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Influence of Aprotic Cosolvents on the Thermophysical Properties of Imidazolium-Based Ionic Liquid Fuxin Yang, Qian Ma, Xiaopo Wang,* and Zhigang Liu Key Laboratory of Thermo-Fluid Science and Engineering, Ministry of Education, Xi’an Jiaotong University, Xi’an, Shaanxi 710049, China S Supporting Information *

ABSTRACT: Imidazolium-based ionic liquids (ILs) have been widely investigated in the biofuels process. In this work, the experimental densities and viscosities of binary mixtures of IL, 1-octyl-3-methylimidazolium chloride, with aprotic cosolvents (i.e., N,N-dimethylacetamide, N,N-dimethylformamide, dimethyl sulfoxide, and pyridine) were studied at temperatures ranging from (303.15 to 353.15) K at atmospheric pressure. The Vogel−Fulcher−Tammann equation is used to correlate the viscosity data. Toward further understanding the influences of cosolvents on the properties of IL, the excess properties of density and viscosity deviations are calculated as well as the energy barrier. The effects of cosolvents on the density and viscosity are discussed.

1. INTRODUCTION Because of the low vapor pressure, large conductivity, high thermal stability, and large liquidus range, ionic liquids (ILs) are considered as green solvents and widely investigated during the past few years in chemical and biochemical applications. Moreover, there are only about 600 molecular solvents in use nowadays, while there are at least a million binary ionic liquids available due to the tunable and designable features of ILs by combining various of organic cations and organic or inorganic anions.1 Rogers et al. reported the pioneering work that the cellulose can be dissolved in ILs of 1-alkyl-3-methylimidazolium chloride ([Cnmim][Cl], where n is the number of carbon in alkyl chain).2 And then, more and more studies have focused on the use of ILs as the alternative solvents to make biofuels.3−7 In recent research, hydrophilic ILs, for example, 1-allyl-3-methylimidazolium chloride, 1-butyl-3-methylimidazolium chloride ([C4mim][Cl]), and 1-ethyl-3-methylimidazolium acetate ([C2mim][OAc]), show the ability for cellulose dissolution.8 Though ILs have many unique properties, they have disadvantages for making biofuels: high viscosity and limited miscibility with reaction products or nonpolar reagents.9 The viscosity will influence the dissolution rate of cellulosecontaining natural products in ILs; furthermore, undesirable high viscosity would impede the process of making biofuels from biomass.10 One possible way to decrease the viscosity of IL is to use cosolvent or diluent as the additive. Therefore, appropriate cosolvents are always selected to improve the influences of these two problems related to the use of ILs.9 Gericke et al. systematically studied the effects of 18 cosolvents on cellulose dissolution; among these solvents, N,N-dimethylacetamide © XXXX American Chemical Society

(DMA), N,N-dimethylformamide (DMF), dimethyl sulfoxide (DMSO), and pyridine (PYR), show great promise.9 The literature is rich in the study of cosolvents (e.g., DMSO) with ILs, [Cnmim][Cl], for bioenergy;11−21 however, little attention has been given to quantitatively study the effects of cosolvents on the thermophysical properties of ILs that are of importance, and necessary for the biochemical process design. In this work, due to its unique properties and wide use in chemical and biochemical processes, IL of [C8mim][Cl] was conducted. The aprotic cosolvents of DMA, DMF, DMSO, and PYR were selected as the additives to reduce the viscosity of IL. The experimental densities and viscosities for these binary mixtures were determined in the temperature range from (303.15 to 353.15) K at atmospheric pressure. The effects of these cosolvents on the densities and viscosities of IL are investigated as well as the interactions of the IL with cosolvents.

2. EXPERIMENTAL SECTION 2.1. Materials. The specifications for the cosolvents and the IL studied in this work are presented in Table 1. All of the cosolvents were purchased from Sigma-Aldrich (St Louis, MO). IL [C8mim][Cl] was obtained from the Center for Green Chemistry and Catalysis (CGCC), Lanzhou Institute of Chemical Physics (LICP), Chinese Academy of Sciences (CAS). The cosolvents were used without further purification. 2.2. Procedure. The IL was dried by 3A molecular sieves (Sigma-Aldrich, St Louis, MO) that were well-washed by acetone and methanol to eliminate the ion leaching from the sieves. The Received: January 1, 2017 Accepted: April 4, 2017

A

DOI: 10.1021/acs.jced.7b00002 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

Article

Table 1. Specifications of the Substances Studied in This Work substance

abbreviation

CAS no.

source

initial mole fraction purity

purification method

1-octyl-3-methylimidazolium chloride N,N-dimethylacetamide N,N-dimethylformamide dimethyl sulfoxide pyridine

[C8mim][Cl] DMA DMF DMSO PYR

64697-40-1 127-19-5 68-12-2 67-68-5 110-86-1

CAS Sigma-Aldrich Sigma-Aldrich Sigma-Aldrich Sigma-Aldrich

≥0.98 ≥0.98 ≥0.99 ≥0.99 ≥0.998

drying none none none none

molecular sieves were dried at 473.15 K in a tube furnace overnight, and then the sieves were added into the IL. The IL with the molecular sieves was dried in a vacuum oven at least for 24 h at 353.15 K under the pressure of 2 ± 0.1 kPa. The binary mixtures were made gravimetrically by using an analytical balance (Mettler-Toledo, AB204-N and ME204) with the uncertainty of 0.0001 g. All the samples were prepared and stored in a glovebox that was filled with nitrogen to minimize the contamination of air moisture. The moisture of the sample before and after experimental measurement was determined by using a Karl Fischer moisture titrator (Coulometric titration, MKC710B, Kyoto Electronics Manufacturing Co., Ltd.). The water contents before and after the experiment in all of the samples are under 0.3% in mass. 2.3. Density Measurements. The densities of the pure liquids and mixtures were measured using an Anton Paar digital vibrating U-tube densimeter (model DMA 5000 M) with the repeatability of 1.0 × 10−6 g·cm−3 provided by the manufacturer. The temperature was controlled and determined by two integrated Pt 100 platinum thermometers together with Peltier elements that can provide a more precise thermostat. The temperature of the Anton Paar densimeter was regulated to ±0.01 K. The densimeter was calibrated by using the dry air as well as the degassed and bidistilled water obtained from Anton Paar GmbH. The relative standard uncertainty of the density is 0.001. During the experiment, the sample density was measured in triplicate. 2.4. Viscosity Measurements. The viscosity values were determined by using an Ubbelohde capillary viscometer (model 9721-R62, 9721-R71, 9721-R77, and 9721-R80) purchased from Cannon instrument company (State College, PA, USA). The capillary viscometer was put into a Lauda oil thermostat bath, and the uncertainty of the temperature was within 0.01 K. The measurement was not conducted until the temperature was stable. The flow-time of the sample through the capillary was measured by using an electronic stopwatch with an uncertainty of 0.01 s. The dynamic viscosity was calculated using the following equation: η = ktρ

Table 2. Density Values of Pure Substances in the Temperature of 303.15 K at 0.0976 MPaa density/g·cm−3

a

substance

this work

literature


1.00620 0.93185 0.93918 1.09029 0.97309

1.0068,22 1.0059423 0.93169,24 0.9316225 0.93900,25 0.9394626 1.09041,24 1.0904927 0.973228,28 0.97298029

The standard uncertainties (u) are u(T) = 0.01 K and ur(ρ) = 0.001.

of the pure substances and binary mixtures of IL with cosolvents were determined in the temperatures from (303.15 to 353.15) K at atmospheric pressure (0.0976 MPa) and presented in Tables 3 to 6. The density of pure IL is considered as a function of temperature and correlated using the following equation: ρ = a0 + a1·T + a 2 ·T 2

(2)

−3

−3

−3

−1

where ρ (g·cm ) is the density; a0 (g·cm ), a1 (g·cm ·K ), and a2 (g·cm −3·K −2) are the parameters fitted by the experimental data; T is the temperature in Kelvin. The average absolute relative deviation (AARD) between the correlated values and the experimental data is calculated by the following equation: AARD(%) =

100 N

N

∑ i=1

(Ecal, i − Eexp, i) Eexp, i

(3)

where N is the number of experimental points, Ecal and Eexp are the calculated and experimental data, respectively. The densities calculated by eq 2 fitted with the experimental data in this work are compared with those collected from the literature. The AARD values are as follows: 0.001%, this work; 0.70%, Seddon et al.;30 0.03%, Gomez et al.;23 0.04%, AlTuwaim et al.;31 0.10%, Ning et al.;32 0.65%, Yan et al.;33 0.32%, He et al.;34 0.04%, Dowell et al.;22 0.03%, Singh et al.35 The relative deviations are depicted in Figure S1 (in Supporting Information). It is clear that there are no significant differences between the experimental densities and the reported values. 3.2. Experimental Viscosities and Derived Properties. The experimental viscosities of pure IL and cosolvents in this work and those in the literature are presented in Table 7.31,36−41 Similarly, no significant discrepancies are observed between the experimental viscosities of the cosolvents and the literature values; however, considerable divergences are displayed among different authors for the IL viscosities. The viscosities of the pure substances and binary mixtures of IL with cosolvents were determined in the temperatures ranging from (303.15 to 353.15) K at atmospheric pressure and are reported in Tables 8 to 11. The Vogel−Fulcher−Tammann (VFT) equation is used to fit the measured viscosities as a function of temperature.42−44

(1)

where η is the dynamic viscosity, k is the viscometer constant, t is the sample efflux time, ρ is the density of the sample corresponding to the same temperature. Because all of the flow times were greater than 100 s and the capillary diameter (0.47 mm) was far less than its length (290 mm), in the experimental measurements, the kinetic energy and the end corrections were ignored. The uncertainty of the viscosity in this study is less than 2% with a confidence level of 0.95. During the experiment, the sample viscosity was measured in triplicate.

3. RESULTS AND DISCUSSIONS 3.1. Experimental Density Data. The experimental density values of IL and pure cosolvents in this work and those data already studied are reported in Table 2.22−29 The density values B



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Table 3. Experimental Densities of IL with the Co-solvent of DMA at 0.0976 MPaa x [C8mim][Cl] + (1- x) DMA, ρ/g·cm−3 x T/K 303.15 308.15 313.15 318.15 323.15 328.15 333.15 338.15 343.15 348.15 353.15 a

1.000

0.877

0.786

0.602

0.468

0.361

0.273

0.139

0

1.00620 1.00333 1.00048 0.99764 0.99482 0.99200 0.98917 0.98638 0.98363 0.98088 0.97813

1.00375 1.00080 0.99786 0.99492 0.99200 0.98912 0.98624 0.98337 0.98049 0.97762 0.97477

1.00137 0.99836 0.99534 0.99237 0.98943 0.98649 0.98355 0.98062 0.97769 0.97478 0.97187

0.99487 0.99172 0.98857 0.98542 0.98229 0.97916 0.97605 0.97296 0.96987 0.96679 0.96372

0.98864 0.98530 0.98199 0.97868 0.97538 0.97209 0.96881 0.96554 0.96228 0.95903 0.95578

0.98193 0.97845 0.97497 0.97150 0.96804 0.96459 0.96114 0.95771 0.95428 0.95085 0.94743

0.97502 0.97138 0.96774 0.96411 0.96049 0.95687 0.95326 0.94965 0.94605 0.94244 0.93885

0.95809 0.95407 0.95012 0.94626 0.94247 0.93872 0.93499 0.93123 0.92734 0.92335 0.91935

0.93185 0.92726 0.92266 0.91806 0.91344 0.90881 0.90416 0.89953 0.89487 0.89019 0.88551

The standard uncertainties (u) are u(x) = 2.0 × 10−4, u(T) = 0.01 K and ur(ρ) = 0.001.

Table 4. Experimental Densities of IL with the Co-solvent of DMF at 0.0976 MPaa x [C8mim][Cl] + (1 − x) DMF, ρ/g·cm−3 x T/K 303.15 308.15 313.15 318.15 323.15 328.15 333.15 338.15 343.15 348.15 353.15 a

1.000

0.858

0.739

0.557

0.424

0.321

0.241

0.119

0

1.00620 1.00333 1.00048 0.99764 0.99482 0.99200 0.98917 0.98638 0.98363 0.98088 0.97813

1.00431 1.00136 0.99842 0.99547 0.99257 0.98970 0.98683 0.98397 0.98110 0.97824 0.97539

1.00201 0.99895 0.99595 0.99297 0.99000 0.98703 0.98405 0.98109 0.97813 0.97519 0.97225

0.99635 0.99319 0.98993 0.98682 0.98392 0.98080 0.97765 0.97448 0.97156 0.96834 0.96507

0.99200 0.98863 0.98526 0.98190 0.97856 0.97522 0.97188 0.96856 0.96524 0.96192 0.95862

0.98599 0.98247 0.97905 0.97572 0.97237 0.96903 0.96562 0.96216 0.95867 0.95518 0.95169

0.97910 0.97540 0.97170 0.96799 0.96431 0.96061 0.95695 0.95332 0.94969 0.94602 0.94233

0.96557 0.96127 0.95731 0.95345 0.94936 0.94525 0.94111 0.93697 0.93283 0.92869 0.92455

0.93918 0.93441 0.92964 0.92486 0.92007 0.91526 0.91044 0.90560 0.90074 0.89587 0.89098


Table 5. Experimental Densities of IL with the Co-solvent of DMSO at 0.0976 MPaa x [C8mim][Cl] + (1 − x) DMSO, ρ/g·cm−3 x T/K 303.15 308.15 313.15 318.15 323.15 328.15 333.15 338.15 343.15 348.15 353.15 a

1.000

0.862

0.750

0.574

0.439

0.338

0.253

0.130

0

1.00620 1.00333 1.00048 0.99764 0.99482 0.99200 0.98917 0.98638 0.98363 0.98088 0.97813

1.01118 1.00824 1.00531 1.00240 0.99947 0.99658 0.99372 0.99088 0.98804 0.98520 0.98237

1.01541 1.01238 1.00941 1.00646 1.00351 1.00056 0.99759 0.99464 0.99175 0.98890 0.98596

1.02526 1.02209 1.01892 1.01575 1.01258 1.00943 1.00628 1.00315 1.00003 0.99693 0.99383

1.03429 1.03091 1.02754 1.02418 1.02083 1.01751 1.01419 1.01088 1.00759 1.00431 1.00104

1.04277 1.03920 1.03565 1.03211 1.02858 1.02506 1.02155 1.01805 1.01456 1.01107 1.00760

1.05157 1.04784 1.04409 1.04039 1.03665 1.03295 1.02924 1.02555 1.02186 1.01818 1.01451

1.06788 1.06371 1.05955 1.05541 1.05126 1.04712 1.04298 1.03885 1.03472 1.03060 1.02647

1.09029 1.08528 1.08030 1.07531 1.07032 1.06533 1.06034 1.05534 1.05035 1.04534 1.04033


η = η0 exp[B /(T − T0)]

Table 12 presents the fitted parameters for eq 4 along with the average absolute relative deviation between the correlated values and the experimental data. It is shown that the correlated viscosities using the equation of VFT agree well with the experimental values. The viscosities calculated using the VFT equation fitted in this work are compared with the values reported in the literature. The AARD from AlTuwaim et al.31 is 7%, from Gomez et al.23 is 26%,

(4)

where η is the viscosity in mPa·s; T is the temperature in Kelvin; η0 (mPa·s), B (K), and T0 (K) are the empirical parameters determined from the experimental data. The fitted parameter of T0 (K) is considered as the “ideal glass transition temperature” that should be lower than the experimental glass transition temperature.30 C



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Table 6. Experimental Densities of IL with the Co-solvent of PYR at 0.0976 MPaa x [C8mim][Cl] + (1 − x) PYR, ρ/g·cm−3 x T/K 303.15 308.15 313.15 318.15 323.15 328.15 333.15 338.15 343.15 348.15 353.15 a

1.000

0.865

0.755

0.578

0.445

0.339

0.255

0.128

0

1.00620 1.00333 1.00048 0.99764 0.99482 0.99200 0.98917 0.98638 0.98363 0.98088 0.97813

1.00551 1.00258 0.99966 0.99675 0.99380 0.99092 0.98806 0.98522 0.98237 0.97956 0.97669

1.00583 1.00279 0.99979 0.99682 0.99386 0.99090 0.98795 0.98499 0.98204 0.97911 0.97618

1.00443 1.00129 0.99815 0.99501 0.99188 0.98875 0.98564 0.98253 0.97944 0.97636 0.97329

1.00278 0.99945 0.99614 0.99285 0.98957 0.98628 0.98300 0.97967 0.97642 0.97305 0.96948

1.00033 0.99684 0.99334 0.98985 0.98638 0.98290 0.97942 0.97593 0.97240 0.96890 0.96527

0.99825 0.99455 0.99084 0.98714 0.98345 0.97975 0.97606 0.97237 0.96868 0.96500 0.96131

0.98890 0.98471 0.98051 0.97631 0.97210 0.96786 0.96358 0.95929 0.95503 0.95066 0.94627

0.97309 0.96806 0.96302 0.95797 0.95290 0.94782 0.94272 0.93760 0.93246 0.92730 0.92212


The energy barrier Eη must be overcome to move ions up on the other ions and it is described as follows:43

Table 7. Viscosity Values of Pure Substances in the Temperature of 303.15 K at 0.0976 MPaa viscosity/mPa·s substance

this work

literature


7339.71 0.89 0.77 1.80 0.84

7770.431 0.877,36 0.8737 0.766,38 0.7675,36 1.8034,36 1.78839 0.820,40 0.857141

Eη = R ×

∂ ln η

( T1 )

∂

⎛ ⎜ = R⎜ ⎜ ⎝

B

(

T02 T2

−

2T0 T

⎞ ⎟ ⎟ +1 ⎟ ⎠

)

(5)

where R is the ideal gas constant (approximate 8.3145 × 10−3 kJ· K−1·mol−1); η (mPa·s), B(K), and T0(K) are determined from eq 4. The energy barriers of the samples at 303.15 K are presented in Table 12. 3.3. Effects of Cosolvents on the Densities and Viscosities of IL. The effects of temperature on the density values for the binary mixtures of IL with cosolvents are depicted in Figures S3 to S6. It is shown that the densities in all the systems present a linear dependence with the temperature, where the values decrease with the increase of the temperature. Clearly, DMSO has the highest density value among the studied cosolvents and the pure IL, and hence, the binary mixture of DMSO with IL has the highest density compared to the corresponding mixtures of IL with other cosolvents. Toward further understanding the effect of cosolvent on the density of IL, the excess molar volume is introduced to evaluate the deviation of the binary molar volume from the ideality:

a The standard uncertainty (u) is u(T) = 0.01 K, and ur(η) = 0.02 with a 0.95 level of confidence.

from Seddon et al.30 is 31%, and from Ning et al.32 is 37%. The relative deviations are depicted in Figure S2. Considerable deviation is observed. It is well-known that the thermophysical properties, especially the viscosity, will be affected by the impurities. One of the possible reasons for the discrepancies between the experimental values and literature data is the source where IL of [C8mim][Cl] was purchased. Because different manufacturers have different production procedures, the classes and contents of impurities are generally different which may cause the divergences. Furthermore, the sample preparation and the experimental measurement are the additional factors to cause the divergences.

Table 8. Experimental Viscosities of IL with the Co-solvent of DMA at 0.0976 MPaa x [C8mim][Cl] + (1 − x) DMA, η/mPa·s x

a

T/K

1.000

0.877

0.786

0.602

0.468

0.361

0.273

0.139

0

303.15 308.15 313.15 318.15 323.15 328.15 333.15 338.15 343.15 348.15 353.15

7339.71 4272.86 2766.92 1788.59 1198.65 828.30 588.69 426.28 316.66 237.14 181.41

1865.36 1254.97 853.26 594.67 425.36 310.37 231.30 176.46 136.74 107.08 85.14

1049.50 714.83 502.24 358.94 263.25 198.27 150.45 116.53 91.77 73.67 59.69

288.58 210.24 156.62 119.22 92.49 72.76 58.34 47.46 38.89 32.38 27.28

92.29 70.68 55.22 44.20 35.89 29.51 24.70 20.81 17.72 15.26 13.20

32.45 26.51 21.93 18.32 15.51 13.24 11.44 9.98 8.73 7.70 6.85

14.75 12.52 10.65 9.20 8.02 7.04 6.22 5.53 4.96 4.47 4.05

3.63 3.26 2.94 2.65 2.42 2.21 2.04 1.88 1.74 1.62 1.51

0.89 0.83 0.78 0.74 0.70 0.66 0.63 0.60 0.57 0.54 0.52

The standard uncertainties (u) are u(x) = 2.0 × 10−4, u(T) = 0.01 K, and ur(η) = 0.02 with a 0.95 level of confidence. D



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Table 9. Experimental Viscosities of IL with the Co-solvent of DMF at 0.0976 MPaa x [C8mim][Cl] + (1 − x) DMF, η/mPa·s x

a

T/K

1.000

0.858

0.739

0.557

0.424

0.321

0.241

0.119

0

303.15 308.15 313.15 318.15 323.15 328.15 333.15 338.15 343.15 348.15 353.15

7339.71 4272.86 2766.92 1788.59 1198.65 828.30 588.69 426.28 316.66 237.14 181.41

1892.76 1257.96 856.69 596.05 426.48 312.38 233.56 176.82 137.16 108.17 86.03

766.79 536.62 381.17 277.21 206.44 156.59 120.90 94.93 75.65 61.29 50.08

160.48 122.67 95.06 74.18 59.06 47.65 38.97 32.38 27.13 22.93 19.60

49.08 39.48 32.20 26.50 22.11 18.67 15.92 13.71 11.91 10.41 9.16

18.99 15.91 13.51 11.56 10.00 8.72 7.67 6.79 6.06 5.43 4.90

8.90 7.69 6.75 5.96 5.29 4.75 4.25 3.85 3.50 3.20 2.93

2.56 2.32 2.13 1.95 1.81 1.67 1.56 1.45 1.36 1.27 1.20

0.77 0.72 0.68 0.65 0.62 0.59 0.56 0.54 0.51 0.49 0.47

The standard uncertainties (u) are u(x) = 2.0 × 10−4, u(T) = 0.01 K, and ur(η) = 0.02 with a 0.95 level of confidence.

Table 10. Experimental Viscosities of IL with the Co-solvent of DMSO at 0.0976 MPaa x [C8mim][Cl] + (1 − x) DMSO, η/mPa·s x

a

T/K

1.000

0.862

0.750

0.574

0.439

0.338

0.253

0.130

0

303.15 308.15 313.15 318.15 323.15 328.15 333.15 338.15 343.15 348.15 353.15

7339.71 4272.86 2766.92 1788.59 1198.65 828.30 588.69 426.28 316.66 237.14 181.41

3235.63 2150.90 1404.41 943.52 653.90 466.77 338.65 250.15 189.55 145.53 113.93

1084.71 742.47 508.33 363.38 264.93 197.17 149.71 115.48 90.80 72.27 58.87

371.36 264.22 192.92 143.95 109.49 84.97 67.04 53.77 43.60 35.76 29.90

143.43 107.70 82.42 63.95 50.83 41.01 33.28 27.38 22.97 19.28 16.49

55.12 43.80 34.98 28.15 23.25 19.36 16.31 13.91 11.95 10.37 9.06

24.14 19.97 16.54 13.89 11.80 10.12 8.78 7.66 6.75 5.98 5.34

6.85 6.00 5.24 4.62 4.10 3.67 3.31 3.00 2.72 2.50 2.29

1.80 1.66 1.53 1.41 1.30 1.21 1.13 1.06 0.99 0.93 0.88


Table 11. Experimental Viscosities of IL with the Co-solvent of PYR at 0.0976 MPaa x [C8mim][Cl] + (1 − x) PYR, η/mPa·s x

a

T/K

1.000

0.865

0.755

0.578

0.445

0.339

0.255

0.128

0

303.15 308.15 313.15 318.15 323.15 328.15 333.15 338.15 343.15 348.15 353.15

7339.71 4272.86 2766.92 1788.59 1198.65 828.30 588.69 426.28 316.66 237.14 181.41

2828.10 1823.55 1210.73 829.67 583.82 419.77 309.69 233.02 177.92 138.26 108.95

863.91 599.38 426.00 309.61 230.68 174.40 134.84 105.82 84.08 68.04 55.20

290.77 215.57 161.31 123.18 96.06 75.98 61.06 49.81 41.06 34.18 28.74

84.05 65.58 51.96 42.02 34.40 28.53 23.98 20.37 17.42 15.01 13.03

29.41 24.13 20.12 16.97 14.49 12.56 10.81 9.51 8.37 7.40 6.63

12.71 10.87 9.39 8.17 7.17 6.34 5.65 5.06 4.56 4.13 3.76

3.46 3.11 2.81 2.56 2.34 2.15 1.98 1.83 1.70 1.58 1.48

0.84 0.79 0.74 0.70 0.66 0.62 0.59 0.56 0.53 0.51 0.48


V E = Vm −

∑ xiVi =

∑ xiMi − ρm

∑

xiMi ρi

In general, the excess molar volumes of the binary mixtures can be correlated by using the Redlich−Kister equation. In this study, a Redlich−Kister type equation is applied:45

(6)

where VE is the excess molar volume; Vm and ρ are the molar volume and density of the binary mixture, respectively; xi, Vi, Mi, and ρi are the mole fraction, molar volume, molar mass, and density of pure component i, respectively. In this study, component i presents the ionic liquid.

V E = xi(1 − xi)(A 0 + A1T + A 2 xi)

(7)

where VE is the excess molar volume; xi is the mole fraction of pure component i; T is the temperature in Kelvin; A0 (cm3· E



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Table 12. Parameters for eq 4, Average Absolute Relative Deviations, and the Energy Barriers at 303.15 K x IL

η0/mPa·s

B/K

1 0.877 0.786 0.602 0.468 0.361 0.273 0.139 0

0.0593 0.0082 0.0343 0.0702 0.1512 0.0829 0.0888 0.0693 0.0379

1274.9917 1820.9856 1338.9987 1050.0349 741.4858 846.5357 754.3681 692.8760 757.5740

0.858 0.739 0.557 0.424 0.321 0.241 0.119 0

0.0198 0.0283 0.0336 0.0752 0.0984 0.0909 0.0698 0.0415

1540.4585 1387.2945 1271.6566 925.8031 758.9275 718.2117 676.2715 736.8911

0.862 0.75 0.574 0.439 0.338 0.253 0.13 0

0.0056 0.0221 0.0570 0.0534 0.0565 0.0628 0.0654 0.0322

1958.2803 1441.5211 1088.6706 1043.9978 961.5062 871.5626 750.5071 917.6832

0.865 0.755 0.578 0.445 0.339 0.255 0.128 0

0.0341 0.0429 0.0472 0.1107 0.1091 0.0930 0.0745 0.0355

1400.2071 1288.8451 1202.6921 851.9269 773.4445 747.5010 675.2365 747.9764

T0/K

AARD/%

Eη (303.15 K)/kJ·mol−1

194.5479 155.5762 173.5005 176.9677 187.5435 161.3662 155.6190 128.1749 63.1220

0.77 1.04 0.26 0.09 0.21 0.09 0.10 0.07 0.08

82.60 63.89 60.87 50.39 42.39 32.18 26.48 17.29 10.05

168.8250 167.2598 153.0706 160.3289 158.9123 146.4131 115.2616 50.4572 DMSO 155.6255 169.6892 179.1888 170.9433 163.5097 156.7773 141.8843 75.2515 PYR 179.5347 173.0949 165.3473 174.7047 164.9210 151.1550 127.2284 67.0499

0.52 0.47 0.48 0.14 0.10 0.11 0.08 0.07

65.24 57.40 43.14 34.68 27.87 22.34 14.64 8.82

1.35 0.73 0.11 0.18 0.37 0.23 0.18 0.21

68.75 61.84 54.13 45.64 37.68 31.08 22.05 13.50

0.20 0.14 0.35 0.16 0.24 0.02 0.08 0.04

70.02 58.22 48.39 39.46 30.93 24.72 16.67 10.25

DMA

DMF

mol−1), A1 (cm3·mol−1·K−1), and A2 (cm3·mol−1) are the empirical parameters determined from the experimental data. The parameters along with the standard deviations in excess molar volume σ(VE) are presented in Table 13. Table 13. Parameters for the Redlich−Kister Equation and the Standard Deviations in Excess Molar Volume σ(VE) A0 system IL-DMA IL-DMF IL-DMSO IL-PYR

cm ·mol 3

A1 −1

2.6519 2.5995 0.8456 4.3130

−1

cm ·mol ·K 3

−0.0289 −0.0301 −0.0141 −0.0366

A2 −1

standard deviation −1

cm ·mol 3

4.6396 5.7176 2.3989 5.8202

σ(VE) 0.04 0.06 0.02 0.04

The excess molar volumes as a function of IL mole fraction are depicted in Figures 1 to 4 and Figure S7. The interactions that occur in these binary mixtures are complex.9 Regarding the excess molar volumes, all of the samples exhibit negative deviations compared to the ideal binary mixtures, indicating that denser molecular packing might be formed.46 Moreover, the maximum negative deviations occur at the mole fraction of IL in

Figure 1. Excess molar volumes for binary mixture of IL with DMA. Experimental: ■, 303.15 K; □, 308.15 K; ●, 313.15 K; ○, 318.15 K; ▲, 323.15 K; △, 328.15 K; ▼, 333.15 K; ▽, 338.15 K; ◆, 343.15 K; ◇, 348.15 K; ★, 353.15 K; , calculated by Redlich−Kister type equation.

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Figure 2. Excess molar volumes for binary mixture of IL with DMF. Experimental: ■, 303.15 K; □, 308.15 K; ●, 313.15 K; ○, 318.15 K; ▲, 323.15 K; △, 328.15 K; ▼, 333.15 K; ▽, 338.15 K; ◆, 343.15 K; ◇, 348.15 K; ★, 353.15 K; , calculated by Redlich−Kister type equation.

Figure 4. Excess molar volumes for binary mixture of IL with PYR. Experimental: ■, 303.15 K, □, 308.15 K, ●, 313.15 K, ○, 318.15 K, ▲, 323.15 K, △, 328.15 K, ▼, 333.15 K, ▽, 338.15 K, ◆, 343.15 K, ◇, 348.15 K, ★, 353.15 K; , calculated by Redlich−Kister type equation.

Figure 3. Excess molar volumes for binary mixture of IL with DMSO. Experimental: ■, 303.15 K; □, 308.15 K; ●, 313.15 K; ○, 318.15 K; ▲, 323.15 K; △, 328.15 K; ▼, 333.15 K; ▽, 338.15 K; ◆, 343.15 K; ◇, 348.15 K; ★, 353.15 K; , calculated by Redlich−Kister type equation.

Figure 5. Viscosities of IL with the cosolvent of DMA as a function of IL mole fraction: ■, 303.15 K; ●, 313.15 K; ▲, 323.15 K; ▼, 333.15 K; ◆, 343.15 K; ★, 353.15 K.

the intensity of the hydrogen-bonding interactions, and then the viscosity of IL decreases substantially. Regarding the binary mixtures of the IL with cosolvents, the viscosities studied in this work present a similar order to that of reported by Prausnitz et al.10 Tables 8 to 11 and Figures 5 to 8 show that the cosolvents remarkably lower the viscosities of pure IL. In Table 8, when a mole fraction of cosolvents of DMA is 0.123 (weight fraction of 0.05), the viscosity of IL falls from 7339.71 mPa·s to 1865.36 mPa·s at 303.15 K. Moreover, in Figure 5, when the mole fraction of cosolvent of DMA is higher than 0.398 (weight fraction of 0.2), the curve of IL viscosity becomes gentle, and the values of viscosities do not change much. Similar to the influence of DMA, the viscosity of IL will be significantly reduced even when a small concentration of DMF, DMSO, or PYR is loaded. As shown in Figure 9, for a fixed temperature, for instant at 303.15 K, the influences of cosolvents on lowering the viscosity of IL can be ordered as follows: DMA > DMF > PYR > DMSO. Figure S9 depicts the energy barrier as a function of IL mole fraction. When the IL mole fraction x is higher than 0.8, the energy barriers

0.1−0.5. Additionally, among these four cosolvents, in absolute values, small VE deviations for the mixture of IL with DMSO from the ideality are observed, indicating a less strength of interaction between IL and the cosolvent of DMSO. The viscosities of cosolvents as a function of temperature are depicted in Figure S8. Among the four cosolvents, DMSO has the highest viscosity and the other three cosolvents show no considerable differences in values. Figures 5−8 present the viscosities of binary mixtures as a function of IL mole fraction. Clearly, the viscosities of IL with cosolvents decrease with the increase of temperature. The viscosity of IL drops from 7339.71 mPa·s at 303.15 K to 181.41 mPa·s at 353.15 K. It is well-known that the composition of anions and cations has a favorable influence on the viscosity of IL. Furthermore, the IL viscosity is mainly governed by the intermolecular forces of van der Waals and Coulombic interactions, especially the formation ability of hydrogenbonding. The increase of temperature will efficiently reduce G



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Figure 6. Viscosities of IL with the cosolvent of DMF as a function of IL mole fraction: ■, 303.15 K; ●, 313.15 K; ▲, 323.15 K; ▼, 333.15 K; ◆, 343.15 K; ★, 353.15 K.

Figure 9. Viscosities of IL with the cosolvents as a function of IL mole fraction at 303.15 K: ■, DMA; ●, DMF; ▲, DMSO; ▼, PYR.

decrease as a sequence: DMSO > PYR > DMF > DMA. The energy barrier sequence coincides with the order of the viscosity. Therefore, the higher is the viscosity, the more energy barrier is required to move ions. For further understanding the interactions between the IL and cosolvents, the viscosity deviations were obtained based on the following equation: Δη = ηm −

∑ xiηi

(8)

where Δη is viscosity deviation; ηm is the viscosity of the mixture; xi and ηi are the mole fraction and the viscosity of pure component i, respectively. Figures 10−13 show the viscosity deviations as a function of IL mole fraction. It is observed that, in the four binary systems, the

Figure 7. Viscosities of IL with the cosolvent of DMSO as a function of IL mole fraction: ■, 303.15 K; ●, 313.15 K; ▲, 323.15 K; ▼, 333.15 K; ◆, 343.15 K; ★, 353.15 K.

Figure 10. Viscosity deviations for ILs with the cosolvent of DMA as a function of IL mole fraction: ■, 303.15 K; □, 308.15 K; ●, 313.15 K; ○, 318.15 K; ▲, 323.15 K; △, 328.15 K; ▼, 333.15 K; ▽, 338.15 K; ◆, 343.15 K; ◇, 348.15 K; ★, 353.15 K.

viscosity deviations are negative at all temperatures in the studied composition. Larger negative viscosity deviations are displayed at low temperatures, indicating the strong intermolecular interactions in these mixtures, and the absolute values of viscosity deviations increase with the decrease of temperature in all binary mixtures.

Figure 8. Viscosities of IL with the cosolvent of PYR as a function of IL mole fraction: ■, 303.15 K; ●, 313.15 K; ▲, 323.15 K; ▼, 333.15 K; ◆, 343.15 K; ★, 353.15 K.

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Figure 11. Viscosity deviations for IL with the cosolvent of DMF as a function of IL mole fraction: ■, 303.15 K; □, 308.15 K; ●, 313.15 K; ○, 318.15 K; ▲, 323.15 K; △, 328.15 K; ▼, 333.15 K; ▽, 338.15 K; ◆, 343.15 K; ◇, 348.15 K; ★, 353.15 K.

Figure 13. Viscosity deviations of IL with the cosolvent of PYR as a function of IL mole fraction: ■, 303.15 K; □, 308.15 K; ●, 313.15 K; ○, 318.15 K; ▲, 323.15 K; △, 328.15 K; ▼, 333.15 K; ▽, 338.15 K; ◆, 343.15 K; ◇, 348.15 K; ★, 353.15 K.

mixtures, and among the four cosolvents, small negative VE deviations for the mixture of IL with DMSO are observed, indicating that a less strength of interaction between IL and cosolvent of DMSO might be formed. In the conversion of biomass to biofuels, the high viscosity of IL has a great influence on the process. One possible way for reducing the viscosity is to use cosolvent or diluent as the additive. In the study, even when a small concentration of cosolvents is loaded, the viscosity of IL is significantly reduced. Three cosolvents of DMA, DMF, and PYR have similar influences on reducing the viscosity of IL, while DMSO presents the highest viscosity deviations from the ideal binary mixture and holds the highest values of the solvent parameters. This study indicates that thermophysical characteristics of IL [C8mim][Cl], especially the viscosity, can be improved by using the cosolvents as additives for the chemical process to make biofuels from biomass.

■

Figure 12. Viscosity deviations of IL with the cosolvent of DMSO as a function of IL mole fraction: ■, 303.15 K; □, 308.15 K; ●, 313.15 K; ○, 318.15 K; ▲, 323.15 K; △, 328.15 K; ▼, 333.15 K; ▽, 338.15 K; ◆, 343.15 K; ◇, 348.15 K; ★, 353.15 K.

ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jced.7b00002. Relative deviations between the calculated values using the correlation equation fitted with the experimental data in this work and literature values for pure IL; effects of temperature on the density values for the binary mixtures of IL with cosolvents; excess molar volumes as a function of IL mole fraction; viscosities of cosolvents as a function of temperature; energy barrier as a function of IL mole fraction (PDF)

Solvatochromism is widely performed to evaluate the interactions in the solutions by using the empirical solvent parameters: normalized empirical polarity (ENT ), solvent acidity (α), and solvent basicity (β).9 The parameters for these four cosolvents are DMA (ENT = 0.377, α = 0, β = 0.76), DMF (ENT = 0.386, α = 0, β = 0.69), DMSO (ENT = 0.444, α = 0, β = 0.76), and PYR (ENT = 0.302, α = 0, β = 0.64). The values of α for the cosolvents are zero, indicating that the solvents lack the ability to perform as the hydrogen-bond donor; while the values of β for the cosolvents are more than 0.4, indicating that the solvents possess relatively high basicity and capacity to act as hydrogenbond acceptor. Among these four cosolvents, DMSO performs the highest values of the solvent parameters as well as the highest viscosity deviations as shown in Figures 10 to 13.

■

AUTHOR INFORMATION

Corresponding Author

*Tel: +86 (0)29-82668210. Fax: +86 (0)29-82663584. E-mail: [email protected].

4. CONCLUSIONS In this work, experimental density and viscosity data for the binary mixtures of IL with cosolvents (i.e., DMA, DMF, DMSO, and PYR) are presented at atmospheric pressure from (303.15 to 353.15) K. Negative molar volumes are observed in all binary

ORCID

Fuxin Yang: 0000-0001-9640-3231 Funding

The authors thank the National Natural Science Foundation of China for financial assistance (No. 51606147). I



Article

Notes

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The authors declare no competing financial interest.

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