J. Phys. Chem. B 2004, 108, 19451-19457
19451
Mixing Schemes in Ionic Liquid-H2O Systems: A Thermodynamic Study Hideki Katayanagi† Center for Frontier Electronics and Photonics, Chiba UniVersity, Chiba 263-8522, Japan
Keiko Nishikawa‡ and Hideki Shimozaki‡ Graduate School of Science and Technology, Chiba UniVersity, Chiba 263-8522 Japan
Kumiko Miki§,| Department of Liberal Arts and Basic Sciences, College of Industrial Technology, Nihon UniVersity, Narashino, Chiba 275-8575, Japan, and Department of Life Science and Chemistry, Roskilde UniVersity, Roskilde DK-4000, Denmark
Peter Westh| Department of Life Science and Chemistry, Roskilde UniVersity, Roskilde DK-4000, Denmark
Yoshikata Koga*,⊥ Department of Chemistry, The UniVersity of British Columbia, VancouVer, British Columbia, Canada V6T 1Z1 ReceiVed: May 24, 2004; In Final Form: September 27, 2004
We studied the hydration characteristics of room-temperature ionic liquids (IL). We experimentally determined the excess chemical potentials, µEi , the excess partial molar enthalpies, HEi , and the excess partial molar entropies SEi in IL-H2O systems at 25 °C. The ionic liquids studied were 1-butyl-3-methylimidazolium tetrafluoroborate ([bmim]BF4) and the iodide ([bmim]I). From these data, the excess (integral) molar enthalpy E and entropy, HEm and SEm, and the IL-IL enthalpic interaction, HIL-IL , were calculated. Using these thermodynamic data, we deduced the mixing schemes, or the “solution structures”, of IL-H2O systems. At infinite dilution IL dissociates in H2O, but the subsequent hydration is much weaker than for NaCl. As the concentration of IL increases, [bmim]+ ions and the counteranions begin to attract each other up to a threshold mole fraction, xIL ) 0.015 for [bmim]BF4 and 0.013 for [bmim]I. At still higher mole fractions, IL ions start to organize themselves, directly or in an H2O-mediated manner. Eventually for xIL > 0.5-0.6, IL molecules E form clusters of their own kind, as in their pure states. We show that HIL-IL , a third derivative of G, provided E E finer details than Hi and Si , second derivatives, which in turn gave more detailed information than HEm and SEm, first derivative quantities.
Introduction The unique nature of ionic liquids (IL) has drawn much attention of late.1-4 Unlike such a typical ionic compound as NaCl, ILs have low melting points and remain liquidlike at about room temperature. Thus, they are used for special reaction media for various organic syntheses and for nonaqueous media for electrochemistry.1-4 Furthermore, their negligible vapor pressures render them suitable as green chemistry solvents.1-4 They also invoke an academic interest, where answers are sought for why ionic compounds are liquidlike at moderate temperatures. To this end, structural investigations have just recently begun.5-9 * Corresponding author: Telephone: (604) 822-3491. Fax: (604) 8222847. E-mail:
[email protected]. † Center for Frontier Electronics and Photonics, Chiba University. ‡ Graduate School of Science and Technology, Chiba University. § Nihon University. | Roskilde University. ⊥ The University of British Columbia.
We measured excess chemical potentials, µEi , and excess partial molar enthalpies, HEi , in aqueous solutions of 1-butyl3-methylimidazolium tetrafluoroborate ([bmim]BF4), and the iodide ([bmim]I), where i stands for IL or H2O (W). Hence, the excess partial molar entropies, SEi , were obtained. The special nature of H2O and aqueous solutions is well documented.10-12 Thus, a study of the mixture of IL and H2O may shed some light on the equally unique properties of IL. Furthermore, there is a practical interest in aqueous solutions of IL. Namely, ILs tend in some cases to be too viscous to be useful as a reaction medium. However, their mixtures with H2O show reduced viscosity without jeopardizing their advantages as green chemistry solvents.13 In other applications, however, extensive efforts have been directed to remove H2O from IL for certain reactions.1-4 Hence, the thermodynamic information about ILH2O systems is important. Thermodynamic studies of the mixtures of IL with other liquids including H2O have also just begun.14-16
10.1021/jp0477607 CCC: $27.50 © 2004 American Chemical Society Published on Web 11/18/2004
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Using the data of HEi and SEi , we calculated numerically three levels of thermodynamic quantities: [1] The excess (integral) molar enthalpy and entropy of the entire system, HEm and SEm calculated as
HEm ) xIL HEIL + xW HEW
(1a)
SEm ) xIL SEIL + xW SEW
(1b)
The subscript i is now replaced with IL and W signifying IL and H2O, respectively. [2] The excess partial molar enthalpy and entropy themselves, HEi and SEi . E , graphically [3] What we call the enthalpic interaction, Hi-i 10-12 calculated as E ≡N Hi-i
( )
( )
∂HEi ∂HEi ) (1 - xi) ∂ni ∂xi
(2)
where ni and xi are respectively the amount and the mole fraction of i in the mixture and N is the total amount. The partial differentiation is performed by keeping other variables constant. The level [1] quantities contain the first derivative of Gibbs energy, G, with respect to temperature, for example, SEm ) -(∂GEm/∂T). HEi and SEi , the level [2] quantities have one more derivative with respect to ni, that is, SEi ) -(∂2GE/∂T ∂ni), the E , eq 2 above, contains yet another second derivative of G. Hi-i derivative with respect to ni. Hence, it is a third derivative quantity. We show below what can be deduced regarding the mixing schemes, or the solution “structures”, using the thermodynamic data at each level described above. We thus demonstrate that finer details of the mixing schemes are obtained from higher-order derivative quantities. Experimental Section [bmim]BF4 was used as supplied (Solvent Innovation GmbH, >98%). [bmim]I was donated by Professor H. Hamaguchi. Because of their hygroscopic nature, ILs were evacuated at room temperature for a few days to remove possible contaminant H2O prior to use. Their vapor pressures became less than 0.001 Torr. This translates into a water content of xW < 5 × 10-5, using our vapor pressure data with the assumption of Henry’s law behavior. The excess partial molar enthalpies were determined using a TAM-2277 titration calorimeter. The uncertainty is (0.03 kJ mol-1 except for sporadic scatters by unknown causes. Typically, a few microliters of a component was titrated into about 1 mL of a solution in the cell. This ratio of the titrant to the solution in the cell was shown to be sufficiently small to ensure an acceptable approximation to the partial molar enthalpy determination.17,18 The total vapor pressures of the solutions were determined by a static method using a Barathron gauge with a sensitivity of (0.001 Torr.19 The composition of the sample solution was varied by adding known amounts of H2O from the gas-handling manifold. The temperature of the sample cell was kept constant within (0.001 K for the duration of measurement. The dayto-day variation of the controlled temperature was observed as (0.05 K, the effect of which was corrected for by the GibbsKonovalov relation.19,20 A small difference of the control set point from 25 °C was also corrected similarly to obtain vapor pressure data at 25 °C. Data Analysis and Results Excess Chemical Potentials, µEi . Since the vapor pressures of pure IL were observed to be zero within the sensitivity of
Figure 1. Excess chemical potential, excess partial molar enthalpy and entropy in [bmim]BF4-H2O at 25 °C. (a) For H2O: µEW (circle), HEW (square), and TSEW (triangle). (b) For [bmim]BF4: µEBF (circle), HEBF (square), and TSEBF (triangle).
our measurement, we assume that the total vapor pressure is equal to the partial pressure of H2O, pW. Hence, the excess chemical potential of H2O is immediately calculated as
µEW ) RT ln
( ) pW
xWp0W
(3)
where p0W is the vapor pressure of pure H2O at 25 °C. We assume that the correction term due to nonideality of the gas phase is negligibly small in comparison with the value calculated by the right side of eq 3. Here, we take the pure state as the reference. The mole fraction of H2O is xW. The results for [bmim]BF4-H2O and [bmim]I-H2O are plotted in Figures 1a and 2a. The uncertainty is estimated as (0.05 kJ mol-1. The excess chemical potential of [bmim]BF4, µEBF, and that for [bmim]I, µEID, were calculated using the µEW data by the GibbsDuhem relation
xW δµEW + xIL δµEIL ) 0
(4)
We chose the symmetric reference state. Hence, µEBF ) µEID ) 0 at xBF ) xID ) 1. The subscripts BF and ID refer to [bmim]BF4 and [bmim]I, respectively, and IL stands collectively for ionic liquids. Because of the uncertainty in µEW, in the H2O-rich region, for xIL < 0.07 or so, the increment δ µEW in eq 4 becomes comparable or smaller than the estimated uncertainty, (0.05 kJ mol-1, which is magnified with the large ratio, xW/xIL
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J. Phys. Chem. B, Vol. 108, No. 50, 2004 19453
Figure 3. Excess (integral) molar quantities for [bmim]BF4-H2O at 25 °C. GEm (circle), HEm (square), and TSEm (triangle).
The second equality of eq 6 is due to the Gibbs-Duhem relation. Hence,
∑(δHEIL) + constant ) ∑[δh/{1 - (xIL/xW)/r}] + constant
HEIL )
Figure 2. Excess chemical potential, excess partial molar enthalpy, and entropy in [bmim]I-H2O at 25 °C. (a) For H2O: µEW (circle), HEW (square), and TSEW (triangle). (b) For [bmim]I: µEID (circle), HEID (square), and TSEID (triangle).
in the H2O-rich region. Thus, in this range, µEIL could not be determined with confidence. The results are shown in Figures 1b and 2b. Excess Partial Molar Enthalpies, HEi . The excess partial molar enthalpy of H2O, HEW, was determined in the standard manner17,18,21 by titrating pure H2O into the mixture. The excess partial molar enthalpy of IL, HEIL, can be calculated by the Gibbs-Duhem relation as above, with the same difficulty in the H2O-rich region. To circumvent this difficulty, we titrated, instead of pure IL, an already diluted aqueous IL into H2O in the calorimetric cell. This is also to avoid another difficulty associated with the high viscosity of pure IL, which seems to prevent an accurate delivery due to some rheological effects. Conversion of these results, using diluted titrant to HEIL, is straightforward.22,23 Briefly, the enthalpy change δq is determined by titrating a diluted titrant consisting of δnIL(t) and δnW(t) into the cell containing nIL of IL and nW of H2O. The superscript (t) signifies the respective quantity of the titrant. The quotient, (δq/δnIL(t)), is written as
h ≡ (δq/δnIL(t)) ) HEIL + HEW/r - ((t)HEIL + (t)HEW/r)
(5)
where r ) δnIL(t)/δnW(t). The increment in h for the successive data points can be written as
δh ) δHEIL + δHEW/r ) δHEIL{1 - (xIL/xW)/r}
(6)
(7)
The uncertainty in HEIL thus obtained is (0.05 kJ mol-1. The constant term on the right of eq 7 is determined if HEIL is known at one point. For this purpose, the calculated values of HEIL by the Gibbs-Duhem relation using the HEW data in the H2O-poor region was used. The excess partial molar entropies were then calculated from the excess chemical potential and excess partial molar enthalpy data. All these level [2] quantities are shown Figures 1 and 2. The level [1] quantities, HEm and SEm, are numerically calculated by eq 1 and shown in Figure 3 for [bmim]BF4, and in Figure 4 for [bmim]I. Also shown are GEm ) HEm - TSEm. Discussion Physical Meanings of Thermodynamic Quantities at Each Level. We briefly review the information contained in the thermodynamic quantities at each level.10-12 The level [1] quantities, HEm and SEm, show the total enthalpy and entropy of the entire system relative to those of the constituents in their pure states. The same is true for GEm ) HEm - TSEm. The level [2] quantities, HEi and SEi , are formally defined as
HEi ) SEi )
( ) ( ) ∂HE ∂ni
(8a)
∂SE ∂ni
(8b)
Hence, they are the response of the system in terms of enthalpy and entropy when perturbed by the infinitesimal increase in ni. They therefore indicate the actual contribution of i, or the actual thermodynamic situation of i in the mixture, in terms of enthalpy and entropy. The net results of these two, µEi ) HEi - TSEi , is
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Figure 4. Excess (integral) molar quantities for [bmim]I-H2O at 25 °C. GEm (circle), HEm (square), and TSEm (triangle). E also included at this level. Hi-i , the level [3] quantity defined by eq 2 above, shows the effect of the infinitesimal increase in ni on HEi , the actual enthalpic situation of i in the solution. E Thus, it signifies the i-i enthalpic interaction. Si-i can be E E here, defined analogous to Hi-i. We did not evaluate Si-i however, since SEi were calculated using two measured quantities, HEi and µEi . Hence, the uncertainty in SEi is about twice that of HEi , which was directly measured. Equivalent Thermodynamic Quantities for NaCl-H2O. We now compare the thermodynamic behaviors of the present IL-H2O systems with those of the aqueous solution of a typical ionic compound, NaCl. Our earlier study showed that an Na+ and Cl- ion pair hydrates 7-8 molecules of H2O and leaves the bulk H2O outside the hydration shell mostly unchanged from the pure H2O state.24,25 A first principle simulation study also indicated that the hydration shell of Na+ contains 5.2 molecules of H2O and there is no effect of Na+ on the orientational dynamics of H2O beyond the hydration shell.26 We use this knowledge in comparing the thermodynamic behaviors of ILs and NaCl. We calculate thermodynamic quantities in aqueous NaCl, treating it as a compound NaCl, as we did for IL above. NaCl is generally believed to dissociate completely in H2O. If Na+ and Cl- ions are perfectly independent species, the thermodynamic treatment must reflect this. Namely, the system has to be treated as a ternary Na+-Cl--H2O system. Hence, the mixing entropy would have an extra term. However, recent theoretical studies suggested some pairing of Na+ and Clions: direct-contact ion pairs and H2O-separated ion pairs.27,28 Furthermore, for the present IL-H2O systems, it is not known if and to what extent ILs dissociate in aqueous solution, although substantial dissociation is expected at a very dilute range. Therefore, we tentatively treat IL as a single entity. To facilitate comparison, we also treat NaCl as a single entity. For this purpose, we use the tabulated data by Clarke and Glew.29 They critically compiled all existing data of a total of 2428 measurements up to 1985 and summarized in the extended Pitzer-type parametrically linear equation. While the fit is impressive, we did not calculate the excess chemical potential of NaCl, µENC, in the very H2O-rich region, xNC < 0.005, for the
Figure 5. Excess chemical potential, excess partial molar enthalpy, and entropy in NaCl-H2O at 25 °C. (a) For H2O: µEW (circle), HEW E E (square), and TSEW (triangle). (b) For NaCl: µNC (circle), HNC (square), E and TSNC (triangle).
same reason as for IL-H2O mentioned above. We take pure solid NaCl as the reference as we did for IL-H2O above. Hence, µENC (saturation) ) -RT ln xNC(saturation). The relative excess partial molar enthalpies of NaCl given in ref 29 are converted for the pure solid NaCl reference state, by HENC(saturation) ) 0 at xNC(saturation). The tabulated data of the excess partial molar enthalpy of H2O are given in the same reference. Thus, the excess partial molar entropies were calculated, and all these excess partial molar quantities are plotted in Figure 5, treating NaCl as an entity in the mixture. Figure 6 shows HEm and SEm, level [1] data, for NaCl-H2O calculated by eq 1. We now compare IL-H2O systems with NaCl-H2O in terms of thermodynamic quantities at each level. Level [1] Quantities, HEm, SEm, and GEm. Comparing Figures 3 and 4 with 6, it is striking that the net thermodynamic effects for IL-H2O systems are much larger, 10- to 20-fold, than those for NaCl-H2O. On mixing NaCl into H2O, (a) NaCl pair breaks away from the crystal and then (b) settles in H2O by hydration. The thermodynamic effects of each step must be very large. For example, the process, NaCl(s) f Na+ + Cl-, costs an enthalpy of about 330 kJ mol-1.30 Thus, the second process, the hydration for NaCl-H2O, must gain about the same amount of enthalpy for the net results of about 0.05 kJ mol-1 as seen in Figure 6. Similar compensation must also be working for the net results of 0.1 kJ mol-1 in TSEm. In comparison, Figures 3 and 4 show a larger enthalpy loss, 2 kJ mol-1 for [bmim]BF4H2O and about 1 kJ mol-1 for [bmim]I-H2O. The entropy gains amount to 1.8 kJ mol-1 for the former and 1.4 kJ mol-1 for the
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J. Phys. Chem. B, Vol. 108, No. 50, 2004 19455
Figure 6. Excess (integral) molar quantities for [bmim]BF4-H2O at 25 °C. GEm (circle), HEm (square), and TSEm (triangle).
latter. While the ionic bonding in IL could be weaker than that of NaCl, IL must also be dissociated in ions at least at infinite dilution. Therefore, the above observation suggests that the second process of hydration is not as complete as that of NaCl, perhaps due to the size effects or IL’s hydrophobicity. In comparing [bmim]BF4-H2O, Figure 3, and [bmim]I-H2O, Figure 4, the general trend for the latter is closer to NaClH2O, that is, TSEm > HEm, and hence, GEm < 0. This may reflect the hydration of the I- ion. We note that the recent experimental data of HEm for [bmim]BF4-H2O at 60 and 5 °C16 seem consistent with our data at 25 °C, Figure 3. Level [2] Quantities, HEi , SEi , and µEi . We now turn to a comparison of the excess partial molar quantities. Comparison of Figures 1a, 2a, and 5a indicates that the thermodynamic state of H2O in the IL-rich region is more drastically altered than that in NaCl-H2O near saturation. As mentioned above, for NaCl-H2O, the bulk H2O away from the hydration shell remains almost the same as in pure liquid H2O.24,25 Figure 1a shows that for [bmim]BF4-H2O, HEW > TSEW with a net positive µEW. For [bmim]I-H2O on the other hand, HEW < TSEW, the behavior is closer to NaCl-H2O though the values are different. These could be related to hydration of I-, the detail of which is yet to be investigated. It is striking in Figures 1b and 2b that HEIL and TSEIL are almost zero with the net µEIL also zero in the IL-rich region, xIL > 0.5-0.6. This means that in this composition range, IL molecules are in an environment very similar to that of their pure states. Namely, they exist as clusters of their own kind, with a size not large enough to form a separate phase. In this range, therefore, IL molecules are not dissociated, and the system may be more appropriately described as aqueous nonelectrolytes in the H2O-poor region.10-12 Instead of precipitating out as for NaCl-H2O, IL molecules remain in solution as aqueous nonelectrolytes. How H2O interacts with these IL clusters should manifest itself in the behavior of HEW and SEW. Figures 1a and 2a indicate that such interactions are different between [bmim]BF4-H2O and [bmim]I-H2O, though the details of interaction are not immediately obvious from these data alone.
Figure 7. Enthalpic interaction function for [bmim]BF4-H2O at 25 E E °C. (a) Between ionic liquids, HBF-BF . (b) Between H2O, HW-W .
In the H2O-rich region in Figures 1b and 2b, HEIL data show a large enthalpy loss, 18 kJ mol-1 for [bmim]BF4-H2O and 10 kJ mol-1 for [bmim]I-H2O at the infinite dilution. It is unfortunate that we could not evaluate µEIL and hence, TSEIL data for xIL < 0.05; however, at about xIL ) 0.1, the entropy gains amount to about 7 kJ mol-1 as TSEIL for both. In our recent X-ray diffraction study on the pure liquid [bmim]I,5 the basic feature of the structural arrangement of I- remains almost the same as that of Cl- in the solid [bmim]Cl.9 Hence, for the first step process of an IL molecule breaking away from its pure environment, there is a large enthalpy loss and entropy gain, though the absolute values may not be as large as for NaCl. However, depending on the degree of dissociation, the subsequent interaction with liquid H2O, direct hydration or “iceberg” formation around IL ions, would be weaker, resulting in the net results of the enthalpy loss and entropy gain of the magnitude observed. If it is hydration, it is much weaker than NaCl. While the entropy gains are about the same, the fact that enthalpy loss for [bmim]I is about 8 kJ mol-1 smaller than that for [bmim]BF4 may reflects hydration of I-. A recent study on [bmim]BF4-H2O indicated the existence of a phase separation with the upper critical solution temperature (UCST) at about 4 °C and 0.07 in the mole fraction.16 Figure 1b shows that in this composition range, both HEBF and TSEBF decrease on increasing xBF. This observation, together with the fact that µEBF is positive and also decreases, is consistent with a thermodynamic property of the system that has a phase separation with a UCST, as we argued at some length.10,31
19456 J. Phys. Chem. B, Vol. 108, No. 50, 2004
Katayanagi et al. start to organize themselves. Nonetheless, it seems clear from analogy of the studies of aqueous nonelectrolytes10-12 that some change in mixing scheme is occurring at this boundary. Indeed, there is an observation that the surface tension takes the minimum at xBF ) 0.015 for [bmim]BF4-H2O at 25 °C.33 HEW data were also directly measured in the range, xIL > 0.1. E We therefore evaluated HW-W in the same manner. The results are shown in Figures 7b and 8b. Both show weak anomalies at about xIL ) 0.5 to 0.6, which may be the boundary beyond which IL molecules form clusters. This corresponds to the saturation for NaCl-H2O. For IL-H2O, however, the system stays mixed as aqueous solutions of nonelectrolytes in the H2O-poor region.10-12 Namely, instead of precipitating out as for NaCl, IL molecules remain in solution as clusters, and this no doubt reflects a unique nature of IL. Conclusion On dissolution into H2O at infinite dilution, IL molecules break away from their pure environment and settle in the H2O environment presumably as ions. However, subsequent hydration of the ions, if any, is much weaker than for NaCl-H2O.24-26 As the concentration of IL increases, IL ions and the counteranions begin to interact with each other, either directly or in an H2O-mediated manner. At about xIL ) 0.015, some organization among the ions takes place. Eventually at about xIL ) 0.5 to 0.6 IL molecules cluster together with a very similar local arrangement as in the pure IL state. For [bmim]I-H2O, however, some effect of hydration of I- is evident. We point out that finer details of the above interpretation were possible by the composition dependence of higher-order derivative quantities.
Figure 8. Enthalpic interaction function for [bmim]I-H2O at 25 °C. E E . (b) Between H2O, HW-W (a) Between ionic liquids, HID-ID .
It is noteworthy that for both IL, the enthalpy loss and the entropy gain sharply diminish as xIL increases. This could suggest that some enthalpically favorable and entropically unfavorable interaction may be setting in, which could be an increasing direct or H2O-mediated attraction between [bmim]+ and the respective counteranion. Thus, these ions are beginning to organize themselves, and eventually for xIL > 0.5-0.6, IL molecules form clusters of their own kind mentioned above. E . In the most H2O-rich region, Level [3] Quantity, Hi-i xIL < 0.07, we measured HEIL data directly and accurately. On close inspection, the values of HEIL do not decrease monotoniE (i ) IL), would show cally as xIL increases. Equation 2, Hi-i this trend more clearly. In studies of aqueous solutions of E nonelectrolytes, Hi-i among other third derivatives of the Gibbs energy, G, has provided important information in elucidating the detailed nature of aqueous solutions.10-12 Here, E for both IL-H2O, by we also evaluate graphically10-12 HIL-IL eq 2. The method of graphical differentiation without sacrificing much precision was discussed at some length earlier.32 The results are plotted in Figures 7a and 8a. It is evident from the figures that something drastic occurs at xBF ) 0.015 for [bmim]BF4-H2O and at xID ) 0.013 for [bmim]I-H2O. In our earlier E and in studies in aqueous solutions, these anomalies in Hi-i other third derivatives of G, were found associated with the transition of mixing scheme.10-12 For the present IL-H2O systems, the details of the mixing schemes on either side of anomalies are yet to be elucidated. A possible explanation could be that at this point, [bmim]+ and the respective counteranion
Acknowledgment. We thank Prof. H. Hamaguchi for donating the [bmim]I sample, and Prof. Y. Ouchi for sharing their unpublished data. We also thank Prof. L. P. N. Rebelo for directing us to ref 16. This research was supported by Ministry of Education, Science, and Culture, Nihon University, Danish Research Council, and the Carlsberg Foundation. Supporting Information Available: Table S-1 showing the vapor pressure data. This material is available free of charge via the Internet at http://pubs.acs.org. References and Notes (1) Wasserscheim, P.; Welton, T.; Eds. Ionic Liquids in Syntheses; VCH-Wiley: Weinheim, 2003. (2) Jessop, P. G. J. Synth. Org. Chem. Jpn. 2003, 61, 483-488. (3) Welton, T. Chem. ReV. 1999, 99, 2071. (4) Jessop, P. G.; Stanley, R. R.; Brown, R. A.; Echert, C. A.; Lietta, C. L.; Ngo, T. T.; Pollet, P. Green Chem. 2003, 5, 123-128. (5) Katayanagi, H.; Hayashi, S.; Hamaguchi, H.; Nishikawa, K. Chem. Phys. Lett., 2004, 392, 460-464. (6) Ozawa, R.; Hayashi, S.; Saha, S.; Kobayashi, A.; Hamaguchi, H. Chem. Lett. 2003, 32, 948-949. (7) Saha, S.; Hayashi, S.; Kobayashi, A.; Hamaguchi, H. Chem. Lett. 2003, 32, 740-741. (8) Hayashi, S.; Ozawa, R.; Hamaguchi, H. Chem. Lett. 2003, 32, 498499. (9) Holbrey, J. D.; Reichert, W. M.; Nieuwenhuyzen, M.; Johnson, S.; Seddon, K. R.; Rogers, R. D. Chem. Commun. 2003, 1636. (10) Koga, Y. Netsusokutei (J. Jpn. Cal. Therm. Anal.) 2003, 30, 5465. Available in pdf format (from the author,
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