LETTER pubs.acs.org/Langmuir
Hydrophilic Quaternary Ammonium Type Ionic Liquids. Systematic Study of the Relationship among Molecular Structures, Osmotic Pressures, and Water-Solubility Koji Kawai,†,‡ Kotaro Kaneko,‡ and Tetsu Yonezawa*,† †
Division of Material Science and Engineering, Faculty of Engineering, Hokkaido University, Kita 13, Nishi 8, Kita-Ku, Sapporo, Hokkaido, 060-8628, Japan ‡ Miyoshi Oil and Fat Co., Ltd, 4-66-1, Horikiri, Katsushika-ku, Tokyo, 124-8510, Japan
bS Supporting Information ABSTRACT: This Letter examines the relationship between the structures of ionic liquids and their water-solubility or osmotic pressure with a number of synthesized quaternary ammonium type ionic liquids and organic salts containing a hydroxyl group as hydrophilic substituted groups on ammonium group cations, and bromide or methylsulfonate as anions. The study found a linear relation between the amount and osmotic pressure of the water-soluble ionic liquids synthesized here, strongly indicating that these water-soluble ionic liquids are perfectly ionized in water like inorganic salts with small diameter ions.
’ INTRODUCTION A room temperature ionic liquid (RTIL) is an organic salt which is in liquid form at room temperature and which has specific physical properties such as noncombustibility, no vapor pressure, high heat resistance, and high ionic conductivity. With such unique physical features, RTILs have attracted interest in a number of scientific fields including electronics, chemistry, analysis, as catalysts, and others.110 For example, an RTIL is used as an organic reaction solvent instead of conventional organic solvents due to better safety and the potential for recycling.4 In lithium-ion rechargeable batteries, the usage of RTIL as the electrolyte is being investigated due to its noncombustibility, nonvolatility, as well as its high conductivity, allowing portability and ensuring the safety of the batteries.5 Further, it is considered to offer promising characteristics as an electrolyte in dye-sensitized solar cells.6 Also, an RTIL is now a highly important agent contributing to reductions in environmental loads. For the applications mentioned here, one requirement is chemical inactivity of the RTILs. Thus, most RTILs do not have reactive groups in the molecular structures. However, quite recently, preparation and applications of functionalized ionic liquids have been reported. For example, a terminal hydroxyl group of quaternary ammonium type RTILs with a hydroxyl group and an alkyl group has been used as a Lewis acid catalyst.7 The application of imidazolium type RTILs with hydroxyalkyl groups has been anticipated to increase in the field of biochemistry and medicine. These can be used as the basic compounds for transdermal drug delivery systems due to their skin permeability,8 cellulose decomposition,9 as well as protein refolding agents.10 To apply RTILs to bio-related fields, it is r 2011 American Chemical Society
necessary to consider their osmotic pressures because cells are covered by cell membranes which are semipermeable. However, there have been few reports on the correlation between ionic liquid and osmotic pressure,11 and especially no report of systematic studies on the relationship among structures, counterions and osmotic pressure has been located. This study systematically synthesized a number of hydrophilic ionic liquids with hydroxyl groups, and the osmotic pressures were measured to discuss the relationships among molecular structure, water solubility, and osmotic pressures.
’ EXPERIMENTAL SECTION Detailed information of the synthesis is described in the Supporting Information, and here one such process will be described: 2-Bromoethanol (TCI) and N,N-dimethylethylamine (Kokusan) were mixed in acetonitrile and kept stirred for 24 h at room temperature. The generated 2-hydroxyethylethyldimethylammonium bromide was obtained as a white solid. Bromide anion was exchanged by mixing with methanesulfonic acid (CH3SO3H, Kanto) or tetrafluoroborate acid (HBF4, Kanto) in water at room temperature and the mixture was kept stirred for 24 h. The corresponding methanesulfonate salt and tetrafluorobrate salt were successfully obtained as products of this process. The products were characterized with the 1H NMR, 13C NMR, and FT-IR spectra, and elemental analyses; the water content, melting temperature, and decomposition temperature of each product were measured with DSC or TG/DTA; and the osmotic pressure was measured Received: March 22, 2011 Revised: May 11, 2011 Published: May 23, 2011 7353
dx.doi.org/10.1021/la201070n | Langmuir 2011, 27, 7353–7356
Langmuir
LETTER
by a freezing-point depression method with an osmotic-pressure meter. The water-solubility of the products was established visually after mixing with water by stirring at 25 °C for 30 min and leaving the solution for 10 min. In some cases, micelles or other aggregated structures of ILs were generated in water. Diameters of such products in aqueous solutions were obtained with a dynamic light scattering method. Some samples adsorbed water molecules in air, and the concentration of the aqueous solutions of the products was calculated by considering the water amount of the neat products obtained by the TG/DTA measurement. We here attempted to synthesize simple structured RTILs to systematically compare the relationship between the molecular structures and osmotic pressures. Two routes have been reported to synthesize quaternary ammonium salts with hydroxyethyl groups: (i) the reaction of a secondary aminoalcohol and an alkylhalide12 and (ii) the reaction of a tertiary amine and 2-bromoethanol.7b,13 The latter route was employed here, because 2-bromoethanol reagent is cheaper and is simpler to handle.
’ RESULTS AND DISCUSSION The definition of ILs requires that they are in liquid form below 100 °C and that RTILs are in liquid form at room temperature.14 None of the bromides synthesized in this study were RTILs when they were dried under reduced pressure. Some bromides (1, 3, 6, and 8) were white solids at room temperature. However, they instantly absorbed H2O after exposure to air to change to transparent liquids like RTILs. The DSC or TG/DTA results of these bromides show their melting points and decomposition points all corresponding to the reported values (90300 °C).7b,15,16 The melting points of 1, 2, 5, and 9 are below 100 °C but higher than room temprature. They are classified in ILs but not in RTILs.17 The obtained ammonium salts with unsymmetrical structures, ethyldimethylhydroxyethyland diethylmethylhydroxyethyl-methylsulfonate (4 and 7, respectively), showed lower melting temperatures than the symmetrical structured methylsulfonates (2 and 9), due to the weaker intermolecular interactions,18 and they are categorized in RTILs. However, the melting points of symmetrical quaternary ammonium bromides (1 and 8) were lower than those of unsymmetrical quaternary ammonium bromides (3 and 6). This difference in the melting points may be explained by the molecular structure of the cations and the bulkiness of the anions. Bromide anions exhibit a higher negative charge density than methylsulfonate anions which had a delocalized electron in the conjugated π-electron orbital and a bulkier structure. Therefore, the interaction between cationic ammonium and bromide anions is very strong and the bromide anion would locate near the cation. With unsymmetrical quaternary ammonium bromides, the location of the small bromide anion is fixed, that is, it is at the nearest position. With symmetrical quaternary ammonium bromides, the location of the bromide anion is flexible, and the melting points of symmetrical ammonium bromides become lower than those of unsymmetrical ammonium bromides. The melting temperatures of the ammonium methylsulfonates (4, 7, and 9) are lower than the melting temperatures or the decomposition points of the corresponding bromides (3, 6, and 8) and a tetrafluoroborate (5). This situation can be explained by the bulkiness of the anions. The 1 (bromide) and 2 (methylsulfonate) products each with a small cation (trimethylammonium) showed similar melting temperatures. Figure 1 is a plot of the osmotic pressures and the concentrations of the ammonium salts and the ILs synthesized here. The methylsulfonate ILs (4 and 7) show a linear relation like inorganic salts with small ions, such as NaCl. The osmotic pressures of these methylsulfonate ILs (4 and 7) were very high
Figure 1. Osmotic pressure of ionic liquids and quaternary ammonium salts. R1R2R3Nþ(CH2)2OH X; (O) R1, R2 = CH3, R3 = C2H5, X = CH3SO3; (b) R1 = CH3, R2, R3 = C2H5, X = CH3SO3; (4) R1, R2 = CH3, R3 = C2H5, X = Br; (2) R1 = CH3, R2,R3 = C2H5, X = Br; (0) R1, R2 = CH3, R3 = C2H5, X = BF4; () Imz BF4.
Figure 2. Sizes of the aggregated structures of compounds 3, 5, 6, and 10 in an aqueous solution (100 mmol of salts/100 g of water) measured by a dynamic scattering method.
compared with those of the bromides (3 and 6) and tetrafluoroborate (5). If these ILs are completely ionized in water, the osmotic pressure at a concentration of 100 mmol of the salts in 100 g of water would be 2 Osmol/kg. However, the osmotic pressures of 4 and 7, as shown in Figure 1, are above 2. The higher values may be attributed to the measurement method as well as to the molecular structures of 4 and 7. The osmotic 7354
dx.doi.org/10.1021/la201070n |Langmuir 2011, 27, 7353–7356
Langmuir
LETTER
Table 1. Melting Point and Water-Solubility of Ionic Liquids and Quaternary Ammonium Salts [R1R2R3N(CH2)2OH]þY compd
Y
mp (°C)
water-solubility (g 100 g water1, mmol 100 g water1)
R1
R2
R3
1
CH3
CH3
CH3
Br
91
2
CH3
CH3
CH3
CH3SO3
96
>900, >4516
3
CH3
CH3
C2H5
Br
285 (d.c.p.)
400500, 20192524
400500, 21732716
282285 (d.c.p.)15a 4
CH3
CH3
C2H5
CH3SO3
87
5
CH3
CH3
C2H5
BF4
97
100200 488976
6
CH3
C2H5
C2H5
Br
275(d.c.p.)
400500, 18862357 >900 200300, 8841327
>900, >4219
7
CH3
C2H5
C2H5
CH3SO3
252 (d.c.p.)15b 75
8
C2H5
C2H5
C2H5
Br
124
>3959 1207b 9
C2H5
10
ImzBF4
C2H5
C2H5
CH3SO3
70
>900, >3729
89
500600, 22122655
7116
pressures were measured by the freezing-point depression method, and when a sample has high viscosity, its osmotic pressure tends to be slightly higher than the expected 2.19 The viscosity of quaternary ammonium type ILs is high in appearance. As both the anions and cations of methylsulfonate ILs (4 and 7) form hydrogen bonds with water molecules, the interaction between the ILs and water molecules is relatively strong. Since methylsulfonate anion is more basic than bromide anion (HBr, pKa = 9; CH3SO3H, pKa = 0.6),20 methylsulfonate can hydrate with water readily compared to bromide. The characteristic molecular structure of the methylsulfonate ILs, a hydrophilic CH3SO3 anion and a cation with a hydroxyl terminal group, must be considered a factor in their high degree of ionization. The obtained osmotic pressures of the bromides (3 and 6) and tetrafluoroborates (5), and 10 (ImzBF4) are much smaller than those of the methylsulfonates, and they do not show linear changes against the interaction. The slope of the curve decreased at 2060 mmol of the salts in 100 g of solution. Generally, breaks in the curves of the osmotic pressure and the concentration occur at critical micelle concentrations (cmc), indicated by the arrows in Figure 1, because the osmotic pressure has a colligative property to the molar concentration of surfactant.21 In order to observe the formation of aggregated structures of the bromides (3 and 6) and tetrafluoroborate (5) in aqueous solutions (100 mmol of salts/ 100 g of water), dynamic light scattering measurement was carried out (Figure 2). The particle diameters of the aggregated molecules were large, because the molecular sizes of the compounds are very small. Such aggregate structure formation of the bromide salts and the tetrafluoroborate salt was also confirmed by the fact that they did not completely dissolve in water at a concentration of ions higher than 2060 mmol of salts/100 g of water. Compared with a general imidazolium type hydrophobic ionic liquid (10, ImzBF4), the synthesized highly hydrophilic alkylammonium ILs (37) showed considerably higher osmotic pressures. The ILs with similar cationic structure, methylsulfonates (4 and 7), showed higher osmostic pressures than the bromides (3 and 6) and the tetrafluoroborate (5). This tendency corresponds well to that of the water-solubility (methylsulfonates > bromides and tetrafluoroborate), as shown in Table 1. It is to be
expected that a compound with a higher ionization ratio shows a higher solubility, and the methylsulfonates are fully ionized in water.
’ CONCLUSION This study synthesized a number of aliphatic hydrophilic ammonium salts and ILs and observes their osmotic pressures and water solubility. The osmotic pressures and degree of electrolytic dissociation of the ionic liquids can be simply controlled and adjusted by molecular design. Among the ammonium salts synthesized here, the salts with an asymmetric hydrophilic cation and methylsulfonate completely dissociated in water and showed the highest osmotic pressure. This suggests that molecular design of ILs can control their permeability through cell membranes. These ILs contain hydrophilic oxyanions and cations which are similar to choline with hydroxyl groups. With dissolution of cellulose with ILs, hydrogen bonds between the ILs and cellulose are important.22 Because anions and cations on aliphatic hydrophilic ammonium can form hydrogen bonds, a contribution to carbon neutrality with biomass would be a possibility. The findings in this study will be useful in a range of fields including the field of nanotechnology. ’ ASSOCIATED CONTENT
bS
Supporting Information. Experimental details of syntheses and characterization of the compounds, water-solubility, and osmotic pressures. This material is available free of charge via the Internet at http://pubs.acs.org.
’ AUTHOR INFORMATION Corresponding Author
* Fax: þ81-11-706-7881. E-mail:
[email protected].
’ ACKNOWLEDGMENT Helpful discussions with Dr. A. Hyono, Dr. S. Abe (Hokkaido Univ.), K. Kamio, K. Sugiyama, and H. Kawakami (Miyoshi Oil and Fat Co., Ltd) are gratefully acknowledged. 7355
dx.doi.org/10.1021/la201070n |Langmuir 2011, 27, 7353–7356
Langmuir
LETTER
’ REFERENCES (1) Yuan, H.; Shimotani, H.; Ye, J.; Yoon, S.; Aliah, H.; Tsukazaki, A.; Kawasaki, M.; Iwasa, Y. J. Am. Chem. Soc. 2010, 132, 18402–18407. (2) Fitzgerald, J. J. D.; Kunnath, P.; Walker, A. V. Anal. Chem. 2010, 82, 4413–4419. (3) Mikkola, J.; Warna, J.; Virtanen, P.; Salmi, T. Ind. Eng. Chem. Res. 2007, 46, 3932–3940. (4) (a) Welton, T. Chem. Rev. 1999, 99, 2071–2083. (b) Earle, M. J.; Seddon, K. R. Pure Appl. Chem. 2000, 72, 1391–1398. (c) Wilkes, J. S. J. Mol. Catal. A 2004, 214, 11–17. (5) (a) Scheers, J.; Johansson, P.; Jacobsson, P. J. Electrochem. Soc. 2008, 155, A628–A634. (b) Noto, V. D.; Negro, E.; Sanchez, J. Y; Iojoiu, C. J. Am. Chem. Soc. 2010, 132, 2183–2195. (6) Armel, V.; Pringle, J. M.; Forsyth, M.; MacFarlane, D. R.; Officer, D. L.; Wagner, P. Chem. Commun. 2010, 46, 3146–3148. (7) (a) Zhu, A.; Jiang, T.; Han, B.; Zhang, J.; Xie, Y.; Ma, X. Green Chem. 2007, 9, 169–172. (b) Sun, J.; Zhang, S.; Cheng, W.; Ren, J. Tetrahedron Lett. 2008, 49, 3588–3591. (8) Moniruzzaman, M.; Tahara, Y.; Tamura, M.; Kamiya, N.; Goto, M. Chem. Commun. 2010, 46, 1452–1454. (9) (a) Kamiya, N.; Matsushita, Y.; Hanaki, M.; Nakashima, K.; Narita, M.; Goto, M.; Takahashi, H. Biotechnol. Lett. 2008, 30, 1037– 1040. (b) Fukaya, Y.; Hayashi, K.; Wada, M.; Ohno, H. Green Chem. 2008, 10, 44–46. (c) Turner, M. B.; Spear, S. K.; Holbrey, J. D.; Daly, D. T.; Rogers, R. D. Biomacromolecules 2005, 6, 2497–2502. (10) (a) Summers, C. A.; Flowers, R. A., II. Protein Sci. 2000, 9, 2001–2008. (b) Lange, C.; Patil, G.; Rudolph, R. Protein Sci. 2005, 14, 2693–2701. (c) Yamaguchi, S.; Yamamoto, E.; Tsukiji, S.; Nagamune, T. Biotechnol. Prog. 2008, 24, 402–408. (11) (a) Gonzalez, B.; Calvar, N.; Dominguez, A.; Macedo, E. A. J. Chem. Thermodyn. 2008, 40, 1346–1351. (b) Shekaari, H.; Mousavi, S. S. J. Chem. Eng. Data 2009, 54, 823–829. (c) Shekaari, H.; Mousavi, S. S. Fluid Phase Equilib. 2009, 279, 73–79. (d) Calvar, N.; Gonzalez, B.; Dominguez, A.; Macedo, E. A. J. Chem. Thermodyn. 2009, 41, 11–16. (e) Shekaari, H.; Armanfar, E. J. Chem. Eng. Data 2010, 55, 765–772. (12) (a) Wilk, K. A. J. Phys. Chem. 1992, 96, 901–904. (b) Dasgupta, A.; Das, D.; Das, P. K. Biochimie 2005, 87, 1111–1119. (13) Lukac, M.; Mojzis, J.; Mojzisova, G.; Mrva, M.; Ondriska, F.; Valentova, J.; Lacko, I.; Bukovsky, M.; Devinsky, F.; Karlovska Eur. J. Med. Chem. 2009, 44, 4970–4977. (14) Endres, F.; Abedin, S. Z. E. Phys. Chem. Chem. Phys. 2006, 8, 2101–2116. (15) (a) Parkkari, J. H.; Bannard, R. A. B.; Coleman, I. W. Can. J. Chem. 1965, 43, 3119–3128.(b) Aktieselskabet Grindstedvaerket. Kongeriget Danmark. Patent DK78457, 1954. (16) Yamamoto, H. J. Comput. Aided Chem. 2006, 7, 18–30. (17) Constantinescu, D.; Schaber, K. J. Chem. Eng. Data 2007, 52, 1280–1285. (18) Dean, P. M.; Pringle, J. M.; MacFarlane, D. R Phys. Chem. Chem. Phys. 2010, 12, 9144–9153. (19) Tsuge, H.; Nakajima, T.; Ohuchi, T.; Aoki, M.; Ohkubo, T.; Yomota, C. Pharm. Regul. Sci. 2009, 40, 136–142. (20) Brownstein, S.; Stillman, A. E. J. Phys. Chem. 1959, 63, 2061–2062. (21) (a) Gonzalez-Perez, A.; Czapkiewicz, J.; Del Castillo, J. L.; Rodriguez, J. R. Colloid Polym. Sci. 2003, 281, 556–561. (b) Jungnickel, C.; Luczak, J.; Ranke, J.; Fernandez, J. F.; Muller, A.; Thoming, J. Colloids Surf., A 2008, 316, 278–284. (22) Zhang, J.; Zang, H.; Wu, J.; Zang, J.; He, J.; Xiang, J. Phys. Chem. Chem. Phys. 2010, 12, 1941–1947.
7356
dx.doi.org/10.1021/la201070n |Langmuir 2011, 27, 7353–7356
