. The lines correspond to the best fit lines.

Figure 6. High pressure CO2 gas solubility isotherms at 25 °C for the various thiazolium-based ionic liquids.

While reports indicate that fractional free volume, cation−anion interactions, and CO2-anion interactions may all influence CO2 solubility in ionic liquids,12 it is interesting to explore the relationship between viscosity and CO2 solubility. From Figure 7a, which shows the CO2 solubility as a function of viscosity for both the thiazolium- and imidazolium-based ionic liquids, there is a clear decrease in solubility as the viscosity increases for both types of ionic liquids. These results suggest that the thiazolium class of ionic liquids have the potential for CO2 separations in much the same way as the imidazolium ionic liquids. For the CO2 permeability, the thiazolium ILs have lower values as shown in Table 2 than alkyl-imidazolium compounds. For example, Emim(Tf2N) and C6mim(Tf2N) have reported CO2 permeability values of 1702 barrer and 1136 barrer, respectively,37 while the thiazolium compounds reported here all have CO2 permeabilities below 450 barrer. One possible reason for the decrease in permeability is the larger viscosity associated with the thiazolium ILs, which may slow the gas diffusion through the ionic liquid and consequently reduce permeability. Figure 7b shows the CO2 permeability as a function of viscosity for the thiazolium ionic liquids. There is a clear nonlinear trend relating viscosity to gas permeability for these thiazolium compounds in agreement with the viscosity dependence for imidazolium ionic liquids which are also shown in Figure 7b. The inset shows the same data on a log−log plot. In a recent review, Scovazzo noted a similar viscosity dependence for imidazolium, ammonium, and phosphonium

Figure 7. a) CO2 solubility as a function of viscosity for thiazolium (filled squares) and imidazolium (open circles) ionic liquids. b) Plot of permeability vs viscosity for thiazolium-based ionic liquids (filled squares) and imidazolium-based ionic liquids (open circles). The inset shows the same data on a log−log plot.

ILs and reported two different equations that correlated CO2 permeability with viscosity.37 From Table 2, compound 2b showed the lowest CO2 permeabilty and the highest viscosity, whereas the highest CO2 permeabilties were found for the compounds with viscosities lower than 216 cP. Again, these results are consistent with previous reports of gas permeability in ionic liquids. The CO2/N2 selectivity values are similar to those typically observed for the more common imidazolium based ILs. For example, CO2/N2 selectivity values for imidazolium ionic 11535

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liquids such as Emim(Tf2N), Bmim(Tf2N), and C6Mim(Tf2N) that contain the same anion as the thiazolium ILs studied in this work are 23.1, 19.7, and 15.4, respectively.37 The CO2/N2 selectivity for 1a, a thiazolium IL that is analogous to the alkylimidazolium ionic liquids, was measured to be 28 which is higher than the alkyl-imidazolium.10 Interestingly, the selectivity of a benzyl-imidazolium IL analogous to the thiazolium 1b has been reported to be 31.4 which is higher than the selectivity of the benzylthiazolium at 21. However, the viscosity of 1b, the benzyl-thiazolium, was nearly an order of magnitude higher than the benzyl-imidazolium. Moreover, compound 2a, which had the highest permeability among the liquids tested, also exhibited the highest selectivity. Similar to gas permeability, a nonlinear trend was found relating the selectivity to the viscosity of these thiazolium ILs, with the highest selectivity found for the less viscous compounds. Excluding compound 2b which had a very high viscosity, the selectivity values for the thiazolium-based ILs ranged from 18 to 36, suggesting a potential use in gas separations. In general, these thiazolium ionic liquids exhibit comparable CO2 solubility values to most commonly reported alkylimidazolium ionic liquids. Certain functional groups such as carbonyl groups found in 1c are known to have high affinity for CO2. The addition of the functional groups, however, resulted in negligible changes for CO2 solubility.38 A balance between the functional chain substitutions to the thiazolium rings and resulting viscosity of these ILs was crucial to achieve the highest CO2 permeability and CO2/N2 selectivity results.

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4. CONCLUSIONS A series of six novel thiazolium-bis(trifluoromethanesulfonimide) based ionic liquids have been synthesized and their physicochemical properties investigated. FTIR and NMR spectra for these compounds show that thiazolium-based ILs share similar structural properties with their imidazolium based congeners.24,26−33,39,40 While a melting transition was observed by DSC for only one compound, all synthesized ILs are freeflowing liquids at room temperature, effectively broadening the scope of room-temperature ionic liquids. Also, trends associating the latter with the final viscosity and the degree of substitution of the IL cores revealed that further improvements toward gas diffusivity and selectivity can be achieved, thus increasing the potential of thiazolium-based ILs in composite membranes for gas separations.

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Article

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected] (S.D.), [email protected] (S.M.M.). Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS P.C.H. was supported by the U.S. Department of Energy, Advanced Research Projects Agency − ENERGY and the U.S. Department of Energy. S.M.M., D.J., and S.D. were supported by the Division of Chemical Sciences, Geosciences, and Biosciences, Office of Basic Energy Sciences, U.S. Department of Energy. We would like to thank Dr. Michelle Kidder for her assistance with analysis of infrared spectra. 11536

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