Solution Thermodynamics of Imidazolium-Based Ionic Liquids and

Article pubs.acs.org/jced

Solution Thermodynamics of Imidazolium-Based Ionic Liquids and Volatile Organic Compounds: Benzene and Acetone Chang-Hao Li,†,‡ Kai-Xuan Gao,† Yu-Ning Meng,† Xian-Kun Wu,† Feng Zhang,*,† and Zhi-Xiang Wang‡ †

School of Chemistry and Chemical Engineering, Nanjing University, Nanjing 210093, China School of Engineering, China Pharmaceutical University, Nanjing 210009, China

‡

S Supporting Information *

ABSTRACT: Vapor−liquid equilibria (VLE) of benzene and acetone in the selected eight imidazole ionic liquids (ILs) was studied with a static method. VLE measurements were carried out over an extremely low concentration range and at temperatures ranging from 303.2 to 333.2 K. On the basis of the solubility data, the infinite dilution activity coefficient (γ∞) and Henry’s constant (H) were derived and described formally by using the nonrandom two-liquid equation. The obtained results indicated that the anion species of ionic liquids play a great role in determining the solubility of benzene and acetone relative to cationic species, and increase following the sequence [Tf2N] > [PF6] > [BF4]. The difference of solubility between acetone and benzene mainly depends on the difference of their chemical and structural properties. Additionally, the partial molar excess enthalpies at infinite dilution (HE,∞ i ) have been derived from the temperature dependence of the limiting activity coefficients, and partial molar enthalpies and entropies at infinite dilution for benzene and acetone absorbed in ILs have also been derived from the temperature dependence of the Henry’s constant.

1. INTRODUCTION Volatile organic compounds (VOCs) released from industrial processes cause serious environmental problems, such as photochemical smoke and acid rain. VOCs emission is becoming one of the most stringent environment challenges in industrial processes, thus many approaches have been tried to cut down the emission. Absorption of VOCs by various solvents has been proven to be an effective way. Different kinds of solvents such as organic compounds,1 silicone oil, and emulsion2,3 have been extensively studied for VOCs absorption. However, their volatile nature and regenerative property limit their application. In recent years, ionic liquids (ILs) have received special attention as potential candidates of a solvent for VOCs, owing to their unique characteristics such as extremely low volatility, high thermal stability, and designable structure.4,5 Moreover, ILs are constructed with organic factors and thus affiliative to VOCs. ILs are regarded as environmentally benign green solvents with wide applications such as catalysis,6 extraction,7 membrane separation,8 gas separation,9−11 and electrolysis.12 Owing to their high affinity, ILs show high efficiency for the removal of industrial waste gases. Milota’s13 research group have screened out an IL with excellent absorption capability and high stability, tetradecy(trihexyl)phosphonium dicyanamide, for the removal of industrial gases. The ILs achieved a 78 % reduction of total VOCs. Additionally, absorption of VOCs in ILs could significantly decrease energy consumption, meanwhile providing the collected pollutants in a concentrated stream that can be used for fuel or distilled into products. VOCs that are dissolved in ILs can be selectively and readily removed under low pressure, and the ILs also can be easily recovered. © XXXX American Chemical Society

To remove VOCs effectively with ILs, it is important to prepare the appropriate ionic liquids by optimizing their chemical structures and characterizing their physicochemical properties. In general, the concentration of VOCs in industry exhaust gas is much lower than the saturation concentration. Thus, an investigation of low pressure VOCs solubility in ILs is necessary to obtain thermodynamic properties under real application conditions. This is also important to improve the industrial use of IL in VOCs removal. In the real industrial process, ILs can be mixed with low viscosity solutions to prepare new kinds of absorbents which aid in the stability and lipophilicity of ILs. The thermodynamic properties of mixtures derived from VOCs solubility, include vapor−liquid equilibria (VLE) data, infinite dilution activity coefficient (γ∞), Henry’s constant, enthalpy and entropy of solvation in ILs. Generally, the dynamic methods14 are adopted to investigate the kinetics characteristic of VOCs absorption, and static methods are convenient for determining the absorbency described as Henry’s law constant.1−3 The kinetics characteristic and the absorbency property give a guidance to select suitable absorbent for VOC removal. As for ILs, they are also helpful to design the structure of ILs with high absorb capacity and ILbased separation process.15−17 The activity coefficient of a solute, at infinite dilution, is an important property. It can be used to estimate the volatility of the solute as well as the intermolecular energy between solute and solvent, and to provide information about selection of Received: October 26, 2014 Accepted: April 27, 2015

A

DOI: 10.1021/je500986b J. Chem. Eng. Data XXXX, XXX, XXX−XXX

Journal of Chemical & Engineering Data

Article

suitable solvents for extraction and extractive distillation.18,19 Thermodynamic properties, which act as one evaluation criterion, have been analyzed for selecting cations and anions to enhance the absorption of VOCs at low concentration by ILs. The lower is the γ∞ in ILs, the more favorable are the intermolecular interactions between ILs and solute (the higher solubility in ILs). On the basis of the difference of the γ∞, a criteria that the VOCs-IL system corresponding to γ∞ < 1 stands for attractive interaction with ILs, and oppositely, γ∞ > 1 presents unfavorable interaction, is proposed to select the IL structures that promote favorable solute−solvent intermolecular interactions.16 The γ∞ can be determined by several methods including gas−liquid chromatography (GLC),20 the dilutor technique,21 ebulliometry,22 and the static method.23 Until now, the experimental data for γ∞ in ILs have been reported for the numerous organic compounds, such as aromatic hydrocarbons, alkanes, ketones, and alcohols.20,21,24,25The VLE data covering the whole range of less volatile solutes concentrations in ILs has been reported using a static vapor pressure method.26,27 Moreover, the γ∞ can be obtained by extrapolating VLE data to infinite dilution. Compared with the values obtained by the GLC method, there is a deviation in γ∞ value, which arises most probably from insufficient VLE data with small mole fractions.28−30 The physicochemical properties such as density, viscosity, heat capacities, and vapor pressure of conventional imidazolium-based ILs have been extensively investigated. Imidazoliumbased ILs have been confirmed to have high affinity for VOCs compounds, and the partition coefficients are comparable or even higher than those of typical organic solvents.15,17 However, there is no systemic investigation on the affinity of low concentration VOCs toward a wide variety of ILs. In this work, by using a static method, the solubilities of benzene and acetone in several ILs were evaluated to obtain their thermodynamic properties under low pressure range, and directions to select ILs for the industrial process of VOCs removal is provided. Low vapor pressure liquid equilibrium behavior of benzene and acetone in eight ionic liquids: 1-nethyl-3-methylimidazolium tetrafluoroborate ([emim][BF4]), 1-n-ethyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide ([emim][Tf2N]), 1-n-butyl-3-methylimidazolium tetrafluoroborate ([bmim][BF4]), 1-n-butyl-3-methylimidazolium hexafluorophosphate ([bmim][PF6]), 1-n-butyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide ([bmim][Tf2N]), 1n-hexyl-3-methylimidazolium tetrafluoroborate ([hxmim][BF4]), 1-n-hexyl-3-methylimidazolium hexafluorophosphate ([hxmim][PF6]), and 1-n-hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide ([hmim][Tf2N]) over the temperature range (303.2 to 333.2) K were also investigated. On the basis of the nonrandom two-liquid (NRTL) activity coefficient model, the activity coefficients in infinite dilution and Henry’s constants were obtained from VLE data. In addition, the partial molar excess enthalpies at infinite dilution (HiE,∞) and partial molar enthalpies and entropies of absorption process were derived.

Shanghai Chengjie Reagent Co., Ltd., in the highest available purity (purity > 99 %). Before using them, all the ionic liquids were further degassed by vacuuming at 80 °C over 48 h. 2.2. Vapor−Liquid Equilibrium Apparatus and Measurements. Apparatus used for the VOCs absorption (shown in Figure 1) is similar to the one used for CO2 absorption.31

Figure 1. Schematic diagram of absorption apparatus.

However, there is a slight change that the VOCs source vessel replaced the CO2 gas cylinder to store liquid benzene or acetone and generate the vapor. The volumes of the storage vessel and absorption vessel were 118.95 cm3 (V1) and 49.35 cm3 (V2), respectively. The storage vessel, is used to store the vapor of VOCs derived from the source vessel and isolates the VOCs before they contact with the absorbents in the absorption vessel. The latter used as the equilibrium cell is equipped with a magnetic stirrer. The temperatures of the setup are controlled by a water bath with an uncertainty of ± 0.1 K. The pressure in the two vessels is monitored by two pressure transducers (Fujing Wideplus Precision Instruments Co., Ltd.) with ± 0.02 % uncertainty in the relation to the full scale. The pressure transducers are connected to a Numeric Instrument to record the pressure online every 3 s. In a typical run of experiment, the VOCs source vessel was degassed after closing valve 1, then the selected liquid VOC was sucked into the vessel through the vacuum degree. To acquire pure liquid vapor without foreign gas, the VOCs source vessel was degassed three times by a vacuum pump at low temperature (2 °C). Absorbent of a known amount (m) was placed in the absorption vessel and the air in the two chambers was evacuated. Then valves 3 and 2 were closed. The pressure of storage and absorption vessels was recorded to be P0. VOC vapor was introduced into the storage vessel through valve 1 until a predetermined pressure (P) was reached. Then after valve 2 was opened, the pressure (P) decreased as the gas was absorbed by absorbent until the pressure became constant for at least 15 min. The absorption was deemed to reach equilibrium, and this pressure was denoted as equilibrium pressure (Pe). Solubility data was continuously determined by feeding more VOCs vapor into the storage vessel to reach new equilibrium. In each solubility measurement, it would take about 60 s to get to the equilibrium state. The upper pressure limit of VOCs was dependent on the saturation pressure at the experimental temperature. The VOCs uptake, nVOCs, can be calculated by measuring the variation of the pressure in the absorption vessel. The solubility of VOCs in the ILs was defined in terms of Henry’s constant.

■

EXPERIMENTAL SECTION 2.1. Materials. Analytical grade benzene and acetone (purity > 99.5 %) were obtained from Nanjing Chemical Reagent Co., Ltd. and used without further purification. The ionic liquids, namely, [emim][BF4], [emim][Tf2N], [bmim][BF 4 ], [bmim][PF 6 ], [bmim][Tf2N], [hxmim][BF 4 ], [hxmim][PF6], and [hmim][Tf2N] were purchased from B



Article

⎤ ⎡ ⎛ ⎞2 G21 τ12G12 ⎥ ln γ1 = x 22⎢τ21⎜ ⎟ + ⎢⎣ ⎝ x1 + x 2G21 ⎠ (x 2 + x1G12)2 ⎥⎦

3. RESULTS AND DISCUSSION 3.1. Solubility of Benzene and Acetone in ILs. The solubility of VOCs in eight ionic liquids (ILs) ([emim][BF4], [emim][Tf2N], [bmim][BF4], [bmim][PF6], [bmim][Tf2N], [hmim][BF4], [hmim][PF6] and [hmim][Tf2N]) was obtained at temperatures ranging from 303.2 K to 333.2 K. Considering the condensation of benzene and acetone vapor, the vapor equilibrium was measured only up to approximately 30 % of their saturated vapor pressure. The solubility [VLE (vapor liquid equilibria) (T, P, x)] data of benzene and acetone, along with their uncertainties, are shown in tables in the Supporting Information. The amount of absorbed VOCs can be calculated by the following equation: nVOCs =

PV1 + P0(V2 − m /ρ) − Pe(V1 + V2 − m /ρ) RT

where Gij = exp(−αijτij), τij = (gij − gjj)/RT (i,j = 1,2; i ≠ j) and the nonrandomness parameter α is assumed to be a constant of 0.2 in this work; the two energy parameters, τ12 and τ21, are supposed to be linear functions of temperature, which can be rewritten as eqs 7 and 8. The adjustable parameters in these equation are listed in Table 2.

where M is the molecular weight of ionic liquid. 3.2. Data Correlation. The solubility of VOCs in ILs is a result of the phase equilibrium, therefore the low pressure VLE for absorbed VOCs can be described by equation given below: (3)

where yi and xi are the vapor phase mole fraction and the liquid phase mole fraction, respectively; P represents the system pressure; Psi stands for the saturated vapor pressure, and γi is the activity coefficient. ϕVi and ϕsi are the vapor phase fugacity coefficients of experimental pressure and saturated vapor pressure, respectively. In this work, because of the nonvolatility of ILs, it is reasonable to assume yi = 1. Considering the low system pressure, the vapor fugacity coefficients (ϕVi and ϕsi ) can be supposed to be unity. Hence, eq 3 can be rewritten as P ≈ γixiPis

(4)

B T + C − 273.15

(5)

The parameters in eq 5 are shown in Table 1. Table 1. Parameters in Equation 5 species

A

B

C

benzene acetone

3.98523 4.2184

1184.24 1197.01

217.572 228.06

(7)

(0) (1) τ21 = τ21 + τ21 /T

(8)

system

τ(0) 12

τ(1) 12

τ(0) 21

τ(1) 21

benzene/[emim][BF4] benzene/[bmim][BF4] benzene/[hmim][BF4] benzene/[bmim][PF6] benzene/[hmim][PF6] benzene/[emim][Tf2N] benzene/[bmim][Tf2N] benzene/[hmim][Tf2N] acetone/[emim][BF4] acetone/[bmim][BF4] acetone/[hmim][BF4] acetone/[bmim][PF6] acetone/[hmim][PF6] acetone/[emim][Tf2N] acetone/ [bmim][Tf2N] acetone/ [hmim][Tf2N]

0.9363 0.8349 0.5212 0.9119 0.8646 0.3873 0.3281 0.4401 0.267 1.2269 0.0249 0.0247 0.3167 0.2474 0.2351 0.4201

56.127 −24.517 26.192 −76.596 −165.57 −81.794 −104.07 −200 50.677 −334.93 −37.244 −90.354 −118.42 −326.48 −333.66 −400.8

0.3442 0.1142 0.4411 0.1949 −0.2985 −0.0863 −0.0995 −0.4296 0.3995 −0.7994 −0.0301 −0.0572 −0.0297 −0.5285 −0.7145 −0.8399

−123.27 −67.356 −165.32 −90.78 75.076 −27.37 6.0359 116.26 −184.4 179.82 −30.833 −20.199 −121.33 191.24 246.6 281.86

The calculated mole fraction solubility data (xcal) as presented in the Tables, and the isothermal P−x diagrams of benzene and acetone in investigated ILs are added in the Supporting Information. The average relative deviations (ARD) under experimental conditions are shown in Table 3. In short, the experimental results agreed well with the NRTL equation, with a relative deviation less than 2 %. The vapor liquid equilibria data for benzene and acetone have been widely reported previously.28,35−40 However, few solubility data have been reported for low concentration range (x < 0.2) in our work. As depicted in Figure 2, low vapor liquid equilibria data of benzene in [bmim][BF4] and [bmim][Tf2N] were compared with the reference data. Under similar conditions, the equilibria pressure of benzene binary mixtures with [bmim][BF4] obtained in this work are lower than those of ref 35, while for [bmim][Tf2N], the pressure values are slightly higher than those of ref 28. 3.3. Infinite Dilution Activity Coefficient and Henry’s Constant. The low pressure solubility of VOCs in ILs can be reported as infinite dilution activity coefficients and Henry’s constant. The γ∞ for benzene and acetone in the ionic liquid could be calculated from the infinite dilution form of the NRTL equation at a given temperature.

With respect to the saturation pressure of VOCs, an Antoinetype equation for temperature variations32 was used as ln Pis = A −

τ12 = τ12(0) + τ12(1)/T

Table 2. Parameters of NRTL Activity Coefficient Model

(1)

where m and ρ is the weight and density of ionic liquid, respectively. P0 stands for the initial pressure of the absorption vessel. The solubility of benzene or acetone can be expressed in mole fraction (x): nVOCs x= nVOCs + m /M (2)

ϕi V yP = γixiϕisPis i

(6)

Several activity models are available in the literature,33,34 and the NRTL activity coefficient model was confirmed to give the best empirical description of the activity coefficients.28 Thus, it was used to correlate the solubility results of VOCs−ILs (eq 6). For each isothermal solubility datum, the activity coefficient γ1 was calculated at each observed x1 point using eq 4.

ln(γ ∞) = τ21 + τ12 exp( −ατ12)

(9)

The Henry’s constant, Hi(T) is defined as

Hi(T ) = γ ∞Pis C

(10) DOI: 10.1021/je500986b J. Chem. Eng. Data XXXX, XXX, XXX−XXX


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Table 3. Average Relative Deviationa (ARD %) between Calculated and Experimental Values in VOCs-IL Systems system benzene/[emim][BF4] benzene/[bmim][BF4] benzene/[hmim][BF4] benzene/[bmim][PF6] benzene/[hmim][PF6] benzene/[emim][Tf2N] benzene/[bmim][Tf2N] benzene/[hmim][Tf2N] acetone/[emim][BF4] acetone/[bmim][BF4] acetone/[hmim][BF4] acetone/[bmim][PF6] acetone/[hmim][PF6] acetone/[emim][Tf2N] acetone/[bmim][Tf2N] acetone/[hmim][Tf2N]

experimental condition 0 kPa to 11 kPa, 303.2 K to 333.2 0 kPa to 15 kPa, 303.2 K to 333.2 0 kPa to 13 kPa, 303.2 K to 333.2 0 kPa to 11 kPa, 303.2 K to 333.2 0 kPa to 15 kPa, 303.2 K to 333.2 0 kPa to 13 kPa, 303.2 K to 333.2 0 kPa to 15 kPa, 303.2 K to 333.2 0 kPa to 14 kPa, 303.2 K to 333.2 0 kPa to 12 kPa, 303.2 K to 333.2 0 kPa to 10 kPa, 303.2 K to 333.2 0 kPa to 14 kPa, 303.2 K to 333.2 0 kPa to 12 kPa, 303.2 K to 333.2 0 kPa to 16 kPa, 303.2 K to 333.2 0 kPa to 12 kPa, 303.2 K to 333.2 0 kPa to 14 kPa, 303.2 K to 333.2 0 kPa to 19 kPa, 303.2 K to 333.2

Table 4. Henry’s Constant (Hi) for Benzene in Various ILs Hi (kPa)

N

ARD %

ILs

303.2 K

313.2 K

323.2 K

333.2 K

32

0.72

34

0.68

31

0.58

32

0.93

31

0.65

[emim][BF4] [bmim][BF4] [hmim][BF4] [bmim][PF6] [hmim][PF6] [emim][Tf2N] [bmim][Tf2N] [hmim][Tf2N]

33.04 26.81 24.40 25.37 19.78 17.45 14.05 11.75

51.42 41.79 38.01 39.53 30.89 27.21 21.97 18.38

77.45 62.93 57.31 59.65 46.64 41.17 33.38 27.96

113.29 91.97 83.88 87.24 68.65 60.55 49.15 41.17

31

1.03

28

0.58

30

0.46

K K K K K

Table 5. Henry’s Constant (Hi) for Acetone in Various ILs

K

Hi (kPa)

K K 31

0.54

28

0.24

30

0.30

30

1.23

29

0.41

28

0.27

29

0.46

31

0.84

K K K K

ILs

303.2 K

313.2 K

323.2 K

333.2 K

[emim][BF4] [bmim][BF4] [hmim][BF4] [bmim][PF6] [hmim][PF6] [emim][Tf2N] [bmim][Tf2N] [hmim][Tf2N]

45.71 34.71 30.00 25.04 21.99 15.74 14.92 14.04

69.17 52.44 45.06 37.81 33.70 24.11 22.75 21.57

101.66 76.87 65.75 55.44 50.11 35.80 33.67 32.11

144.99 109.72 93.48 79.21 72.53 51.74 48.48 46.49

K

counteranion changed from [BF4] to [PF6] or [Tf2N], the γ∞ and Hi gradually decreased. The length of the alkyl chain does slightly enhance the solubility of benzene and acetone as exemplified by alkyl-substituted groups of ILs from ethyl to butyl or hexyl. Thus, the solubility of benzene and acetone can be greatly affected by the anion species of ILs relative to the cationic species. The γ∞ values of acetone are much smaller than those of benzene due to the difference in their chemical and structural properties. Acetone has strong polar functional groups which generate attractive hydrogen bonding between the solute and the ILs. As for benzene, the high electron density above and below the plane of the ring produces quadrupole and higher electrostatic multipolar moment. Mainly, ILs and benzene have π−π interactions between aromatic and imidazolium rings. When anion species of ILs are altered from [BF4] to [PF6] or [Tf2N], the behavior of the benzene vapor dissolved in the ILs gradually changes sign from a positive deviation from Raoult’s law to negative. The acetone absorbed in ILs all commonly exhibit negative deviation from ideality. The VOCs-ILs system exhibit activity coefficients γ∞ < 1, which indicate the attractive forces between molecules of solute and IL are greater than the solute−solute or IL−IL cohesive forces. Thus, it is potentially more favorable to absorb VOCs by selecting those cations and anions that provide the high negative deviation from ideality. The γ∞ values of benzene and acetone reported previously by other researchers are summarized in Table 6. Taking the difference of experimental methods and data processing into consideration, the values of γ∞ obtained in our work agree well with literature results, validating the reliability of the apparatus and data processing employed in this work. 3.4. Partial Molar Excess Enthalpy at Infinite Dilution and Enthalpy and Entropy of Absorption. According to the Gibbs−Helmholtz equation, the value of the partial molar excess enthalpy at infinite dilution HE,∞ can be directly i obtained from the slope of a straight line derived from eq 13:

K K K

exp exp Average relative deviations, ARD % = 100/N∑iN= 1|(xcal i − xi )/xi |; N, Number of vapor−liquid equilibrium data points.

a

Figure 2. Comparison of solubility data of benzene in [bmim][BF4] and [bmim][Tf2N] (filled symbol: ■ and ▲, ref 35: ▼ and ◀, ref 28. Open symbol, experimental data).

Table 4 and 5 list the calculated values of Hi(T) for benzene and acetone, respectively. The experimental results of the γ∞ are presented in Table 6. As can be seen, Henry’s constant increases with a rise in temperature, indicating the typical physical absorption of the VOCs in ILs. Generally, the larger the Hi is, the less the VOCs solubility is in ILs. Infinite dilution activity coefficients (γ∞) slightly increase when temperature rises. Obviously, as the D



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Table 6. Values of γ∞ Calculated in This Work and Reported in the Literature ILs

303.2 K

313.2 K

323.2 K

333.2 K

ref

Benzene

a

[emim][BF4]a [emim][BF4]b [bmim][BF4]a [bmim][BF4]b [hmim][BF4]a [hmim][BF4]b [bmim][PF6]a [bmim][PF6]b [hmim][PF6]a [hmim][PF6]b [emim][Tf2N]a [emim][Tf2N]b [bmim][Tf2N]a [bmim][Tf2N]b [hmim][Tf2N]a [hmim][Tf2N]b

2.09 2.42 (303 K) 1.69 1.739 (303.15 K) 1.54 1.63 (303 K) 1.60

[emim][BF4]a [emim][BF4]b [bmim][BF4]a [bmim][BF4]b [hmim][BF4]a [hmim][BF4]b [bmim][PF6]a [bmim][PF6]b [hmim][PF6]a [emim][Tf2N]a [emim][Tf2N]b [bmim][Tf2N]a [hmim][Tf2N]a [hmim][Tf2N]b

1.21

1.63 1.82 (313.15 K) 1.27 1.04 (313.15 K) 1.12 1.177 (313.8 K) 0.90 0.881 (303.15 K) 0.76 0.76 (313.15 K) Acetone 1.23

0.92 1.157 0.77

0.93 1.186 0.79

0.66

0.67 0.707 (315.9 K) 0.60 0.43 0.404 (303.15 K) 0.40 0.38 0.357 (312.25 K)

1.25 1.03 (298.15 K) 1.10 0.89 0.872 (293.15 K) 0.74 0.75 (303.15 K)

0.59 0.42 0.403 (293.15 K) 0.39 0.37 0.352 (301.75 K)

2.12

2.15 2.51 (323 K) 1.75 1.77 (323.15 K) 1.59 1.65 (323 K) 1.65 1.96 (323.15 K) 1.29 1.07 (323.15 K) 1.14 1.179 (323.1 K) 0.92 0.892 (313.15 K) 0.77 0.77 (323.15 K)

1.72 1.758 (313.15 K) 1.57

2.17 25 1.77 1.798 (333.15 K) 1.61

41 25

1.68 2.13 (333.15 K) 1.32

42 43

1.16 1.179 (333.6 K) 0.94 0.903 (323.15 K) 0.79 0.78 (333.15 K)

1.24 1.11 (323 K) 0.94 1.209 0.80 0.85 (323 K) 0.68 0.709 (325.3 K) 0.61 0.44 0.41 (313.15 K) 0.41 0.39 0.362 (322.85 K)

44 21 33

1.26 25 0.95 1.232 0.81

45 25

0.69 0.699 (334.8 K) 0.62 0.45 0.423 (323.15 K) 0.42 0.40 0.366 (333.35 K)

24

21

46

The infinite dilution activity coefficients were obtain from experimental correlation. bThe values of γ∞ are collected from the literature.

⎛ ∂ln γ ∞ ⎞ H E, ∞ i ⎟= i ⎜ R ⎝ ∂(1/T ) ⎠

Table 7. Gibbs Energy, Partial Molar Enthalpy, and Partial Molar Entropy of absorption for Benzene Vapor in ILs

(11)

ΔG|T=303.2 K

HE,∞ i

where R is the gas constant. The values of for benzene and acetone are listed in Tables 7 and 8, respectively. To make a comparison, the literature results are also summarized in Tables 7 and 8. Owing to the small slope of ln γ∞ i versus 1/T plots and considering the experimental uncertainty of the γ∞ i values, the obtained here are acceptable. The negative values of HE,∞ i values of the partial excess enthalpies mean that the interactions between the solute−ionic liquid pairs are larger than those of the solute−solute pairs. Meanwhile, the lower is the partial molar excess enthalpy, the higher is the interaction between followed the trend VOCs and ionic liquids. The values of HE,∞ i [Tf2N] > [PF6] > [BF4]. Hence, the [Tf2N] of three anions show slightly higher strength of the ion-induced interactions with ILs, possessing the higher solubility capability. Owing to their special chemical structure and interaction with ionic for benzene and acetone vapor do not liquid, the values of HE,∞ i show any significant difference. As discussed in the above section, absorption of VOCs in ILs is strongly temperature dependent. The analysis about thermodynamic of absorption, which can elucidate the behavior of absorption, usually focuses on the free energy (ΔG), enthalpy (ΔH), and entropy (ΔS). In general, these can be

IL [emim] [BF4] [bmim] [BF4] [hmim] [BF4] [bmim] [PF6] [hmim] [PF6] [emim] [Tf2N] [bmim] [Tf2N] [hmim] [Tf2N]

−1

kJ·mol

ΔH

ΔS|T=303.2 K −1

kJ·mol

−1

−1

J·mol ·K

a HE,∞ i

−1

J·mol

b HE,∞ i

J·mol−1

8.82

−34.50

−142.87

−1099.61

8.29

−34.51

−141.16

−1143.92

8.05

−34.57

−140.58

−1024.78

8.15

−34.10

−140.93

−1212.68

7.52

−34.81

−139.62

−1444.47

−102043

7.21

−34.84

−138.67

−1442.65

−65933

6.66

−35.06

−137.59

−1685.00

−92033

6.21

−35.14

−136.38

−1739.79

−151133

−107141

a

The values of partial molar excess enthalpies are calculated from experimental results. bThese values are reported in the literature.

calculated from the partial derivations of the Henry’s coefficients with respect to temperature: ΔG = RT ln Hi E

(12) DOI: 10.1021/je500986b J. Chem. Eng. Data XXXX, XXX, XXX−XXX


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Table 8. Gibbs Energy, Partial Molar Enthalpy, and Partial Molar Entropy of Absorption for Acetone Vapor in ILs ΔG|T=303.2 K IL [emim][BF4] [bmim][BF4] [hmim][BF4] [bmim][PF6] [hmim][PF6] [emim][Tf2N] [bmim][Tf2N] [hmim][Tf2N] a

kJ·mol

−1

9.64 8.94 8.51 8.11 7.82 6.95 6.79 6.66

ΔH

ΔS|T=303.2 K −1

−1

−1

J·mol ·K

kJ·mol

−32.42 −32.35 −32.65 −32.52 −32.71 −33.32 −33.29 −33.51

−138.49 −136.18 −135.75 −133.99 −133.67 −132.81 −132.20 −132.49

a HE,∞ i

J·mol

b HE,∞ i

−1

J·mol−1

−1179.26 −1103.18 −1388.02 −1000.34 −1529.44 −2079.165 −1748.684 −2265.32

−172345

−124333 −266533

b

The values of partial molar excess enthalpies are calculated from experimental results. The values are reported in the literature.

⎛ ∂ ln Hi ⎞ ΔH = R ⎜ ⎟ ⎝ ∂(1/T ) ⎠ P

ΔS =

⎡ ⎛ ∂ ln Hi ⎞ ⎤ ΔH − ΔG ⎟ ⎥ = −R ⎢ln Hi + T ⎜ ⎝ ∂T ⎠ P ⎥⎦ ⎢⎣ T

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ASSOCIATED CONTENT

S Supporting Information *

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Figures S1−S8 and S9−S16 list the isothermal P−x diagrams of benzene and acetone in eight imidazolium ionic liquids at a given temperature, respectively. Tables S1−S8 and S9−S16: experimental and calculated solubility (T, P, x) data of benzene and acetone with deviation, respectively. The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/je500986b.

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Generally, the enthalpy reflects the thermal effect of physical dissolution of benzene and acetone in ILs, whereas the entropy represents the level of ordering that takes place in the liquid/ gas mixture. For the VOCs absorption system, the absorption behavior is similar at different temperatures, so the absorption was taken at 303.2 K as a case. The data of ΔG, ΔH, and ΔS are presented in Tables 7 and 8. The entropy of absorption is a measure of the number of specific ways in which a thermodynamic system may be arranged, often taken to be a measure of disorder. The higher are the values of ΔS, the more ordered configuration is formed in the ILs, meanwhile, the least degree of freedom of restricted VOCs occurs after it absorbed into ILs. Obviously, the entropy of benzene and acetone vapor gradually increase when changing the anion species from [BF4] to [PF6] or [Tf2N] as well as lengthening the length of the substituted alkyl. In the gas absorption process, the larger entropy is more thermodynamically favorable, implying a strong molecular interaction. The [Tf2N] anion has the greater entropy of the three anions, resulting in better absorption capability. In conclusion, to improve trapping efficiency of VOCs, the [Tf2N] ILs with longer length of alkyl substituted cation are preferred.

■

AUTHOR INFORMATION

Corresponding Author

*Tel: 86-25-83596665. Fax: 86-25-83593772. E-mail: zf@nju. edu.cn. Funding

The authors are grateful for the financial support of the National Natural Science Foundation of China (No. 21306078 and No. 21476105). Notes

The authors declare no competing financial interest.

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REFERENCES

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4. CONCLUSION In this study, the solubility of benzene and acetone gas at low pressure in eight selected ILs was determined in detail. The observed vapor−liquid equilibrium behavior of VOCs has been very well correlated with the NRTL activity coefficient model. Henry’s constant, infinite dilution activity coefficients, partial molar excess enthalpy, and thermodynamic properties (e.g., enthalpy and entropy of dissolution) were derived. From thermodynamic analysis, it was found that the process of benzene and acetone vapor absorption is regulated by entropy, and the larger entropy is more thermodynamically favorable. It appears that anionic species play the most significant role in determining the solubility of benzene and acetone in ILs and increases in following order [Tf2N] > [PF6] > [BF4]. The longer is the length of alkyl substitutions of imidazolium cation, the greater is the solubility of benzene and acetone in ILs. F



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H


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