Thermophysical Properties of Aqueous Solution of Ammonium-Based

Article pubs.acs.org/JPCB

Thermophysical Properties of Aqueous Solution of AmmoniumBased Ionic Liquids Reddicherla Umapathi,† Pankaj Attri,‡ and Pannuru Venkatesu*,† †

Department of Chemistry, University of Delhi, Delhi 110007, India Plasma Bioscience Research Centre/Department of Electrical and Biological Physics, Kwangwoon University, Seoul, Korea, 139-791

‡

ABSTRACT: Experimental densities (ρ), ultrasonic sound velocities (u), viscosities (η), and refractive indices (nD) of binary mixtures of ammoniumbased ionic liquids (ILs) such as diethylammonium acetate (DEAA) [(CH 3 CH 2 ) 2 NH][CH 3 COO], triethylammonium acetate (TEAA) [(CH3CH2)3NH][CH3COO], diethylammonium hydrogen sulfate (DEAS) [(CH3CH2)2NH][HSO4], triethylammonium hydrogen sulfate (TEAS) [(CH 3 CH 2 ) 3 NH][HSO 4 ], trimethylammonium acetate (TMAA) [(CH 3 ) 3 NH][CH 3 COO], and trimethylammonium hydrogen sulfate (TMAS) [(CH3)3NH][HSO4] with water are reported over the wide composition range at 25 °C under atmospheric pressure. The excess molar volumes (VE), deviation in isentropic compressibilities (Δκs), deviation in viscosities (Δη) and deviation in refractive indices (ΔnD) are calculated from experimental values and are correlated by Redlich−Kister polynomial equations. The VE and Δκs values for the aforesaid systems are negative over the entire composition range while the Δη and ΔnD values are positive under the same experimental conditions. The intermolecular interactions and structural effects were analyzed on the basis of measured and derived properties. A qualitative analysis of the results is discussed in terms of the iondipole, ion-pair interactions and hydrogen bonding between ILs and water. Furthermore, the hydrogen bonding features between ILs with water were analyzed by using a molecular modeling program with the help of HyperChem7.

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INTRODUCTION The physicochemical properties such as densities (ρ), ultrasonic sound velocities (u), viscosities (η), and refractive indices (nD) of liquids and liquid mixtures are used to identify the behavior of molecular interactions and its nature of structure making or breaking effects in the binary system. These properties can expand the range of structural properties and the scope of the molecular interactions between the solvent molecules. In recent years, ionic liquids (ILs), which are the combinations of ions, have emerged as the promising solvents, are fast growing topic of chemical research on account of their unique properties that include negligible vapor pressure, nonflammability, and good ability to dissolve in organic and inorganic compounds.1−4 The utilization and applications of ILs have been rapidly increased in all scientific fields by several researchers.5−12 Recently, ammonium-based ILs, a subset of protic ionic liquids have shown a wide range of scientific applications in biochemical process,13−15 separation technology,10,16,17 and chemical synthesis.18,19 An adequate knowledge of the thermo-physical properties of ammonium-based ILs with molecular solvents are essentially required to clarify the nature of molecular interactions between these ILs with solvents as well as to design new chemical and technological processes. Recently, we have reported the thermo-physical properties and excess properties of ammonium-based ILs +N,Ndimethylformamide (DMF),20−22 +N,N-dimethyl sulfoxide (DMSO)23−27 and N-methyl-2-pyrrolidone (NMP)28−30 over © 2014 American Chemical Society

the whole composition range. Although, the studies on the thermo-physical properties of ammonium-based ILs with water are essentially required for the design of industrial, simulation process, and theoretical studies. Obviously, these studies for aqueous solutions of ammonium-based ILs yet remain to be understood. An exhaustive literature survey reveals that only a very few researchers have reported the thermo-physical properties of ammonium family of ILs with water.31−34 In 2010, Taib and Murugasen31 explored the ρ and excess molar volumes for bis(2-hydroxyethyl) ammonium acetate with water over the entire composition range. Later, Alvareszet al.32 presented ρ and u data for 2-hydroxy ethylammonium acetate with water throughout the whole concentration range. In addition, Xu research group reported the thermo-physical properties of ethylammonium acetate with water33 and nbutylammonium acetate or n-butylammonium nitrate34 with water over the whole concentration range. Apparently, there are scarce reports, which have presented the measurements of diethylammonium and triethylammonium-based protic ILs with water. Furthermore, a molecular description of interactions of these ILs in aqueous solution is lacking and numerous issues remain unresolved on the thermo-physical properties of these ILs in aqueous solution. Received: March 10, 2014 Revised: May 14, 2014 Published: May 15, 2014 5971

dx.doi.org/10.1021/jp502400z | J. Phys. Chem. B 2014, 118, 5971−5982

The Journal of Physical Chemistry B

Article

Table 1. Solvent Purity, Molecular Mass (MW), Density (ρ), Ultrasonic Sound Velocity (u), Viscosity (η) and Refractive Index (nD) for the Solvents Such as ILs and Water at 25 °C and at Atmospheric Pressure solvent

MW/g mol−1

% purity

ρ/g·cm−3

u/m·s−1

η/mPa·s

nD

water DEAA DEAS TEAA TEAS TMAA TMAS

18.02 133.11 171.06 161.14 199.09 119.05 157.03

99 99 99 99 99 99 99

0.997 04 1.021 41 1.283 95 1.015 86 1.142 89 1.053 85 1.467 58

1496 1608 1432 1840 1874 1544 1564

0.89 16.36 34.10 24.12 235.00 2.76 5.10

1.330 1.431 1.422 1.501 1.516 1.392 1.406

Ultrasonic Sound Velocity Measurements. A single crystal ultrasonic interferometer (model F-05) from Mittal Enterprises, New Delhi, India, at 2 MHz frequency was used for sound velocity measurements for pure solvents and mixtures at 25 °C. Prior to measurements, the interferometer was calibrated with double distilled water and purified methanol. The uncertainty of the sound velocity was 0.2%. Viscosity Measurements. Vibro viscometer (model SV-10, A&D Company Limited, Japan) was used for viscosity measurements. The instrument has been provided with two sensor plates of gold coating. The measurements of η were taken from the digital display device attached to the vibro viscometer. Viscosity measurements of the sample were taken at heating rate of 1 °C/15 min for getting the thermodynamic equilibrium. Viscometer was calibrated with pure water and purified methanol at various temperatures. The uncertainty of the η measurements was less than 1%. Refractive Index Measurements. The refractive index was determined using Abbe refractometer from Mittal Enterprises, New Delhi, India, with an accuracy of ±0.0002. Refractometer was calibrated by measuring the nD values of the high purity water and purified methanol at various temperatures. A thermostatically controlled, well-stirred circulating water bath with a temperature controlled to ±0.01 °C was used for u, η, and nD measurements. Hydrogen Bonding through Molecular Modeling Program. The structures of ammonium-based ILs and water were optimized based on molecular mechanics and semiempirical calculations using the HyperChem7 molecular visualization and simulation program. The detailed procedure is depicted elsewhere.27 Initial molecular geometry of water and ILs were optimized with the AM1 semiempirical calculations and single point calculations were carried out to determine the total energies.36 Now the optimized molecules, water and IL were chosen and then placed on top of each other symmetrically (parallel) with a starting inter planar distance of 2.3 Å and the angle made by covalent bonds to the donor and acceptor atoms less than 120° was fulfilled.36−38 The geometries were optimized using geometry optimizations based on molecular mechanics (using the MM + force field) and AM1 semiempirical calculations, the Polak-Ribiere routine with rms gradient of 0.01 as the termination condition was used.39 Hydrogen bonds were obtained by using HyperChem7 “show hydrogen bonds” and “recompute hydrogen bond” options. We have also calculated the heat of formation (ΔHf) and binding energies (E) of the pure components and their mixtures. Using this calculation, we obtained that the ammonium ILs are engaged in the H-bonding with the water.

In this work, we report the thermo-physical properties such as ρ, u, η, and nD as well as derived properties of excess molar volume (VE), deviation in isentropic compressibilities (Δκs), deviation in viscosities (Δη) and deviation in refractive indices (ΔnD) for aqueous solution of ammonium-based ILs. Further, the study intends to draw molecular level information on the molecular interactions between ILs and water from the macroscopic properties. The ILs investigated in the present study included diethylammonium acetate (DEAA), triethylammoniumacetate (TEAA), diethylammonium hydrogen sulfate (DEAS), triethylammonium hydrogen sulfate (TEAS), trimethylammonium acetate (TMAA) and trimethylammonium hydrogen sulfate (TMAS). The experimental values of VE, Δκs, Δη and ΔnD values were correlated by Redlich−Kister polynomial equation. To the best of our knowledge no efforts have been made in the literature to study the molecular interactions between these ammonium-based ILs and water. Moreover, the hydrogen bonding features between these ILs and water were carried out to obtain deeper insight into intermolecular interactions by the semiempirical calculations using HyperChem7.

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EXPERIMENTAL PROCEDURE

Materials. Six ILs are used in the present study, namely DEAA, TEAA, DEAS, TEAS, TMAA, and TMAS were synthesized in our laboratory. The complete procedure of synthesis of these ILs had been distinctly demonstrated in our previous articles.9,20 The water plays a significant role in the properties of ILs and obviously significant variations can be found in their properties.35 All investigated ILs have low levels of water (below 70 ppm) analyzed by Karl-Fischer titration. Water was obtained from a NANO pure-ultra water system (Rions, New Delhi) which is distilled, deionized, degassed having a resistivity of 18.3 Ω·cm. The purity and the investigated thermo-physical properties of water and ILs are reported in Table 1. Methods and Procedure. Clear binary mixtures were prepared by weighing the components using a Mettler Toledo balance with an accuracy of ±0.0001 g for all measurements. The estimated uncertainty on the mole fraction composition was found to be less than 5 × 10−4. After mixing the sample, the bubble-free homogeneous sample was transferred into the U-tube of the densimeter or the sample cell of ultrasonic interferometer or refractometer through a medical syringe. A bubble-free sample was introduced into the sample cell, and the cell was placed under the sensor plates of the viscometer. Density Measurements. The density measurements of various ILs, water, and ILs + water were carried out using an Anton-Paar DMA 4500 M vibrating-tube densimeter, equipped with a built-in solid-state thermostat and a resident program with an accuracy of temperature of ±0.03 °C. The precision of the density measurements was ±0.00005 g cm−3. The instrument was calibrated once a day with double distilled, deionized water and with air at 20 °C as standards. The excess molar volumes (VE) of IL with water systems over the IL concentration range at 25 °C have been deduced from the ρ of the pure compounds and their mixture (ρm) using the standard equations.

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RESULTS AND DISCUSSION For understanding the effect of water on the thermophysical properties of the ammonium-based ILs, the values of ρ, u, η, and nD for the ILs with water systems were measured at 25 °C under atmospheric pressure. From Table 1, it is clear that the studied ILs have higher ρ values than water. The experimental ρ, u, η, and nD values for the binary mixtures involving water with ammonium family ILs are presented at 25 °C as a function of IL concentration in Table 2 and Figure 1. The results in 5972



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Table 2. Mole fraction (x1) of ILs, Density (ρ), Ultrasonic Sound Velocity (u), Viscosity (η), Refractive Index (nD), Excess Molar Volumes (VE), Isentropic Compressibility (κs), Deviation in Isentropic Compressibility (Δκs), Deviation in Viscosity (Δη), and Deviation in Refractive Index (ΔnD) for the Systems of ILs with Water at 25 °C and at Atmospheric Pressure x1

ρ/g cm−3

u/m·s−1

η/mPa·s

nD

0.0000 0.0366 0.0532 0.0803 0.1059 0.1265 0.1792 0.2567 0.2821 0.3438 0.3818 0.4259 0.4783 0.5410 0.5773 0.6627 0.7134 0.7707 0.8363 0.9119 1.0000

0.997 04 1.003 35 1.005 87 1.008 84 1.010 98 1.012 50 1.015 14 1.017 62 1.018 22 1.019 23 1.019 67 1.020 11 1.020 46 1.020 76 1.020 87 1.021 04 1.021 12 1.021 21 1.021 24 1.021 30 1.021 41

1496 1510 1516 1524 1532 1538 1551 1568 1572 1581 1585 1589 1591 1594 1595 1596 1597 1598 1599 1602 1608

0.89 2.77 3.48 4.68 5.59 6.37 8.01 9.91 10.45 11.58 12.02 12.56 13.01 13.40 13.61 14.06 14.26 14.52 14.82 15.40 16.36

1.330 1.346 1.355 1.365 1.374 1.380 1.395 1.410 1.415 1.422 1.425 1.427 1.428 1.429 1.430 1.427 1.426 1.424 1.425 1.426 1.431

0.0000 0.0259 0.0378 0.0576 0.0767 0.0921 0.1327 0.1949 0.2159 0.2686 0.3021 0.3422 0.3912 0.4524 0.4891 0.5793 0.6356 0.7021 0.7817 0.8788 1.0000

0.997 04 1.048 16 1.066 92 1.093 84 1.115 17 1.130 05 1.161 34 1.194 54 1.203 07 1.220 28 1.228 78 1.237 14 1.245 46 1.253 61 1.257 62 1.265 29 1.268 86 1.272 32 1.275 94 1.279 70 1.283 95

1496 1468 1461 1454 1448 1444 1439 1438 1437 1439 1440 1441 1442 1443 1439 1437 1436 1433 1429 1426 1432

0.89 6.19 8.48 11.06 13.51 15.14 19.31 24.56 26.04 28.77 30.24 31.38 32.09 32.33 32.34 31.26 30.99 30.09 29.61 30.40 34.10

1.330 1.340 1.343 1.349 1.354 1.358 1.367 1.379 1.383 1.390 1.394 1.398 1.402 1.405 1.406 1.409 1.410 1.411 1.412 1.415 1.422

0.0000 0.0288 0.0420 0.0638 0.0847 0.1016 0.1457 0.2125 0.2349 0.2904 0.3254 0.3670 0.4174 0.5162

0.997 04 1.001 22 1.002 72 1.004 63 1.006 11 1.007 15 1.009 15 1.011 13 1.011 60 1.012 58 1.013 04 1.013 50 1.013 95 1.014 59

1496 1532 1549 1571 1594 1610 1651 1701 1720 1753 1769 1784 1800 1810

0.89 6.32 8.12 10.70 13.35 14.91 18.62 23.38 24.44 26.47 27.44 27.82 28.16 27.57

1.330 1.355 1.363 1.377 1.390 1.398 1.418 1.443 1.449 1.462 1.468 1.473 1.478 1.483

VE/cm−3 mol−1 DEAA + Water 0.000 −0.023 −0.043 −0.064 −0.080 −0.096 −0.120 −0.147 −0.155 −0.162 −0.163 −0.164 −0.158 −0.148 −0.138 −0.112 −0.096 −0.078 −0.052 −0.025 0.000 DEAS + Water 0.000 −0.080 −0.113 −0.171 −0.221 −0.259 −0.350 −0.462 −0.492 −0.551 −0.571 −0.581 −0.583 −0.566 −0.548 −0.473 −0.403 −0.310 −0.208 −0.095 0.000 TEAA + Water 0.000 −0.006 −0.012 −0.015 −0.018 −0.023 −0.031 −0.040 −0.042 −0.047 −0.049 −0.050 −0.051 −0.048 5973

κs/TPa−1

Δκs/TPa−1

Δη/mPa·s

ΔnD

448 437 432 426 421 417 409 399 397 392 390 388 387 385 385 384 384 383 382 381 379

0.0 −8.5 −11.9 −15.8 −19.4 −21.8 −26.2 −30.6 −31.1 −31.7 −31.3 −30.3 −27.7 −24.9 −22.9 −17.6 −14.6 −11.1 −7.0 −3.2 0.0

0.000 1.319 1.776 2.547 3.063 3.521 4.344 5.048 5.194 5.369 5.228 5.080 4.715 4.143 3.788 2.917 2.334 1.711 0.992 0.406 0.000

0.000 0.012 0.019 0.027 0.033 0.037 0.047 0.054 0.056 0.057 0.056 0.054 0.049 0.044 0.041 0.030 0.024 0.016 0.010 0.004 0.000

448 443 439 432 427 424 416 405 402 396 392 389 386 383 384 383 382 383 384 384 379

0.0 −3.6 −6.4 −11.7 −15.2 −17.4 −23.2 −29.9 −30.8 −34.1 −35.1 −35.5 −35.3 −34.1 −30.7 −25.8 −22.5 −17.4 −10.9 −3.8 0.0

0.000 4.443 6.337 8.253 10.078 11.194 14.016 17.197 17.982 18.958 19.318 19.128 18.208 16.419 15.204 11.131 8.996 5.885 2.753 0.328 0.000

0.000 0.007 0.009 0.014 0.017 0.019 0.025 0.031 0.033 0.035 0.036 0.036 0.036 0.033 0.031 0.026 0.021 0.016 0.010 0.000 0.000

448 425 415 403 391 383 363 342 334 321 315 310 304 301

0.0 −18.1 −25.8 −34.8 −43.6 −49.1 −61.7 −72.9 −77.0 −81.1 −81.5 −80.4 −78.0 −66.1

0.000 4.761 6.253 8.324 10.488 11.662 14.341 17.556 18.096 18.835 18.994 18.403 17.577 14.687

0.000 0.020 0.026 0.036 0.045 0.051 0.063 0.077 0.079 0.082 0.082 0.080 0.076 0.065



Article

Table 2. continued x1

ρ/g cm−3

u/m·s−1

η/mPa·s

nD

0.5579 0.6055 0.6604 0.7243 0.7996 0.8899 1.0000

1.014 78 1.014 97 1.015 18 1.015 34 1.015 51 1.015 69 1.015 86

1813 1812 1805 1802 1800 1809 1840

26.81 26.07 25.33 24.21 22.93 22.67 24.12

1.481 1.483 1.481 1.480 1.481 1.487 1.501

0.0000 0.0257 0.0373 0.0568 0.0756 0.0909 0.1310 0.1925 0.2134 0.2656 0.2989 0.3387 0.3876 0.4853 0.5272 0.5756 0.6321 0.6988 0.7791 0.8772 1.0000

0.997 04 1.027 48 1.038 52 1.052 98 1.063 88 1.071 75 1.086 94 1.102 27 1.106 19 1.113 88 1.117 72 1.121 53 1.125 27 1.130 80 1.132 58 1.134 36 1.136 11 1.137 82 1.139 52 1.141 18 1.142 89

1496 1507 1526 1542 1562 1581 1616 1677 1702 1748 1772 1795 1818 1843 1844 1837 1836 1829 1819 1828 1874

0.89 14.14 19.88 29.46 39.84 48.75 63.13 87.35 95.21 110.78 120.03 129.26 139.58 155.92 162.36 169.20 175.83 183.80 194.74 210.28 235.10

1.330 1.362 1.373 1.389 1.403 1.417 1.442 1.473 1.483 1.500 1.508 1.515 1.519 1.521 1.518 1.515 1.509 1.504 1.500 1.501 1.516

0.0000 0.0451 0.0651 0.0978 0.1283 0.1524 0.2133 0.3002 0.3280 0.3943 0.4341 0.4797 0.5325 0.5942 0.6292 0.7094 0.7556 0.8068 0.8639 0.9278 1.0000

0.997 04 1.011 94 1.016 96 1.023 69 1.028 34 1.031 45 1.037 43 1.042 95 1.044 25 1.046 66 1.047 78 1.048 86 1.049 80 1.050 71 1.051 15 1.051 94 1.052 34 1.052 71 1.053 04 1.053 43 1.053 85

1496 1499 1503 1504 1507 1510 1517 1525 1527 1532 1534 1535 1537 1536 1538 1539 1540 1539 1541 1542 1544

0.890 1.199 1.31 1.50 1.64 1.75 1.99 2.24 2.31 2.43 2.48 2.52 2.56 2.59 2.61 2.62 2.63 2.64 2.66 2.69 2.76

1.330 1.339 1.342 1.347 1.351 1.354 1.361 1.369 1.371 1.376 1.378 1.380 1.382 1.384 1.385 1.387 1.388 1.389 1.390 1.391 1.392

0.0000 0.0448 0.0648 0.0973 0.1276 0.1516 0.2122 0.2988

0.997 04 1.101 41 1.137 13 1.184 38 1.220 13 1.244 42 1.292 73 1.340 91

1496 1464 1466 1470 1472 1482 1502 1531

0.89 1.30 1.49 1.76 2.01 2.18 2.58 3.08

1.330 1.342 1.346 1.353 1.359 1.363 1.372 1.381

VE/cm−3 mol−1 TEAA + Water −0.046 −0.042 −0.039 −0.032 −0.023 −0.012 0.000 TEAS + Water 0.000 −0.021 −0.042 −0.059 −0.071 −0.094 −0.124 −0.157 −0.169 −0.187 −0.195 −0.202 −0.204 −0.196 −0.186 −0.174 −0.157 −0.131 −0.098 −0.051 0.000 TMAA + Water 0.000 −0.043 −0.064 −0.099 −0.112 −0.137 −0.172 −0.202 −0.208 −0.213 −0.211 −0.206 −0.192 −0.174 −0.162 −0.129 −0.110 −0.086 −0.056 −0.028 0.000 TMAS + Water 0.000 −0.041 −0.067 −0.091 −0.113 −0.138 −0.184 −0.228 5974

κs/TPa−1

Δκs/TPa−1

Δη/mPa·s

ΔnD

300 300 302 303 304 301 291

−60.5 −52.7 −41.8 −30.8 −18.4 −7.2 0.0

12.963 11.112 9.099 6.499 3.467 1.113 0.000

0.056 0.049 0.038 0.026 0.014 0.005 0.000

448 428 413 399 385 373 352 323 312 294 285 277 269 260 260 261 261 263 265 262 249

0.0 −14.5 −27.2 −37.4 −47.8 −56.7 −69.7 −87.2 −93.6 −101.4 −103.7 −104.0 −102.1 −91.2 −83.6 −72.4 −61.2 −46.3 −27.9 −11.3 0.0

0.000 7.256 10.256 15.256 21.236 26.586 31.569 41.383 44.353 47.686 49.146 49.027 47.916 41.365 37.997 33.504 26.901 19.235 11.391 3.949 0.000

0.000 0.027 0.036 0.048 0.059 0.070 0.087 0.107 0.113 0.121 0.122 0.122 0.117 0.101 0.090 0.078 0.061 0.044 0.025 0.008 0.000

448 440 435 432 428 425 419 412 412 407 406 405 403 403 402 401 401 401 400 399 398

0.0 −6.1 −9.6 −11.4 −13.5 −15.3 −18.6 −20.8 −21.0 −21.3 −20.8 −19.5 −18.2 −14.9 −14.4 −11.2 −9.6 −6.6 −4.9 −2.4 0.0

0.000 0.225 0.295 0.424 0.509 0.576 0.700 0.786 0.802 0.799 0.780 0.729 0.674 0.588 0.545 0.404 0.328 0.242 0.153 0.073 0.000

0.000 0.006 0.008 0.011 0.013 0.014 0.018 0.020 0.021 0.022 0.021 0.020 0.019 0.017 0.016 0.013 0.011 0.009 0.006 0.003 0.000

448 423 409 391 378 366 343 318

0.0 −16.9 −27.9 −40.9 −48.2 −56.5 −69.2 −79.3

0.000 0.221 0.327 0.465 0.583 0.651 0.79 0.939

0.000 0.008 0.011 0.016 0.019 0.022 0.026 0.028



Article

Table 2. continued x1 0.3266 0.3927 0.4325 0.4781 0.5309 0.6277 0.6661 0.7081 0.7544 0.8058 0.8631 0.9273 1.0000

ρ/g cm−3 1.352 88 1.376 79 1.388 69 1.400 51 1.412 17 1.429 41 1.435 09 1.440 58 1.446 11 1.451 58 1.456 96 1.462 33 1.467 58

u/m·s−1 1541 1556 1562 1564 1565 1562 1560 1558 1556 1555 1553 1558 1564

η/mPa·s 3.22 3.54 3.69 3.88 4.05 4.34 4.43 4.53 4.64 4.74 4.86 4.98 5.10

VE/cm−3 mol−1

nD 1.384 1.389 1.391 1.393 1.395 1.397 1.398 1.399 1.399 1.400 1.401 1.403 1.406

TMAS + Water −0.237 −0.255 −0.263 −0.266 −0.263 −0.241 −0.229 −0.204 −0.179 −0.148 −0.108 −0.061 0.000

κs/TPa−1

Δκs/TPa−1

Δη/mPa·s

ΔnD

311 300 295 292 289 287 286 286 286 285 284 282 278

−81.5 −81.5 −79.6 −75.2 −68.9 −54.9 −48.8 −42.1 −34.6 −26.6 −17.2 −9.1 0.0

0.959 0.995 0.98 0.977 0.926 0.809 0.738 0.661 0.577 0.455 0.340 0.186 0.000

0.029 0.029 0.028 0.027 0.025 0.019 0.017 0.015 0.012 0.008 0.005 0.002 0.000

Figure 1. Experimental (a) densities, (b) ultrasonic sound velocities, (c) viscosities, and (d) refractive indices for the mixtures of ILs with water as a function of mole fraction of ILs (x1) for (O) DEAA, (Δ) DEAS, (□) TEAA, (●) TEAS, (▲) TMAA, and (■) TMAS with water at 25 °C under atmospheric pressure. The solid lines represent the smoothness of these data.

length of the cation. The data in Figure 1a show that the ρ values of the ILs in water follow the order TMAS > DEAS > TEAS > TMAA > DEAA > TEAA, which indicates that the lower alkyl chain length of cation of ILs is much denser than higher alkyl chain length of ILs (for example 1.02141 g cm−3 for DEAA and 1.01586 g cm−3 for TEAA; 1.28395 g cm−3 for

Figure 1a reveal that the ρ values for the mixtures of all studied ILs with water increase as the concentrations of the IL in water increase. The increase in ρ values for IL with water mixtures is possibly due to increase in the ion-pair interactions between IL and water. The thermo-physical properties of ILs mainly depend on the nature and structure of ions and the alkyl chain 5975



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DEAS and 1.14289 g cm−3 for TEAS). This is mainly due to increase in dispersive interactions in ILs with increase in chain length, resulting in a nanostructural organization in polar and nonpolar regions. The nonpolar regions are buildup of alkyl chains whereas the polar groups contain the cationic head groups and the anions. When we enhance the chain length of cation, the nonpolar regions increase and take up more and more space, as result lower in overall density in higher alkyl chain length of ILs.40−42 As can be seen in Figure 1a the acetate ILs display significantly lower ρ values than the corresponding sulfate-based ILs due to the ionic species and the increased molecular mass of the anion. Usually, ρ values are related to molar mass of the ions of ILs and it is obvious that ILs containing larger molar mass of sulfate-based IL systems found to be most denser than acetate-based ILs systems. Clearly, our results show that the ρ is quite sensitive to the size of the cation and anion of ammonium ILs and also the composition of mixture. With regard to ultrasonic sound velocity, both size of the ions and water content have significant influence on the u values of the IL, and the results have been shown in Figure 1b at 25 °C over the whole composition range. The variation of the values of u for ammonium-based IL with water varies with the composition in an anomalous way. From analysis of the data in Figure 1b, it was found that the u values depend more strongly on the mole fraction of IL in the solution and, the trend for all the systems with respect to composition was not the same. The values of u are found to increase rapidly with increasing the mole fraction of TEAA or TEAS up to x1 ≈ 0.5000 after which the increment was too small (Figure 1b) for TEAA or TEAS with water. On the other hand, the values of u for the DEAA or TMAA or TMAS with water system increased slightly with increasing IL composition. On the contrary, the u values decreases slowly when the mole fraction of DEAS increases up to x1 ≈ 0.1100, later the values obviously became nearly constant. The data in Figure 1b show that the u values of ILs + water follow the order TEAS > TEAA >DEAA > TMAA> TMAS> DEAS. The values of η for the ammonium-based ILs are exhibited in Figure 1c as a function of the mole fraction of ILs. From Figure 1c, it is quite clear that IL water mixtures are less viscous than pure ILs however more viscous than water. An increase in η is obtained when DEAA or TMAA or TMAS composition enhances in mixture of DEAA or TMAA or TMAS with water. Interestingly, the η values for the mixture of TEAS + water increase rapidly with increasing the mole fraction of TEAS. In other words, this result has been interpreted by the fact that the strong columbic interaction between the ions is strengthening upon mixing with the water leading to a low mobility of the ions of ILs. However, the η values for the system of DEAS or TEAA with water increase close to ≈0.4500 or ≈0.4100 mol fraction of DEAS or TEAA, respectively, later abruptly a slight increase is observed for the rest of the IL composition, as shown in Figure 1c. The data in Table 2 and Figure 1c show that the values of η for ILs + water follows the order: TEAS > DEAS > TEAA >DEAA > TMAS >TMAA. Clearly, the acetate-based IL mixtures have lower viscosity than the corresponding sulfatebased IL systems under the same experimental conditions. The anionic viscosity effect and their ionic nature are most obvious in the present study, resulting in significantly higher η for sulfate IL mixture as compared to the corresponding acetate IL mixtures. For all the investigated systems with a common

anion, the η values increase with the increase of the alkyl side chain length of the cation. This trend is due to increase on the van der Waals interactions between the alkyl side chains of the cations and on the proportion of the charged species in the overall mixtures. Among the ILs studied here TEAS system shows the highest η values due to the size of the anion and the number of carbon atoms on the cation (TEA+) groups. As it can be seen in Figure 1, the ρ values decrease with increasing cation alkyl chain length from diethyl to triethyl while an opposite trend was observed, in which the values of η increase with increasing the number of carbon atoms in the alkyl chain length of cation of ILs. This is mainly contributing to anion accommodation closer to the cation. These observed results are quite consistent with the existing results,35,43−46 in which ρ values decrease while η values increase with increasing the cation alkyl chain length of ILs. It appears that, IL systems that possess a higher cation side chain is accompanied by lower ρ and larger η. Clearly, our results might imply that the cation size is responsible for the alteration of the thermophysical properties of protic ILs in mixture with water. Figure 1d depicts the measured values of the nD for the six ammonium-based IL with water at 25 °C over the whole composition range. The nD values for ILs with water mixtures increase with increasing composition of IL, except DEAA or TEAA or TEAS with water systems. Figure 1d reveals that the nD values for DEAA or TEAA or TEAS + water system increase up to x1 ≈0.5700 or ≈0.5500 or ≈0.4800 on addition of DEAA or TEAA or TEAS to water, respectively. Later, the nD values slightly decrease with increasing the mole fraction of IL for these systems. This may be due to diminished ion−ion pair interactions between the IL and water and also self-interaction between the ions of IL. For the ammonium family ILs with water, the nD values followed sequence are TEAS > TEAA > DEAA > DEAS > TMAS > TMAA. This order clearly shows that sulfate anion with same cation ILs show larger nD values as compared to the corresponding acetate IL mixtures, except DEA+ cation ILs. The highest nD values are due to on the ions arrangement and an efficient packing of ions of ILs. In addition, the nD increases with the anion molecular weight that leads the nD of a system is higher, when that system is denser.47−49 The results in parts a and d of Figure1 show that this phenomena is quite consistent for TEA+ or TMA+ cation system, however, it is not obvious for DEA+ systems. The excess and deviation properties provide a good estimate of the strength of unlike molecular interactions in the solution phase. Therefore, the experimental thermo-physical properties, the ρ, u, η, and nD of the ammonium-based ILs and the ILs with water systems are further used to obtain VE, Δκs, Δη and ΔnD according to well-known thermodynamic expressions. These properties were fitted with the following Redlich−Kister polynomial equation. n

Y = x1x 2(∑ ai(x1 − x 2)i ) i=0

⎡ ∑n (Y − Y )2 ⎤1/2 cal i = 1 exp ⎥ σ=⎢ ⎢⎣ ⎥⎦ n

(1)

(2)

where n is the number of experimental data, where Y refers to VE or Δκs or Δη or ΔnD, and where x1 and x2 are mole fractions of pure compounds. Yexp is the experimental property and Ycal is the calculated from eq 1. ai are adjustable parameters and can 5976



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Table 3. Estimated Parameters of eq 1 and Standard Deviation (σ) for the Systems of ILs with water at 25 °C Y

water +

ao

a1

a2

σ

VE/cm3 mol−1

DEAA TEAA DEAS TEAS TMAA TMAS DEAA TEAA DEAS TEAS TMAA TMAS DEAA TEAA DEAS TEAS TMAA TMAS DEAA TEAA DEAS TEAS TMAA TMAS

−0.634 −0.198 −2.176 −0.793 −0.816 −1.065 −105.806 −262.785 −131.326 −353.721 74.329 −297.013 17.999 57.455 54.306 166.869 2.850 3.826 0.194 0.252 0.119 0.366 0.078 0.103

0.316 0.079 1.451 0.382 0.388 0.089 102.617 292.525 104.721 347.928 58.520 194.175 −17.214 −74.020 −83.479 −168.043 −2.229 −1.529 −0.186 −0.323 −0.123 −0.481 −0.044 −0.085

0.113 0.023 0.403 0.085 0.109 0.149 −35.779 −91.771 37.920 −17.766 −18.958 20.359 2.463 28.647 31.928 −22.867 0.207 0.247 0.010 0.106 0.035 0.171 0.019 0.013

0.003 0.001 0.005 0.004 0.003 0.003 0.612 1.337 1.318 1.510 0.558 1.806 0.057 0.672 0.837 1.157 0.005 0.007 0.001 0.002 0.001 0.004 0.001 0.001

Δκs/TPa−1

Δη/mPa·s

ΔnD

be obtained by least-squares analysis. Values of the fitted parameters are collected in Table 3 along with the standard deviations of the fit. The VE, Δκs, Δη, and ΔnD values of six ammonium-based ILs with water are also included in Table 2 and also illustrated in Figure 2 along with as fitting curves from eq 1 as a function of the mole fraction of IL at 25 °C. Figure 2a, which is the VE graph of ILs with water, depicts that the values are negative over a wide mole fraction range at 25 °C under atmospheric pressure and these curves are asymmetric. Analogously, the negative values have been observed for another type ammonium-based ILs with water,33,34,50 ammonium family ILs with DMF20−22 or DMSO,23−27 or NMP28−30 over the whole composition range. The observed negative values reveal more attractive interactions in the mixtures than in the pure components and the systems have a strong packing effect by hetero associations between both molecules through hydrogen bonding. Looking at the data in Figure 2a, it is clear that the magnitude of negative VE values has a broad minimum at the mole fraction of IL ≈ 0.3400 or ≈0.3500 or ≈0.4000 or ≈0.3300 or ≈0.3900 or ≈0.4100 for DEAA or DEAS or TEAA or TEAS or TMAA or TMAS with water, respectively. The negative VE values can be due to hydrogen bonding between ions of ILs with water molecules. Further, we have also observed that the magnitude of negative sign decreases as mole fraction of IL increases in the studied systems. In this regard, we can assert that more efficient packing is due to the differences in size and shape of molecules of the mixtures or attractive interaction occurs in the region of low mole fraction of IL. The negative contribution arises from changes of free volume in the real mixtures, comprising ions and water molecules, such behavior might arise from restriction of rotational motion, when the water molecules are accommodated interstitially with in ions of ILs. Obviously, the VE depends on the interactions and on the size and shape of

molecules of the system. From Figure 2a, it can be noted that the magnitude of VE values for ILs with water at 25 °C showed the following trend: DEAS > TMAS > TMAA > TEAS > DEAA > TEAA

From this order, it appears that the increase of the alkyl chain length of the IL from DEAA to TEAA affected strongly the VE values of the solutions. As noted from Figure 2a that the negative VE values for DEAA + water (VE = −0.164 cm3 mol−1 at x1 = 0.4259) are obviously more negative than those for TEAA + water (VE = −0.051 cm3 mol−1 at x1 = 0.4174). Subsequently, the negative VE values for DEAS + water (VE = −0.583 cm3 mol−1 at x1 = 0.3912) are obviously higher than those for TEAS + water (VE = −0.204 cm3 mol−1 at x1 = 0.3876). As a result, VE values become less negative in higher alkyl length of the IL under the same experimental condition. The less negative VE values for TEA+ mixture serves as an evidence that higher alkyl chain molecules decrease the hydrogen bonding tendency between TEA+ with water. On the other hand, the DEA+ with water mixture reveals more negative values of VE than TEA+ with water system mixture, which imply that DEAA ion−dipole interactions and packing effects with water are stronger than those in the TEAA system. Clearly, we might expect moderate steric hindrance of the alkyl chain in TEA+ molecules. Therefore, it is important to note that the nature of interactions in IL with water systems is highly dependent on nature of the ions. It is quite clear from Figure 2 that anion structure in ammonium-based ILs strongly affect the VE values. It was found that sulfate-based IL systems exhibit more negative VE values than the corresponding acetate-based IL systems. It has been shown that H-bonding of water molecules with ions of IL mainly depend on the basicity of the anion of ILs47 and obvious that water molecules enhance hydrogen bonding with anions of stronger basicity. Moreover, the high VE negative values in 5977



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Figure 2. Plot of (a) excess molar volumes (VE), (b) deviation in isentropic compressibility (ΔκS), (c) deviation in viscosities (Δη) and (d) deviation in refractive indices (ΔnD) against the mole fraction of ILs for (O) DEAA + water; (Δ) DEAS + water; (□) TEAA + water; (●) TEAS + water; (▲) TMAA + water; and (■) TMAS + water at 25 °C under atmospheric pressure. Solid lines are correlated by Redlich−Kister equation.

of hydrogen bonding strength between CH3COO− and water. As a consequence, the acetate-based ILs weakly interact with water molecules whereas sulfate-based ILs interact strongly with water molecules. As can be seen in Figure 2b, the Δκs values for the IL + water system are negative for all of compositions studied at 25 °C. The fitting curves are asymmetric and present a minimum which is obtained in the water-region at x1 ≈ 0.3200 to 0.3900. The negative contributions are attributed to the strong attractive interactions due to the solvation of the ions in the solution. A strong intermolecular interaction through charge transfer, dipole−induced dipole and dipole−dipole interactions, interstitial accommodation and orientational ordering lead to a more compact structure which contributes to negative deviation in Δκs values. The negative values of Δκs of the IL with water imply that solvent molecules around solute are less compressible than the solvent molecules in the bulk solutions. While on further addition of IL there is a decrease in the compressibility of water. This might be due to decrease in attraction of water molecules and ions of IL in IL-rich composition region. The algebraic values of Δκs for ammonium family ILs with water fall in the order:

sulfate-based IL systems may be due to the basicity of HSO4 anion is stronger than that of CH3COO anion that leads to increase the strength of H-bonding between water molecules and sulfate anion. Therefore, taking into account that VE values involve a net creation of interactions upon mixing, it can be stated that for sulfate-based ILs the forces between the ions or molecules that form the dissimilar species, (IL−water interactions), are stronger than between those of similar ones (IL−IL and water−water interactions). As a result, it can be said that packing effects and ion−ion pair interactions in the sulfate-based IL systems are dominated than the acetate-based IL systems. The analysis of the studied ammonium-based ILs presents a considerably different behavior in aqueous solution, which is mostly due to the chemical structure of the anion. Since, water molecules interact mainly through hydrogen bonding, this results can be explained by the different capability of each anion to form hydrogen bonds with water molecules: the HSO4 functional group presents four accessible, negatively charged oxygen atoms to form a hydrogen bond, on the other hand the CH3COO anion present only two oxygen atoms as interacting sites with water molecules. In addition, the −CH3 group in the acetate anion decreases the electron affinity of oxygen due to + I effect of methyl group that results in decrease 5978



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Figure 3. Hydrogen bonding interaction between (a) DEAA + water, (b) DEAS + water, (c) TEAA + water, (d) TEAS + water, (e) TMAA + water, and (f) TMAS + water molecules, which is predicted by semiempirical calculation with the help of HyperChem7.

TEAS > TEAA ≥ TMAS > DEAS > DEAA > TMAA

than those between sulfate-based ILs interact and water molecules. As shown in Figure 2c, the Δη values for ILs with water are positive over the whole range of compositions at 25 °C under atmospheric pressure the maximum existed at IL region; i.e., x1 ≈ 0.3000−0.3500. The positive deviation for these systems indicates strong interaction between ammonium-based IL and

It is clear that different phenomena of Δκs were observed for the various ILs with water. As can be seen from this order the Δκs values for sulfate-based ILs with water systems are higher than those in acetate-based ILs with water. The lower Δκs values for acetate-based ILs serves as further evidence that the interactions between acetate-based ILs with water are weaker 5979



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water molecules. The positive deviation is characteristic of mixtures containing hydrogen bonding between ions of IL and water molecules. The absolute values of Δη for these systems fall in the order:

Table 4. Calculated Heats of Formations (ΔHf) and Estimated Hydrogen Bond Energies (ΔΔHf) (kcal/mol)

TEAS > TEAA > DEAS > DEAA > TMAS > TMAA

The highest Δη was observed for TEAS with water, this may be due to the formation of strong hydrogen bonding interactions of the ion with water molecules as compared with the rest of the ions of IL with water. Further, this order explicitly elucidates that the Δη values increase with the cation alkyl chain length and a crucial effect of the anion was observed. The Δη values of sulfate-based IL systems are more than those of acetate-based IL mixtures due to the remarkable difference in the η of pure ILs and also the ability of the formation of hydrogen bonding with water molecules. A close examination of Figure 2 it should be noted that the main factor for a negative VE and a positive Δη may be association or complex formation between the components of ammonium ILs with water. This phenomenon is similar to the binary systems of another kind of ammonium-based ILs with DMSO27 or NMP30 at 25 °C. It is obvious that the difference in these properties between these ILs is caused mainly due to the conformational changes in the structural interactions and ion−ion pair interactions, size and also shape of the components. Moreover, the structure and nature of cations and anions greatly influence the physicochemical properties of ILs. The obtained ΔnD for the binary mixture of ILs with water are positive in the entire composition range at 25 °C as shown in Figure 2d and the maximum is reached near to 0.3000− 0.4000 mol fraction of IL. These values depend mainly on variation in intermolecular interactions between two components into contact. On the other hand, the values followed sequence is

solvent

ΔHf/(kcal mol−1)

ΔΔHf/(kcal mol−1)

water DEAA DEAS TEAA TEAS TMAA TMAS DEAA + water DEAS + water TEAA + water TEAS + water TMAA + water TMAS + water

−53.273 −36.362 −33.897 −14.817 −26.254 −20.702 −23.548 −165.857 −174.108 −154.308 −255.239 −148.942 −242.372

76.22 86.93 89.21 175.71 72.12 165.55

Table 4 indicate that the energy required for the formation of a weak hydrogen bond is less than required for the formation of a stronger hydrogen bond: ΔΔHf = ΔHf (1) + ΔHf (2) − ΔHf (3)

(3)

where ΔHf (1) is the heat of formation of the water, ΔHf (2) the heat of formation of the ILs, and ΔHf (3) the heat of formation of the complex (water and ILs). Our HyperChem results clearly show that the hydrogen bonding between cation of IL with the water molecules. The estimated hydrogen bond energies of the TEAS + water are found to be 175.71 kcal/mol and TMAS + water found to be 165.55 kcal/mol (Table 4). Hence, the estimated hydrogen bond energy of the above systems (TEAS and TMAS) is highest among all the ILs with water. This might be due to strong interaction of ammonium cation, sulfate anion and water that result in the highest estimated hydrogen bond energies. Whereas the estimated hydrogen bond energy of the DEAS + water is 86.93 kcal/mol, this shows that presence of sulfate group itself does not have much significance in increasing the hydrogen bond energies, the role of cation is also playing an important role. In another case, if we check from Table 4 carefully we observed that TEAA + water has estimated hydrogen bond energies of 89.21 kcal/mol, which is more than TEAS + water. Hence, there is no doubt that sulfate anion molecules have strong hydrogen bond energies value, while the presence of cation also plays important role in the interaction. This phenomena is well consistent with our experimental results, in which sulfate-based ILs strongly interact with water molecules whereas the acetate-based ILs interact weakly with water molecules. In other words, the presence of water in the ILs acts as impurity, because water is trap between the interaction of cation of IL and anion IL, that results in the decrease of interaction between the cation and anion of ILs, as illustrated in Scheme 1. Hence, in different ILs we observed the different values of VE, because these values are not only due to interaction of water with cation of ILs (as shown in HyperChem results), while it is also a part of ion−ion and ion-dipole interactions between the cation and anion of ILs. We observed more negative VE values in sulfate ILs than acetate ILs. Moreover, sulfate anion and ammonium cation have very high interaction due to high electron affinity as compare to the acetate anion and ammonium cation. Although, the interaction of water with cation of ammonium ILs is same, but the

TEAS > TEAA > DEAA > DEAS > TMAS > TMAA

The ΔnD values increase with increase chain length of the cation and a higher ΔnD values for mixtures containing TEA+ than for mixtures with DEA+. As can be seen in Figure 2, parts a and d, there is strong correlation between VE and ΔnD quantities for all the studied systems. We observed negative VE values corresponded to positive ΔnD values, the minimum or maximum of both values exist at almost the same mole fraction of IL. This opposition of signs between VE and ΔnD mainly due to there will be less free volume available (if VE is negative) than in an ideal solution and photons will be more likely to interact with molecules or ions constituting the compound.51−54 Overall, our experimental results explicitly elucidate that there is a hydrogen bonding between ions of ILs with water molecules. To understand hydrogen bonding interactions between ammonium-based ILs and water, we further studied semiempirical calculations with the help of HyperChem 7 and these results displayed in Figure 3. We calculated heat of formation (ΔHf) of the complexes and compared the values with those of the ILs and water (as displayed in Table 4). In all the cases, ΔHf of the complex resulting from hydrogen bonding was higher than the sum of ΔHf ’s of water and ILs. It is reasonable to assume that these differences (ΔHf), calculated according to eq 3, represent the energies of the hydrogen bond. The energies of the hydrogen bonding can also obtained by using the total binding energies of water and ILs, presented in Table 4, instead of ΔHf ’s for these calculations. The results in 5980



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(3) Kumar, A.; Venkatesu, P. Overview of the Stability of αChymotrypsin in Different Solvent Media. Chem. Rev. 2012, 112, 4283−4307. (4) Rantwijk, F. V.; Sheldon, R. A. Biocatalysis in Ionic Liquids. Chem. Rev. 2007, 107, 2757−2785. (5) Seddon, K. R. Ionic Liquids for Clean Technology. J. Chem. Technol. Biotechnol. 1997, 68, 351−356. (6) Wasserscheid, P.; Keim, W. Ionic Liquids-New Solutions for Transistion Metal Catalysis. Angew. Chem., Int. Ed. 2000, 39, 3772− 3789. (7) Yoshizawa, M.; Narita, A.; Ohno, H. Design of Ionic Liquids for Electro Chemical Aplications. Aust. J. Chem. 2004, 57, 139−144. (8) Attri, P.; Venkatesu, P.; Kumar, A. Activity and Stability of αChymotrypsin in Biocompatible Ionic Liquids: Enzyme Refolding by Triethyl Ammonium Acetate. Phys. Chem. Chem. Phys. 2011, 13, 2788−2796. (9) Attri, P.; Venkatesu, P. Thermodynamic Characterization of the Biocompatible Ionic Liquid Effects on Protein Model Compounds and their Functional Groups. Phys. Chem. Chem. Phys. 2011, 13, 6566− 6575. (10) Reddy, P. M.; Venkatesu, P. Ionic Liquid Modifies the Lower Critical Solution Temperature (LCST) of Poly(N-isopropylacrylamide) in Aqueous Solution. J. Phys. Chem. B 2011, 115, 4752−4757. (11) Cooper, E. R.; Andrews, C. D.; Wheatley, P. S.; Webb, P. B.; Wormald, P.; Morris, R. E. Ionic liquids and Eutectic Mixtures as Solvent and Template in Synthesis of Zeolite Analogues. Nature. 2004, 430, 1012−1016. (12) Branco, L. C.; Crespo, J. G.; Afonso, C. A. M. Highly Selective Transport of Organic Compounds by Using Supported Liquid Membranes Based on Ionic Liquids. Angew. Chem. Int. Ed 2002, 41, 2771−2773. (13) Attri, P.; Venkatesu, P. Exploring the Thermal Stability of αChymotrypsin in Protic Ionic Liquids. Process Biochemistry. 2013, 48, 462−470. (14) Wei, W.; Danielson, N. D. Fluorescence and Circular dichroism Spectroscopy of Cytochrome c in Alkylammonium Formate Ionic Liquids. Biomacromolecules. 2011, 12, 290−297. (15) Indrani, J.; Attri, P.; Venkatesu, P. Unexpected effects of the Alteration of Structure and Stability of Myoglobin and Hemoglobin in Ammonium-based Ionic Liquids. Phys. Chem. Chem. Phys. 2014, 16, 5514−5526. (16) Shamsi, S. A.; Danielson, N. D. Utility of Ionic Liquids in Analytical Separations. J. Sep. Sci. 2007, 30, 1729−1750. (17) Reddy, P. M.; Venkatesu, P. Influence of Ionic Liquids on the Critical Micellization Temperature of a Tri-block co-polymer in Aqueous Media. J. Colloid Interface Sci. 2014, 420, 166−173. (18) Weng, J.; Wang, C.; Li, H.; Wang, Y. Novel Quaternary Ammonium Ionic Liquids and their use as Dual Solvent-Catalysts in the Hydrolytic Reaction. Green Chem. 2006, 8, 96−99. (19) Mehnert, C. P.; Dispenziereb, N. C.; Cook, R. A. Preparation of C9-aldehyde via Aldol Condensation Reactions in Ionic Liquid Media. Chem. Commun. 2002, 1610−1611. (20) Attri, P.; Venkatesu, P.; Hofman, T. Temperature Dependence Measurements and Structural Characterization of Trimethyl Ammonium Ionic Liquids with Highly Polar Solvent. J. Phys. Chem. B 2011, 115, 10086−10097. (21) Attri, P.; Reddy, P. M.; Venkatesu, P.; Kumar, A.; Hofman, T. Measurements and Molecular Interactions for N,N-Dimethylformamide with Ionic Liquid Mixed Solvents. J. Phys. Chem. B 2010, 114, 6126. (22) Attri, P.; Venkatesu, P.; Kumar, A. Temperature Effect on the Molecular Interactions between Ammonium Ionic Liquids and N,NDimethylformamide. J. Phys. Chem. B 2010, 114, 13415−13425. (23) Govinda, V.; Attri, P.; Venkatesu, P.; Venkateswarlu, P. Thermophysical Properties of Dimethylsulfoxide with Ionic Liquids at Various Temperatures. Fluid Phase Equilib. 2011, 304, 35−43. (24) Govinda, V.; Attri, P.; Venkatesu, P.; Venkateswarlu, P. Temperature Effect on the Molecular Interactions between Two

Scheme 1. Schematic Depiction of Strong Interactions between Water and Ions of IL and the Weak Interactions between Ions of IL

interaction of cation and anion of ILs is very different that results in the variation of strength results in different VE values. It is quite clear that the hydrogen bonding is clearly appearing between ions of ammonium family ILs and the oxygen atom of water. Interestingly, the semiempirical calculations are very well correlated with our experimental results, clearly affirming that there is an interaction between the water and ammonium ILs.

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CONCLUSIONS In this study, new data of ρ, u, η, and nD were measured at 25 °C and atmospheric pressure for six binary systems of ammonium-based ILs with water. From these experimental data, VE or Δκs or Δη or ΔnD were calculated and fitted to a Redlich−Kister type equation. Our results reveal that the ρ values decrease as the cation alkyl chain length increases whereas an opposite trend was observed, in which the values of η increase with increasing the number of carbon atoms in the alkyl chain length of cation of ILs. The experimental data indicate that cation and anion of ILs have a strong effect on the excess and deviation properties, especially on excess molar volume. The experimental results show that strong hydrogen bonding between ions of ILs with water molecules. These observed interactions are supported by our molecular modeling calculations, which are obtained by HyperChem 7.

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AUTHOR INFORMATION

Corresponding Author

*(P.V.) E-mail: [email protected]; pvenkatesu@ chemistry.du.ac.in. Telephone: +91-11-27666646-142. Fax: +91-11-2766 6605. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We acknowledge the financial support from the Department of Science and Technology (DST), New Delhi, India (Grant No. SB/SI/PC-109/2012).

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REFERENCES

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